east meck chemistry -...

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East Meck Chemistry

STUDENT NAME:___________________

TEACHER/Period:________________ CHEMISTRY: THE STUDY OF MATTER

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Chemistry is the science that investigates and explains the structure and properties of ________________. This

includes its composition, properties and the changes it undergoes.

SCIENTIFIC METHOD

A scientific method is a systematic approach to answer a ______________________ or study a situation. It starts

with _____________________ - noting and recording facts. A _________________________ is a possible

explanation for what has been observed. It is an educated _________________ as to the cause of the problem or

answer to the question. An experiment is a set of controlled observations that _____________ a hypothesis. The

variable that is changed in an experiment is called the ________________________ variable. The variable that you

watch to see how it _________________ as a result of your changes to the independent variable is called the

dependent variable. The cycle (hypothesis followed by experimentation) repeats many times, and the hypothesis

gets more and more certain. The hypothesis becomes a _______________________, which is a thoroughly tested

model that explains why things behave a certain way. Theories can never be ____________________; they are

always subject to additional research. Another outcome is that certain behavior is repeated many times. A scientific

____________ describes a relationship in nature that is supported by many experiments and for which no exception

has been found.

Identify the dependent variable and the independent variable in the following experiments.

a) A student tests the ability of a given chemical to dissolve in water at three different temperatures.

independent variable: _________________________________

dependent variable: _________________________________

b) A farmer compares how his crops grow with and without phosphorous fertilizers.

independent variable: _________________________________

dependent variable: _________________________________

MATTER

Matter is anything that takes up __________________ and has mass. ______________ is the measure of the

amount of matter that an object contains. Virtually all of the matter around us consists of mixtures. A mixture can

be defined as something that has _____________________ composition. Soda is a mixture (carbon dioxide is

dissolved in it), and ____________________ is a mixture (it can be strong, weak or bitter). If matter is not uniform

throughout, then it is a _______________________ mixture. If matter is uniform throughout, it is homogeneous.

Homogeneous mixtures are called ___________________. A heterogeneous mixture contains regions that have

____________________ properties from those of other regions. When we pour sand into water, the resulting

mixture contains two distinct regions. ___________________ pavement, which has small rocks mixed with tarry

goo, is a simple example of a heterogeneous mixture. Oil-and-vinegar salad dressing, which has a layer of oil

floating on a layer of vinegar, is another example. Homogeneous mixtures (also known as solutions) are mixtures in

which the composition is _______________________, there are no chunks or layers. Salt water,

___________________ ___________________ and dust free air (mixture of nitrogen, oxygen, argon, carbon

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dioxide, water vapor and other gases) are examples of homogeneous mixtures. Brass (solid mixture of copper and

______________) is also a homogeneous mixture. Brass is a(n) _________________, which is a mixture of metals.

Since heterogeneous mixtures contain chunks or layers, they are often easier to separate than homogeneous

mixtures. A mixture of solid particles in a liquid can be separated by pouring the mixture through a

___________________ that traps the solid particles while the liquid passes through in a process called filtering.

Some simple methods also exist for separating homogeneous mixtures. A solid dissolved in a liquid solution can be

separated by letting it dry out in the process of ___________________. Mixtures are separated into pure

_____________________. A pure substance always has the same composition. Pure substances are either elements

or _________________________. Elements are substances that cannot be broken down into other substances

chemically or _______________________. Examples include sodium, carbon and aluminum. Compounds are

substances made of two or more ______________________ combined chemically. Compounds have properties

___________________________ from those of the original elements. Examples of compounds include water

(hydrogen and oxygen) and table salt (sodium and chlorine).

Classify each of the following as a pure substance, a homogeneous mixture or a heterogeneous mixture.

A. gasoline ____________________ B. copper metal ____________________

C. a stream with gravel at the bottom ____________________

D. maple syrup _________________ E. chunky peanut butter _____________

F. common salt ____________________ G. margarine ____________________

H. a Spanish omelet ________________ I. a multivitamin tablet ________________

J. oxygen gas ____________________ K. carbon dioxide gas _________________

PROPERTIES

The properties of matter describe the characteristics and behavior of matter, including the changes that matter

undergoes. _____________________ properties are characteristics that a sample of matter exhibits without any

change in its identity. This property can be observed and measured without _____________________ the

substance.

Examples of the physical properties of a chunk of matter include its:

1. __________________________________ 2. _________________________________

3. __________________________________ 4. _________________________________

5. __________________________________ 6. _________________________________

7. __________________________________

Chemical properties are those that can be observed only when there is a change in the

___________________________ of the substance. Rusting is a chemical reaction in which iron combines with

__________________ to form a new substance, iron (III) oxide.

Classify each of the following as a chemical or physical property.

density ___________________________ reactivity ___________________________

color _____________________________ melting point ________________________

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Using the Chemistry Reference Tables, which substance has a

A. density = 19.31 g/cm3 _____________________

B. melting point = -119°C _____________________

C. boiling point = 65°C _____________________

D. melting point = -73°C _____________________

Using the Chemistry Reference Tables, are the following substances soluble or insoluble in water?

A. zinc nitrate __________

B. sodium sulfate __________

C. calcium carbonate __________

D. potassium oxide __________

E. lead (II) fluoride __________

F. barium hydroxide __________

G. copper (II) sulfide __________

H. silver chloride __________

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CHANGES

A physical change is a change in matter that does not involve a change in the chemical identity of individual

substances. The matter only changes in appearance. Examples: ______________, _________________,

__________________, _________________, ___________________, and _____________________. A chemical

property always relates to a chemical change, the change of one or more substances _____________ other

substances. Another term for chemical change is chemical ___________________. Indications of a chemical

reaction: __________________ absorbed or released, _________________ change, formation of a precipitate -

______________ that separates from solution, and formation of a ___________. All matter is made of atoms, and

any chemical change involves only a rearrangement of the atoms. Atoms do not just appear. Atoms do not just

disappear. This is an example of the law of conservation of mass (or matter), which says that in a chemical change,

matter is neither ________________ nor destroyed. All chemical changes also involve some sort of energy change.

Energy is either taken in or __________________ ____________ as the chemical change takes place. Energy is the

capacity to do _________________. Work is done whenever something is moved. Chemical reactions that give off

heat energy are called ____________________ reactions. Chemical reactions that _________________ heat energy

are called endothermic reactions. Freezing, condensation and ___________________ are exothermic. Melting,

_______________________ and sublimation are endothermic.

State whether each of the following is an endothermic or exothermic process.

1. melting of ice __________________________

2. combustion of gasoline __________________________

3. Natural gas is burned in a furnace. __________________________

4. When solid potassium bromide is dissolved in water, the solution gets colder. _____________________

DENSITY

Density is the amount of matter (mass) contained in a unit of ___________________. Styrofoam has a low density

or small mass per unit of volume.

volume

massdenisty =

V

mD =

Solve the following density problems.

1. The density of sugar is 1.59 g/cm3. Calculate the mass of sugar in 15.0 ml. (1 mL = 1 cm

3).

2. The density of helium is 0.178 g/L. Calculate the volume of helium that has a mass of 23.5 g.

3. A 14.95 g sample of gold has a volume of 0.774 cm3. Calculate the density of gold.

4. Balsa wood has a density of 0.12 g/cm3. What is the mass of a sample of balsa wood if its volume is 134 cm

3?

5. The density of a sample of lead is found by the process of water displacement. A graduated cylinder is filled

with water to the 30.0 mL mark. The cylinder with the water is placed on an electronic balance and weighs

106.82 g. A piece of lead is added to the cylinder. The cylinder is reweighed with the water and the lead and

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the scale reads 155.83 g. The volume of all the material in the cylinder is 34.5 mL. Calculate the density of the

lead.

6. The density of an unknown solid was found by the process of water displacement. The object was massed on an

electronic balance. The balance reads 125 g. 50.0 cm3 of water was poured into a 100.0 mL graduated

cylinder. The unknown sample was then gently placed into the graduated cylinder. The volume in the cylinder

rose to 60.7 cm3. Calculate the density of the unknown solid.

MODERN VIEW OF THE ATOM

The atom has two regions and is ___-dimensional. The nucleus is at the ___________________ and contains the

protons and _____________________. The electron cloud is the region where you might find an electron and most

of the volume of an atom. The atomic _________________ of an element is the number of protons in the nucleus of

an atom of that element. The number of protons determines ____________________ of an element, as well as

many of its chemical and physical properties. Because atoms have no overall electrical charge, an atom must have

as many ____________________ as there are protons in its nucleus. Therefore, the atomic number of an element

also tells the number of electrons in a neutral atom of that element. The mass of a neutron is almost the same as the

mass of a ________________. The sum of the protons and neutrons in the nucleus is the ________________

number of that particular atom. _____________________ of an element have different mass numbers because they

have different numbers of _______________, but they all have the same atomic number.

Practice / Homework

Density: Use the ref. packet to identify the substance based on the density value given D = m / V

1. D = 0.66g/cm3

2. D = 2.702g/cm3

3. m = 20 g, V = 4.44 cm3

4. m = 3 g, V = 2.1 L

Melting and Boiling points: Use the ref. packet to identify the substance based on the given temperature value.

5. Melting point = 801oC

6. Boiling point = 79oC

7. Melting point = 1455oC

8. Boiling point = 1413oC

Solubility: Use the ref. packet to identify if the substance is soluble or insoluble.

9. Lithium sulfate

10. Strontium oxide

11. Lead (IV) bromide

12. Ammonium carbonate

Identify each of the following as an element, a compound, a homogeneous mixture or a heterogeneous mixture.

13. Water

14. Cheerios in milk

15. Apple juice

16. Silver

17. Salsa

18. A bag of nuts and bolts

Identify each of the following as a chemical or physical property

19. Combustible

20. Mass

21. Volume

22. Ability to rust

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Identify each of the following as a chemical or physical change

23. Melts 25. Dissolves 27. Tarnishis

24. Burns 26. Rips 28. Shatters

25. Determine the following for the fluorine atom.

a) number of protons b) number of neutrons c) number of electrons

d) atomic number e) mass number

26. If an element has an atomic number of 34 and a mass number of 78, what is the

a) number of protons b) number of neutrons

c) number of electrons d) element

27. If an element has 91 protons and 140 neutrons, what is the

a) atomic number b) mass number

c) number of electrons d) element

28. If an element has 78 electrons and 117 neutrons what is the

a) atomic number b) mass number

c) number of protons d) element

Density Practice: Solve each problem below, writing the equation and showing the substitution. Provide a unit for

each answer.

1. A block of aluminum occupies a volume of 15.0 mL and weighs 40.5 g. What is its density?

2. Find the mass of gold that occupies 965 cm3 of space.

3. Mercury metal is poured into a graduated cylinder that holds exactly 22.5 mL. The mercury used to fill the

cylinder weighs 306.0 g. From this information, calculate the density of mercury.

4. Find the volume occupied by 250.0 g of O2.

5. A cube of metal has a side length of 1.55 cm. If the sample is found to have a mass of 26.7 g, find the density

and identity of the metal.

6. An irregularly-shaped sample of aluminum (Al) is put on a balance and found to have a mass of 43.6 g. The

student decides to use the water-displacement method to find the volume. The initial volume reading is

25.5 mL and, after the Al sample is added, the water level has risen to 41.7 mL. Find the density of the Al

sample in g/cm3. (Remember: 1 mL = 1 cm

3.)

7. A flask that weighs 345.8 g is filled with 225 mL of carbon tetrachloride. The weight of the flask and carbon

tetrachloride is found to be 703.55 g. From this information, calculate the density of carbon tetrachloride.

The Study of Matter Practice Test

Directions: Define and/or describe the following terms relating to the scientific method.

1. Dependent Variable _________________________________________________

2. Hypothesis __________________________________________________________

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Directions: Indicate if the process listed is a physical or chemical change.

3. Food digests _________________________________________________________

4. Bending a piece of copper wire. _______________________________________

5. Two clear liquids react to form a yellow clumps ___________________________

Directions: Solve the following problems. Show all work! Be sure to include the correct unit with your final

answer.

6. What is the density of a substance that has a volume of 2.8 cm3 and mass of 25 grams?

7. What is the density of a solid that has a volume of 4 cm3 and a mass of 6 grams?

Directions: For each sample of matter below, correctly classify it as a pure substance or a mixture.

8. Trail Mix ____________________

9. Helium ____________________

Directions: For each, correctly classify as homogeneous or heterogeneous mixture.

10. Vegetable soup ______________________

11. Gatorade ______________________

12. Orange juice, with pulp ______________________

13. The amount of mass per unit volume refers to the

a. Density

b. Specific weight

c. Volume

d. Weight

14. A substance that can be further simplified may be either

a. An element or a compound

b. An element or a mixture

c. A mixture or a compound

d. A mixture or an atom

15. A substance composed of two or more elements chemically united is called

a. An isotope

b. A compound

c. An element

d. A mixture

16. An example of a chemical change is the

a. Breaking of a glass bottle

b. Sawing of a piece of wood

c. Rusting of iron

d. Melting of an ice cube

17. A substance that cannot be further decomposed by ordinary chemical means is

a. Water

b. Air

c. Sugar

d. Silver

18. An example of a physical change is

a. The fermenting of sugar to alcohol

b. The rusting of iron

c. The burning of paper

d. A solution of sugar in water

19. What Kelvin temperature is equal to 25°C?

a. 248 K

b. 298 K

c. 100 K

d. 200 K

20. Which substance cannot be decomposed into simpler substances?

a. ammonia

b. aluminum

c. methane

d. methanol

Unit 2- Atomic Theory and Structure

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21. A compound differs from a mixture in that a compound always has a

a. homogeneous composition

b. maximum of two components

c. minimum of three components

d. heterogeneous composition

22. Which statement describes a chemical property?

a. Its crystals are a metallic gray.

b. It dissolves in alcohol.

c. It is a violet-colored gas.

d. It reacts with hydrogen.

23. An experiment for a new asthma medication was set up into two groups. Group one was given the new drug for

asthma, while group 2 was given a sugar pill. The sugar pill serves as

a. Control

b. Constant

c. experimental variable

d. dependent variable

24. To determine the density of an irregularly shaped object, a student immersed the object in 21.2 milliliters of

H2O in a graduated cylinder, causing the level of the H2O to rise to 27.8 milliliters. If the object had a mass of

22.4 grams, what was the density of the object.

a. 27.8 g / mL

b. 6.6 g / mL

c. 3.0 g / mL

d. 3.4 g/ mL

25. A scientist plants two rows of corn for experimentation. She puts fertilizer on row 1 but does not put fertilizer

on row 2. Both rows receive the same amount of water and light intensity. She checks the growth of the corn

over the course of 5 months. What is a constant in this experiment?

a. Plant height

b. Corn without fertilizer

c. Corn with fertilizer

d. Amount of water

26. The measurable factor in an experiment is known as the:

a. Control

b. independent variable

c. constant

d. dependent variable


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ATOMIC THEORY

HISTORY OF THE ATOM

The original idea (400 B.C.) came from ______________________, a Greek philosopher. He expressed the belief

that all matter is composed of very small, indivisible particles, which he named atomos. John Dalton (1766-1844),

an English school teacher and chemist, proposed his atomic theory of matter in 1803. Dalton’s Atomic Theory

states that:

1. All matter is made of tiny __________________________ particles called atoms.

2. Atoms of the ____________ element are identical; those of different elements are different.

3. Atoms of different elements combine in whole number ________________ to form compounds

4. Chemical reactions involve the rearrangement of atoms. No _______ atoms are created or destroyed.

PARTS OF THE ATOM

Because of Dalton’s atomic theory, most scientists in the 1800s believed that the atom was like a tiny solid ball that

could not be broken up into parts. In 1897, a British physicist, J.J. Thomson, discovered that this solid-ball model

was not accurate. Thomson’s experiments used a __________________ ray tube. It is a vacuum tube - all the air

has been pumped out. Because these rays originate at the ____________________, they are called cathode rays.

Thomson concluded that cathode rays are made up of invisible, _________________________ charged particles

referred to as electrons. From Thomson’s experiments, scientists had to conclude that atoms were not just neutral

_________________, but somehow were composed of electrically charged particles. Matter is not negatively

charged, so atoms can’t be negatively charged either. If atoms contained extremely light, negatively charged

particles, then they must also contain positively charged particles — probably with a much greater _____________

than electrons. J.J. Thomson said the atom was like ______________ pudding, a popular English dessert.

In 1909, a team of scientists led by Ernest Rutherford in England carried out the first of several important

experiments that revealed an arrangement far different from the plum pudding model of the atom. The

experimenters set up a lead-shielded box containing radioactive polonium, which emitted a beam of positively

charged subatomic particles through a small hole. The sheet of ________________ foil was surrounded by a screen

coated with zinc sulfide, which glows when struck by the positively charged particles of the beam. The

________________ particles were expected to pass through without changing direction very much because

Rutherford thought the mass was evenly distributed in the atom. Because most of the particles passed through the

foil, they concluded that the atom is nearly all _______________ ______________. Because so few particles were

deflected, they proposed that the atom has a small, dense, positively charged central core, called a

____________________. Alpha particles are deflected by it if they get close enough to the nucleus.

In 1910, J.J. Thomson discovered that neon consisted of atoms of two different masses. Atoms of an

element that are chemically alike but differ in mass are called ______________________ of the element. Because

of the discovery of isotopes, scientists hypothesized that atoms contained still a third type of particle that explained

these differences in mass. Calculations showed that such a particle should have a mass ____________________ to


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that of a proton but no electrical _______________. The existence of this neutral particle, called a neutron, was

confirmed in the early 1930s. James _________________ is given credit for discovering the neutron.

NAME SYMBOL CHARGE RELATIVE MASS

1/2000

proton

no

AVERAGE ATOMIC MASS

The atomic mass is the weighted average mass of all the naturally occurring isotopes of that element.

To determine the average atomic mass, first calculate the contribution of each isotope to the average atomic mass,

being sure to convert each ___ to a fractional abundance. The average atomic mass of the element is the sum of the

mass contributions of each isotope.

Elements can be represented by using the symbol of the element, the mass number and the atomic number. The

mass number is the __________________ mass rounded to a whole number.

Practice/Homework

Isotopes and Subatomic Particles

Complete the chart

Isotope symbol Atomic # Mass Protons Neutrons Electrons

8 9

77 54

Pb207

82

12 13

38 50

U238

92

As75

33

S32

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MOLES

We measure ________________ in grams. We measure volume in __________________. We count pieces in

_________________. The number of moles is defined as the number of __________________ atoms in exactly

____ grams of carbon-12. ____ mole is 6.022 x 1023

particles. 6.022 x 1023

is called __________________ number.

Representative particles are the smallest pieces of a substance. For a molecular compound it is a(n)

______________________. For an ionic compound it is a ______________________ ______________. For an

element it is a(n) ________________.


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How many oxygen atoms are in the following?

a) CaCO3 b) Al2(SO4)3

How many total ions are in the following?

a) CaCl2 b) NaF c) Al2S3

MOLE CONVERSIONS

1. How many atoms of carbon are there in 1.23 moles of carbon?

2. How many molecules of CO2 are in 4.56 moles of CO2?

3. How many atoms of iron are in 0.600 moles of iron?

4. How many moles are in 7.78 x 1024

formula units of MgCl2?

5. How many moles of water are 5.87 x 1022

molecules of water?

6. How many moles of aluminum are 1.2 x 1024

atoms of aluminum?

Calculate the number of particles (atoms, ions or molecules) in each of the following.

a) 3.4 moles Na2S b) 0.0020 moles Zn

c) 1.77 x 10-11

moles C d) 92.35 moles O2

Calculate the number of moles in each of the following.

a) 3.4 x 1024

molecules HCl b) 8.7 x 1021

atoms Zn

c) 1.77 x 1018

ions Al+3

d) 2.66 x 1026

atoms Cu

MOLAR MASS

Molar mass is the generic term for the mass of one _____________. It may also be referred to as gram molecular

mass, gram formula mass, and gram atomic mass. The unit is ______________. To determine the molar mass of an

element, find the element’s symbol on the periodic table and round the mass so there is __________ digit beyond

the decimal.

Determine the molar mass of the each of the following elements.

a) sulfur (S) b) chromium (Cr) c) bromine (Br)

To determine the molar mass of a compound, find the mass of all elements in the compound.

If necessary, ___________________ an element’s mass by the subscript appearing beside that element in the

compound’s formula (or ________________ of the subscripts).

Calculate the molar mass of each of the following compounds.

a) Na2S b) N2O4 c) C6H12O6 d) Ca(NO3)2


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MASS-PARTICLE/MOLE CONVERSIONS

1. How many atoms of lithium are in 1.00 g of Li?

2. How many molecules of sodium oxide are in 42.0 g of Na2O?

3. How much would 3.45 x 1022

atoms of uranium (U) weigh?

4. How many moles of magnesium are in 56.3 g of Mg?

5. How many moles is 5.69 g of NaOH?

6. How many grams of sodium chloride are in 3.45 moles of NaCl?

7. How many moles is 4.8 g of CO2?

8. How many grams is 9.87 moles of H2O?

9. How many molecules are in 6.8 g of CH4?

10. What is the mass of 49.0 molecules of C6H12O6?

GASES

Many of the chemicals we deal with are gases. They are difficult to weigh, and we need to know how many moles

of gas we have. Two things affect the volume of a gas: temperature and pressure. Standard temperature is

______ ºC, and standard pressure is ______ atm. Standard temperature and pressure is abbreviated STP. At STP 1

mole of gas occupies ______ L. 22.4 L is called the _____________ volume. Avogadro’s Hypothesis - At the same

temperature and pressure equal volumes of gas have the same number of _______________________.

GAS CONVERSIONS

1. What is the volume of 4.59 mole of CO2at STP?

2. How many moles is 5.67 L of O2 at STP?

3. What is the volume of 8.8 g of CH4 gas at STP?

4. How many grams is 16.2 L of O2 at STP?

Calculate the number of liters in each of the following.

a) 3.10 x 1024

molecules Cl2 b) 8.7 moles Ne

c) 2.77 x 1018

atoms He d) 266 grams SO2

Homework/Practice

Part 1--Convert between particles and moles

1. 24 atoms of sodium = _____ moles of sodium atoms

2. 5 molecules of chlorine gas = _____ moles of chlorine molecules

3. 900 atoms of silver = _____ moles of silver atoms

4. 2.89 x 1023

molecules of ammonia = _____ moles of ammonia molecules

5. 15 moles of arsenic atoms = ______ atoms of arsenic

6. 4.00 x 103 moles of barium atoms = __________ atoms of barium


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Part 2--Convert between mass and moles

7. Calculate the mass of 1.000 mole of CaCl2

8. Calculate grams in 3.0000 moles of CO2

9. Calculate number of moles in 32.0 g of CH4

10. Calculate moles in 168.0 g of HgS

11. Calculate moles in 510.0 g of Al2S3

12. How many moles are in 27.00 g of H2O

13. What is the mass of 2.55 moles Cu2CrO4

Part 3- Multiple steps

14. Arrange the following in order of increasing weight.

a. 10.4 g of sulfur

b. 0.179 moles of iron

c. 6.33 x 1025

atoms of hydrogen

d. 0.77 moles of N2

15. How many grams would 8.1 × 1021

molecules of sucrose (C12H22O11) weigh?

16. How many atoms are in a 2.0 kg ingot of gold? (Note mass units.)

17. What is the mass of 2.3 x 1024

molecules of KCl?

18. Calculate the number of molecules in 50.0 grams of H2SO4

19. Calculate the number of molecules in 100. grams of KClO4

20. Calculate the number of molecules in 8.76 grams of NaOH

21. Calculate the mass of 1.2 x 1022

molecules of Fe3(PO4)2

22. Calculate mass of 7.2 x 1024

molecules of Na2CO3

NUCLEAR CHEMISTRY

Nuclear chemistry is the study of the structure of _________________ nuclei and the changes they undergo. Marie

Curie named the process by which materials such as uranium give off rays radioactivity; the rays and particles

emitted by a radioactive source are called __________________. As you may recall, isotopes are atoms of the same

element that have different numbers of _________________. Isotopes of atoms with unstable nuclei are called

______________________. These unstable nuclei emit radiation to attain more stable atomic configurations in a

process called radioactive ________________. During radioactive decay, unstable atoms lose _________________

by emitting one of several types of radiation.

TYPES OF RADIATION

The three most common types of radiation are alpha (α), ____________ (β), and gamma (γ). An alpha particle (α)

has the same composition as a __________________ nucleus - two protons and ________ neutrons - and is

therefore given the symbol _________. The charge of an alpha particle is 2+ due to the presence of the two

___________________. Because of their mass and charge, alpha particles are relatively slow-moving compared

with other types of radiation. Thus, alpha particles are not very ________________________ - a single sheet of


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paper stops alpha particles. A beta particle is a very-fast moving ______________________ that has been emitted

from a neutron of an unstable nucleus. Beta particles are represented by the symbol _________. The zero

superscript indicates the insignificant mass of an electron in comparison with the mass of a

____________________. The –1 subscript denotes the _____________________ charge of the particle. Beta

radiation consists of a stream of fast-moving electrons. Because beta particles are both lightweight and fast moving,

they have _____________________ penetrating power than alpha particles. A thin metal foil is required to stop

beta particles. Gamma rays are high-energy (_________________ wavelength) electromagnetic radiation. They are

denoted by the symbol __________. As you can see from the symbol, both the subscript and superscript are zero.

Thus, the emission of gamma rays does not change the __________________ number or mass number of a nucleus.

Gamma rays almost always accompany alpha and beta radiation, as they account for most of the energy loss that

occurs as a nucleus decays.

NAME SYMBOL FORMULA MASS CHARGE DESCRIPTION

Alpha He4

2

β -1

0 High energy

radiation

NUCLEAR STABILITY and DECAY

Radioactive nuclei undergo decay in order to gain _____________________. All elements with atomic numbers

greater than 83 are radioactive. Nuclear equations are used to show nuclear transformations. Balanced nuclear

equations require that both the ____________________ number and the mass number must be balanced.

• When beryllium-9 is bombarded with alpha particles (helium nuclei), a neutron is produced. The balanced

nuclear reaction is given as: ________________________________________________

The atomic number (the number on the bottom) determines the identity of the element.

• When nitrogen-14 is bombarded with a neutron, a proton is produced. The balanced nuclear equation can be

written as: _________________________________________________________

• Polonium-230 undergoes alpha decay: ________________________________________________

• Uranium-234 undergoes alpha decay: ________________________________________________

• Cobalt-50 undergoes beta decay: ____________________________________________________

Provide symbols for each of the following: neutron ___________, proton ___________ or ___________, and the

positron ___________.


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• What element is formed when iron-60 undergoes beta decay? Give the atomic number and mass number of the

element. ____________

• Write a balanced nuclear equation for the alpha decay of the following radioisotope, uranium-235.

____________________________________________________________

• Nitrogen-12 decays into a positron and another element. Write the balanced nuclear equation.

____________________________________________________________

• Uranium-238 is bombarded with a neutron. One product forms along with gamma radiation. Write the

balanced nuclear equation.

____________________________________________________________

• Nitrogen-14 is bombarded with deuterium (hydrogen-2). One product forms along with an alpha particle.

Write the balanced nuclear equation.

____________________________________________________________

RADIOACTIVE DECAY RATES

Radioactive decay rates are measured in half-lives. A half-life is the time required for one-half of a radioisotope’s

nuclei to ________________ into its products. For example, the half-life of the radioisotope strontium-90 is 29

years. If you had 10.0 g of strontium-90 today, 29 years from now you would have 5.0 g left. The decay continues

until negligible strontium-90 remains.

• The half-life of iron-59 is 44.5 days. How much of a 2.000-mg sample will remain after 133.5 days?

• Cobalt-60 has a half-life of 5.27 years. How much of a 10.0 g sample will remain after 21.08 years?

• Carbon-14 has a half-life of 5730 years. How much of a 250. g sample will remain after 5730 years?

FISSION and FUSION

Heavy atoms (mass number > 60) tend to break into smaller atoms, thereby increasing their

________________________. Using a neutron to split a nucleus into fragments is called nuclear

_______________________. Nuclear fission releases a large amount of energy and several neutrons. Since

neutrons are products, one fission reaction can lead to more fission reactions, a process called a ________________

reaction. A chain reaction can occur only if the starting material has enough mass to sustain a chain reaction; this

amount is called __________________ mass. The _____________________ of atomic nuclei is called nuclear

fusion. For example, nuclear fusion occurs within the Sun, where hydrogen atoms fuse to form

__________________ atoms. Fusion reactions can release very large amounts of energy but require extremely high

temperatures. For this reason, they are also called _____________________________ reactions.

EFFECTS OF NUCLEAR REACTIONS

Any exposure to radiation can damage living ____________. Gamma rays are very dangerous because they

penetrate ______________________ and produce unstable and reactive molecules, which can then disrupt the


19

normal functioning of cells. The amount of radiation the body absorbs (a dose) is measured in units called rads and

____________. Everyone is exposed to radiation, on average 100–300 millirems per year. A dose exceeding

____________ rem can be fatal.

Atomic Theory, The Mole, and Nuclear Chemistry- Practice Test

1. Given the work of Dalton, please check the box for the postulate(s) that have since been proven to be incorrect.

Explain what we now know to be the true case.

[] All atoms of a specific element are identical.

[] Compounds consist of atoms of different elements combined together.

[] Atoms of different elements have different masses.

Directions: For the scientist listed below, explain what was done in the experiment, what knowledge was

developed as a result.

2. Rutherford

3. Thomson

Directions: Fill in the table for the following isotopes.

Isotope Atomic # Mass # Protons Neutrons Electrons

4. H-1

5. Cu-65

6. 18 40

7. 19 9

8. What is the charge of a beta particle?

Directions: Solve the following problems be sure to include the correct unit with your final answer.

9. Given the equation: X � He4

2 + Po220

84 The nucleus represented by X is

10. How many moles of sodium are 6.02 x 10 23

atoms of sodium?

11. What is the mass of 6 moles of Carbon?

12. How many atoms are in 45 g of Neon?

Multiple Choice Practice

13. What is the approximate formula mass of Ca(NO3)2

a. 70

b. 82

c. 102

d. 150

e. 164

14. How many molecules are in 1 mole of water?

a. 3

b. 54

c. 6.02 x 1023

d. 2 (6.02 x 1023

)

e. 3 (6.02 x 1023

)

15. How many atoms are represented in the formula Ca3(PO4)2

a. 5

b. 8

c. 9

d. 12

e. 13


20

16. What is the mass in grams of 1 mole of KAl(SO4)2•12H2O

a. 132

b. 180

c. 394

d. 474

e. 516

17. Compared to the charge and mass of a proton, an electron has

a. the same charge and a smaller mass

b. the same charge and the same mass

c. an opposite charge and a smaller mass

d. an opposite charge and the same mass

18. When alpha particles are used to bombard gold foil, most of the alpha particles pass through undeflected.

This result indicates that most of the volume of a gold atom consists of ____.

a. deuterons

b. neutrons

c. protons

d. unoccupied space

19. A proton has approximately the same mass as

a. a neutron

b. an alpha particle

c. a beta particle

d. an electron

20. A neutron has approximately the same mass as a

a. an alpha particle

b. a beta particle

c. an electron

d. a proton

21. Which symbols represent atoms that are isotopes?

a. C-14 and N-14

b. O-16 and O-18

c. I-131 and I-131

d. Rn-222 and Ra-222

22. Which atom contains exactly 15 protons?

a. P-32

b. S-32

c. O-15

d. N-15

23. An ion with 5 protons, 6 neutrons, and a charge of 3+ has an atomic number of

a. 5 b. 6 c. 8 d. 11

24. What is the mass number of an atom which contains 28 protons, 28 electrons, and 34 neutrons?

a. 56 b. 62 c. 90 d. 28

25. What is the gram formula mass of K2CO3?

a. 138 g

b. 106 g

c. 99 g

d. 67 g

26. What is the total number of atoms contained in 2.00 moles of nickel?

a. 58.9

b. 118

c. 6.02 x 1023

d. 1.2 x 1024

27. What is the mass in grams of 3.0 x 1023

molecules of CO2?

a. 22 g

b. 44 g

c. 66 g

d. 88 g

28. The amount of substance having 6.022 x 1023

of any kind of chemical unit is called a(n):

a. formula

b. mass number

c. mole

d. atomic weight

29. The total number of atoms in a formula unit of aluminum dichromate, Al2(Cr2O7)3 is:

a. 5 b. 29 c. 17 d. 11


21

30. The formula mass of calcium hydroxide, Ca(OH)2 is:

a. 57.05 grams

b. 74.10 grams

c. 128 grams

d. 97.07 grams

31. What is the molar mass of the gas butane, C4H10?

a. 13.02 grams

b. 485.2 grams

c. 68 24 grams

d. 58.14 grams

32. The formula mass of magnesium hydroxide, Mg(OH)2 is:

a. 42.33 grams

b. 58.33 grams

c. 41.32 grams

d. 5 grams

33. What is the mass in grams of 3 moles of water molecules, H2O?

a. 54.06 grams

b. 21.02 grams

c. 0.166 grams

d. 6.01 grams

34. What is the mass in grams of 10 moles of ammonia, NH3?

a. 170.4 grams

b. 0.587 grams

c. 1.704 grams

d. 27.04 grams

35. How many moles of water molecules, H2O, are present in a 42 gram sample of water?

a. 23.98 moles

b. 0.429 moles

c. 2.33 moles

d. 757 moles

36. How many moles of methane molecules, CH4, are in 80 grams of methane?

a. 0.201 moles

b. 4.98 moles

c. 6.022 x 1080

moles

d. 1284 moles

37. How many moles of calcium hydroxide, Ca(OH)2 are in 150 grams of the compound?

a. 2.02 moles

b. 224.1 moles

c. 0.494 moles

d. 11115 moles

38. How many oxygen atoms are there in one formula unit of Al2(SO4)3?

a. 3 b. 4 c. 7 d. 12

39. Which of the following arrangements represent different isotopes of the same element?

i. 12 protons, 11 neutrons, 12 electrons

ii. 11 protons, 12 neutrons, 11 electrons

iii. 10 protons, 12 neutrons, 12 electrons

iv. 11 protons, 12 neutrons, 10 electrons

v. 12 protons, 12 neutrons, 12 electrons

a. 1 and 5

b. 2 and 4

c. 2, 3, 4 and 5

d. all of these qualify

e. None of these qualify

40. If the abundance of 6Li (6.015121 amu) is 7.500% and the abundance of

7Li (7.016003 amu) is 92.500%,

what is the average atomic mass?

a. 6.0750 amu

b. 6.0902 amu

c. 6.9250 amu

d. 6.9409 amu

41. An alpha (α) particle is essentially a ____________________ nucleus.

a. plutonium

b. helium

c. hydrogen

d. uranium

e. carbon-12


22

42. Which of the following have equal numbers of neutrons?

a. I, II and III

b. II and III

c. I and V

d. I and IV

e. II, III and IV

43. The element hafnium (Hf) has five stable isotopes. The correct number of nuclear particles in an atom of

hafnium-178 is:

a. 72 protons, 178 neutrons

b. 72 protons, 72 electrons

c. 106 protons, 72 neutrons

d. 72 protons, 106 neutrons

e. 72 protons, 106 neutrons,

72 electrons

44. J.J. Thomson's model of the atom can be summarized with the visual image of:

a. planets orbiting the sun

b. plum pudding

c. bees around a hive

d. a small central nucleus and an

electron cloud

e. none of the above

45. Identify the missing particle in the following nuclear reaction: K37

19 → _____ + e0

1+

a. Ar37

18

b. Ar38

18

c. Ar36

18

d. Ca37

20

46. For the most common types of radioactive decay, the order of least penetrating to human tissue, to most

penetrating to human tissue is:

a. gamma, beta, alpha

b. alpha, beta, gamma

c. beta, gamma, alpha

d. gamma, alpha, beta

47. Phosphorus-15 has a half-life of 14 days. What proportion of the original phosphorus-15 remains after

8 weeks?

a. 1/2

b. 1/16

c. 1/4

d. 1/32

e. 1/8

48. The nuclide radium-226 is the daughter nuclide resulting from the α decay of what parent nuclide?

a. radon-222

b. polonium-214

c. thorium-230

d. thorium-228

e. radium-225

49. An electron emitted from the nucleus during some kinds of radioactive decay is known as:

a. A gamma ray

b. A positron

c. A beta (β) particle

d. An alpha (α) particle

50. A process in which a very heavy nucleus splits into more-stable nuclei of intermediate mass is called:

a. nuclear fission

b. radiocarbon dating

c. a chain reaction

d. nuclear fusion

Unit 3- Electrons and Periodicity

23

PERIODIC TABLE

LIGHT

Light is a kind of electromagnetic _____________________. All forms of electromagnetic radiation move at

3.00 x 108 m/s. The ______________ is the baseline of a wave. The crest is the high point on a wave, and the

trough is low point on a wave. The amplitude of a wave is the wave’s _____________ from the origin to a crest, or

from the origin to a trough. Wavelength (represented by λ, the Greek letter lambda) is the ___________________

distance between equivalent points on a continuous wave. Wavelength is the distance from crest to crest or trough

to trough and is usually expressed in meters (m). _____________________ (represented by f ) is the number of

“waves” that pass a given point per second, and the units are cycles/sec or hertz (Hz)

c = f λ c = the speed of light

Frequency and wavelength are __________________ related, which means that as one goes up the other goes

_____________. Different frequencies of light correspond to different colors of light. In 1900, the German

physicist Max Planck began searching for an explanation as he studied the light emitted from

___________________ objects. Matter can gain or lose energy only in small, specific amounts called

_______________. That is, a quantum is the minimum amount of energy that can be gained or lost by a(n)

____________. That is, while a beam of light has many wavelike characteristics, it also can be thought of as a

stream of tiny particles, or bundles of energy, called ________________. Thus, a photon is a particle of

electromagnetic radiation with no _____________ that carries a quantum of energy. Planck went further and

demonstrated mathematically that the energy of a quantum is ___________________ related to the frequency of the

emitted radiation.

Scientists knew that the wave model of light could not explain a phenomenon called the ____________________

effect. In the photoelectric effect, electrons, called __________________________, are emitted from a metal’s

surface when light of a certain _______________________ shines on the surface. Einstein proposed that for the

photoelectric effect to occur, a photon must possess, at a minimum, the energy required to _______________ an

electron from an atom of the metal.

Building on Planck’s and Einstein’s concepts of ____________________ energy (quantized means that only certain

values are allowed), Bohr proposed that the hydrogen atom has only certain allowable energy ______________.

The lowest allowable energy state of an atom is called its _______________ state. When an atom gains energy, it is

said to be in a(n) __________________ state. When the atom is in an excited state, the electron can drop from the

higher-energy orbit to a _______________-energy orbit. As a result of this transition, the atom emits a

____________________ corresponding to the difference between the energy levels associated with the two orbits.

ATOMIC EMISSION SPECTRA

By heating a gas of a given element with electricity, we can get it to give off _______________. Each element

gives off its own characteristic colors. The spectrum can be used to __________________ the atom. These are

called line _______________. Each is unique to an element. The spectrum of light released from excited atoms of


24

an element is called the _________________ spectrum of that element. As the electrons fall from the excited state,

they __________________ energy in the form of light. The further they fall, the ________________ the energy.

This results in a higher frequency.

Use the Chemistry Reference Tables to answer the following:

(a) An electron falls from energy level 5 to 3. What is the wavelength of the light emitted?

(b) An electron falls from energy level 6 to 2. What is the wavelength of the light emitted?

(c) An electron falls from energy level 3 to 1. What type of electromagnetic radiation is emitted (infrared,

visible or ultraviolet)?

(d) An electron falls from energy level 4 to 2. What type of electromagnetic radiation is emitted (infrared,

visible or ultraviolet)?

(e) An electron falls from energy level 5 to 2. What color of visible light is emitted?

(f) An electron falls from energy level 3 to 2. What color of visible light is emitted?

PERIODIC TABLE- HISTORY

The Russian chemist, Dmitri ______________________ was studying the properties of the elements and realized

that the chemical and physical properties of the elements repeated in an orderly way when he organized the elements

according to increasing atomic ___________. Mendeleev later developed an improved version of his table with the

elements arranged in horizontal ___________. This arrangement was the forerunner of today’s periodic table.

Patterns of changing properties repeated for the elements across the horizontal rows. Elements in vertical

___________________ showed similar properties. Mendeleev grouped elements in columns by similar properties

in order of increasing atomic mass. He found some inconsistencies and felt that the properties were more important

than the mass, so he switched order. Mendeleev left some _____________ in his periodic table, deciding there must

be undiscovered elements. He predicted their properties before they were found. Mendeleev is considered to be the

_________________ of the Periodic table. This repeated pattern (when Mendeleev grouped elements in columns by

similar properties) is an example of __________________ in the properties of elements. Periodicity is the tendency

to recur at regular intervals. By 1860, scientists had already discovered _________ elements and determined their

atomic masses.

THE MODERN PERIODIC TABLE

Fifty years after Mendeleev, the British scientist Henry ________________ discovered that the number of protons in

the nucleus of a particular type of atom was always the same. When atoms were arranged according to increasing

atomic ___________________, the few problems with Mendeleev's periodic table disappeared. Because of

Moseley's work, the modern periodic table is based on the atomic numbers of the elements. The statement that the

physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of

increasing atomic number is known as the periodic _____________. On the periodic table a _________________,

sometimes also called a series, consists of the elements in a horizontal row. A __________________, sometimes


25

also called a family, consists of the elements in a vertical column. Elements are placed in columns by similar

properties.

The elements in the A groups are called the __________________ elements. The B groups are called the

____________________ elements. The two rows at the bottom of the table are called the inner transition elements.

Group 1A elements are the _________________ metals. Group 1A elements have ______ valence electron and form

_______ ions after losing the one valence electron. Group 2A elements are the alkaline earth metals. Group 2A

elements have ______ valence electrons and form 2+ ions after losing the two __________________ electrons.

Group 3A is called the _________________ group. Group 3A elements have ________ valence electrons and form

3+ ions after losing the three valence electrons. Group _______ is called the carbon group. Group 4A elements

have four valence electrons and form 4+ ions after ___________________ the four valence electrons or 4- ions after

___________________ four additional electrons. Group 5A is called the _____________________ group. Group

5A elements have five valence electrons and form ________ ions after gaining three more electrons. Group 6A is

called the oxygen group. Group 6A elements have _______ valence electrons and form 2- ions after

____________________ two more electrons. Group 7A is called the ____________________. Group 7A elements

have seven valence electrons and form 1- ions after gaining one more electron. The word halogen is from the Greek

words for “______________ former” so named because the compounds that halogens form with metals are salt-like.

Group 8A elements are the ________________ gases. Group 8A elements have eight valence electrons except for

helium which only has ________. The noble gases, with a full complement of valence electrons, are generally

unreactive. All transition elements have _______ valence electrons.

• How many valence electrons are in an atom of each of the following elements?

a) Magnesium (Mg) ______ b) Selenium (Se) ______ c) Tin (Sn) _____

METALS, NONMETALS AND METALLOIDS

Metals are elements that have ________________, conduct ____________ and electricity, and usually bend without

breaking. Most metals have one, two, or three valence electrons. All metals except _________________ are solids

at room temperature; in fact, most have extremely _____________ melting points. A metal’s

___________________ is its ability to react with another substance.

• Consult the “Activity Series of Metals” in the Chemistry Reference Tables to determine the more active metal.

a) cobalt (Co) or manganese (Mn) ________ b) barium (Ba) or sodium (Na) ________

Although the majority of the elements in the periodic table are _________________, many nonmetals are abundant

in nature. Most nonmetals don’t conduct electricity, are much poorer conductors of heat than metals, and are

__________________ when solid. Many are ______________ at room temperature; those that are solids lack the

luster of metals. Their _____________________ points tend to be lower than those of metals. With the exception

of carbon, nonmetals have five, six, seven, or eight valence electrons. A nonmetal’s reactivity is its ability to react

with another substance.


26

• Consult the “Activity Series of Metals” in the Chemistry Reference Tables to determine the less active

nonmetal.

a) fluorine (F2) or chlorine (Cl2) _________ b) chlorine (Cl2) or iodine (I2) _________

__________________ have some chemical and physical properties of metals and other properties of nonmetals. In

the periodic table, the metalloids lie along the border between metals and nonmetals. Some metalloids such as

silicon, germanium (Ge), and arsenic (As) are _____________________. A semiconductor is an element that does

not conduct electricity as well as a ________________, but does conduct slightly better than a nonmetal.

ELECTRONS IN ATOMS

THE BOHR MODEL OF THE ATOM

Niels Bohr, a young Danish physicist working in Rutherford’s laboratory in 1913, suggested that the single electron

in a ___________________ atom moves around the nucleus in only certain allowed circular orbits. The atom

looked like a miniature _________________ system. The nucleus is represented by the sun, and the electrons act

like the planets. The orbits are circular and are at different levels. Amounts of ___________________ separate one

level from another. (Modern View: The atom has two regions and is 3-dimensional. The nucleus is at the

_________________ and contains the protons and neutrons. The electron _________________ is the region where

you might find an electron and most of the volume of an atom.) Bohr proposed that electrons must have enough

energy to keep them in constant motion around the ___________________. Electrons have energy of motion that

enables them to overcome the attraction of the _________________ nucleus. Further away from the nucleus means

more energy. Electrons reside in ________________ levels.

QUANTUM MECHANICAL MODEL

Like Bohr’s model, the quantum mechanical model limits an electron’s energy to certain values. The space around

the nucleus of an atom where the atom’s electrons are found is called the electron ________________. A three-

dimensional region around the nucleus called an atomic __________________ describes the electron’s probable

location. In general, electrons reside in principal ________________ levels. As the energy level number increases,

the orbital becomes _______________, the electron spends more time ___________________ from the nucleus, and

the atom’s energy level increases. Principal energy levels contain energy ___________________. Principal energy

level 1 consists of a single sublevel, principal energy level 2 consists of __________ sublevels, principal energy

level 3 consists of three sublevels, and so on. Sublevels are labeled s, p, d, or f. The s sublevel can hold 2 electrons,

the p sublevel can hold _____ electrons, the d sublevel can hold 10 electrons, and the f sublevel can hold 14

electrons. Sublevels contain __________________. Each orbital may contain at most ________ electrons. There is

one s orbital for every energy level, and the s orbital is ____________________ shaped. They are called the 1s, 2s,

3s, etc… orbitals. The p orbitals start at the second energy level, reside along ______ different directions and have 3

different ________________ shapes. The d orbitals start at the ________________ energy level and have ____

different shapes. The f orbitals start at the fourth energy level and have ______ different shapes.


27

ELECTRON CONFIGURATIONS

Electron configurations represent the way electrons are arranged in atoms. The Aufbau principle states that

electrons enter the __________________ energy first. This causes difficulties because of the ________________ of

orbitals of different energies. At most there can be only 2 electrons per orbital, and they must have

__________________ “spins.” Hund’s rule states that when electrons occupy orbitals of equal energy, they don’t

_________ up with an electron of opposite spin until they have to.

Let’s determine the electron configuration for phosphorus. ______________________________

Let’s determine the electron configuration for chromium. _______________________________

• Write the electron configuration for aluminum (Al). ________________________________

• Write the electron configuration for neon (Ne). ____________________________________

• Write the electron configuration for calcium (Ca). __________________________________

• Write the electron configuration for iron (Fe). _____________________________________

• Write the electron configuration for bromine (Br). _________________________________

To identify an element with a given electron configuration, add the _________________ numbers together and find

the element with that atomic number.

Directions: Identify the element with the following electron configuration:

a. 1s2 2s

2 2p

6 3s

2 3p

4 _________________________________

b. 1s2 2s

2 2p

6 3s

2 3p

6 4s

2 3d

9 _________________________________

c. 1s2 2s

2 2p

6 3s

2 3p

6 4s

2 3d

10 4p

2 _________________________________

Electron Configuration Using a Noble Gas Abbreviation - In order to write this type of configuration, find the

_______________ gas (from Group 8A) that comes before the element in question. Put the symbol for the noble

gas in _____________________ and then write the part of the configuration that follows to reach the desired

element.

Write the electron configuration using a noble gas abbreviation for:

• magnesium (Mg) _________________________ • nickel (Ni) ___________________

• fluorine (F) _________________________ • silicon (Si) ___________________

• zirconium (Zr) _________________________

VALENCE ELECTRONS

The electrons in the ______________________ energy level are called valence electrons. You can also use the

periodic table as a tool to predict the number of valence electrons in any atom in Groups 1, 2, 13, 14, 15, 16, 17, and

18. All atoms in Group 1, like hydrogen, have __________ valence electron. All atoms in Group 2 have two, in

Group 13 have _______, in Group 14 have four, in Group 15 have five, in Group 16 have six, and in Group 17 have


28

________ valence electrons. All atoms in Group 18 have eight valence electrons, except helium which only has

two. All atoms in sublevels d and f have _________ valence electrons.

How many valence electrons does each of the following elements have?

a) carbon (C)

b) bromine (Br)

c) iron (Fe)

d) potassium (K)

e) aluminum (Al)

LEWIS DOT DIAGRAMS

Because valence electrons are so important to the behavior of an atom, it is useful to represent them with symbols.

A Lewis dot diagram illustrates ___________________ electrons as dots (or other small symbols) around the

chemical symbol of an element. Each dot represents _____________ valence electron. In the dot diagram, the

element’s symbol represents the core of the atom - the nucleus plus all the _______________ electrons.

Write a Lewis dot diagram for a) chlorine b) calcium c) potassium

PERIODIC TRENDS

Because the periodic table relates group and period numbers to valence electrons, it’s useful in predicting atomic

structure and, therefore, ______________________ properties.

Atomic Radius

Atomic radius is half the distance between two __________________ of a diatomic molecule. Atomic size is

influenced by two factors: (1) energy level – A _________________ energy level is further away. (2) charge on

nucleus - More charge (_________________) pulls electrons in closer. As you go down a ___________________,

each atom has another energy level so the atoms get bigger. As you go across a period, the radius gets

____________________. Atoms are in the same energy level, but as you move across the chart, atoms have a

greater ___________________ charge (more protons). Therefore, the outermost electrons are closer.

• Choose the element from the pair with the larger atomic radius.

a) lithium (Li) or beryllium (Be) _________ b) silicon (Si) or tin (Sn) _________

• Choose the element from the pair with the smaller atomic radius.

a) silver (Ag) or gold (Au) _________ b) cesium (Cs) or barium (Ba) _______

Ionic Radius

When an atom gains or loses one or more electrons, it becomes a(n) ______________. Because an electron has a

negative charge, gaining electrons produces a _______________________ charged ion, an anion, whereas losing

electrons produces a positively charged ion, a ________________. As you might expect, the loss of electrons

produces a positive ion with a radius that is ___________________ than that of the parent atom. Conversely, when

an atom gains electrons, the resulting negative ion is larger than the parent atom. Practically all of the elements to

the _____________ of group 4A of the periodic table commonly form positive ions. As with neutral atoms,


29

___________________ ions become smaller moving across a period and become larger moving down through a

group. As you go down a group, you are adding a(n) _________________ level. Ions get bigger as you go down.

Most elements to the right of group 4A (with the exception of the noble gases in group 8A) form negative ions.

These ions, although considerably larger than the positive ions to the left, also decrease in ______________ moving

across a period. Like the positive ions, the negative ions increase in size moving down through a group. Across the

period, nuclear charge __________________ so they get smaller. Energy level changes between anions and cations.

• Choose the element from the pair with the smaller radius.

a) silver (Ag) or the silver ion (Ag1+

) _________

b) oxygen (O) or the oxygen ion (O2-

) _________

• For each of the following pairs, predict which atom is larger.

a) Mg, Sr _________ b) Sr, Sn _________ c) Ge, Sn _________

d) Ge, Br _________ e) Cr, W _________

• For each of the following pairs, predict which atom or ion is larger.

a) Mg, Mg2+

_________ b) S, S2–

_________ c) Ca2+

, Ba2+

______

d) Cl–, I

– _________ e) Na

+, Al

3+ _________

Ionization Energy

Ionization energy (IE) is the amount of energy required to completely _____________________ an electron from a

gaseous atom. Removing one electron makes a ________ ion. The energy required to do this is called the first

ionization energy. The _____________________ the nuclear charge (# of protons), the greater IE. The distance

from the ____________________ increases IE. As you go down a group, first IE decreases because the electron is

further away, thus there is more shielding by the _______________ electrons from the pull of the positive nucleus.

All the atoms in the same period have the same energy level. They have the same shielding, but as you move across

the chart there is a(n) _____________________ nuclear charge. Therefore, IE generally increases from left to right.

• Choose the element from the pair with the greater ionization energy.

a) silver (Ag) or iodine (I) _________ b) oxygen (O) or selenium (Se) ________

• Choose the element from the pair with the smaller ionization energy.

a) chromium (Cr) or tungsten (W) ______ b) sodium (Na) or magnesium (Mg) _______

Electronegativity

Electronegativity is the tendency for an atom to ___________________ electrons to itself when it is chemically

combined with another element. Large electronegativity means it _______________ the electron toward it. The

further you go down a group, the farther the electron is away from the nucleus and the _____________ electrons an

atom has. It is harder to attract extra electrons if the available energy level is far from the nucleus, so the

electronegativity _____________________. As you go across a row, electronegativity increases as the

________________________ character of the elements decreases.


30

• Choose the element from the pair with the greater electronegativity.

a) sodium (Na) or rubidium (Rb) _______ b) selenium (Se) or bromine (Br) _______

• Choose the element from the pair with the smaller electronegativity.

a) magnesium (Mg) or calcium (Ca) _______ b) nitrogen (N) or oxygen (O) _______

Homework / Practice

Write the configuration notation for each of the following elements: 1) sodium

2) iron

3) bromine

4) barium

Write the noble gas notation for each of the following elements:

5) cobalt

6) silver

7) tellurium

8) radium

Determine what elements are denoted by the following electron configurations:

9) 1s22s

22p

63s

23p

4

10) 1s22s

22p

63s

23p

64s

23d

104p

65s

1

11) [Kr] 5s24d

105p

3

12) [Rn] 7s25f

11

Write the orbital notation for the following:

13) C 14) Ne 15) S 16) P 17) B 18) Na

Write configuration notation for atoms containing the following number of electrons:

19) 3 20) 6 21) 8 22) 13

Draw the Lewis Dot Notation for the following elements

23) Sodium

24) Sulfur

25) Silver

26) Aluminum

27) Antimony

28) Argon

Electrons and Periodicity Practice Test

Directions: For questions 1-4, match each of the following terms with a number or chemical symbol from the

periodic table below.

1. Alkaline earth metals: 2. Halogens:


31

3. Noble gases 4. The transition metals

5. Draw the orbital notation for sodium.

6. Given the electron configuration, identify the element 1s2 2s

2 2p

6 3s

2 3p

6 4s

2 3d

7

7. Write the complete configuration notation for silver.

8. Write the shorthand method (Noble Gas notation) for antimony.

9. Give the energy level for the valence electrons in helium.

10. Determine the color of light emitted when an electron jumps from the following quantum levels n=4 to n=2.

11. How many valence electrons does carbon have?

12. Draw the Lewis Dot notation for sodium.

13. Describe why the atomic radius of elements increases as you go down a group.


14. The two main parts of an atom are the

a. Principle energy levels and energy

sublevels

b. Nucleus and kernel

c. Nucleus and energy levels

d. Planetary electrons and energy

levels

15. The sublevel that has only one orbital is identified by the letter

a. s b. p c. d d. f

16. The sublevel that can be occupied by a maximum of ten electrons is identified by the letter

a. s b. p c. d d. f

17. An orbital may never be occupied by

a. 1 electron

b. 2 electrons

c. 3 electrons

d. 0 electrons

18. An atom of beryllium consists of 4 protons, 5 neutrons, 4 electrons. The mass number of this atom is

a. 13 b. 9 c. 8 d. 5

19. Which of the following is the correct electron configuration for the bromide ion, Br1-

?

a. [Ar] 4s24p

5

b. [Ar] 4s23d

104p

5

c. [Ar] 4s23d

104p

6

d. [Ar] 4s23d

104p

65s

1

e. [Ar] 4s23d

103p

6

20. Which is the first element to have 4d electrons in its electron configuration?

a. Ca

b. Sc

c. Rb

d. Y

e. La

21. When electrons in an atom in an excited state fall to lower energy levels, energy is

a. absorbed, only

b. released, only

c. neither released nor absorbed

d. both released and absorbed

22. Which of the following elements has the greatest electronegativity?

a. Mg b. K c. S d. F


32

23. Which of the following elements would have the smallest radius


24. Which of following elements has the lowest first ionization energy


25. Which of the following elements is an alkali metal?


26. Which element's ionic radius is smaller than its atomic radius?

a. neon

b. nitrogen

c. sodium

d. sulfur

27. Which three groups of the Periodic Table contain the most elements classified as metalloids (semimetals)?

a. 1, 2, and 13

b. 2, 13, and 14

c. 14, 15, and 16

d. 16, 17, and 18

28. Which element has the highest first ionization energy?

a. sodium

b. aluminum

c. calcium

d. phosphorus

29. Which of the following elements has the smallest atomic radius?

a. nickel

b. cobalt

c. calcium

d. potassium

30. Which set of elements contains a metalloid?

a. K, Mn, As, Ar

b. Li, Mg, Ca, Kr

c. Ba, Ag, Sn, Xe

d. Fr, F, O, Rn

31. Atoms of elements in a group on the Periodic Table have similar chemical properties. This similarity is most

closely related to the atoms'

a. number of principal energy levels

b. number of valence electrons

c. atomic numbers

d. atomic masses

32. As atoms of elements in Group 16 are considered in order from top to bottom, the electronegativity of each

successive element

a. decreases

b. increases

c. remains the same

33. An atom of which of the following elements has the greatest ability to attract electrons?

a. silicon

b. sulfur

c. nitrogen

d. chlorine

34. At STP, which substance is the best conductor of electricity?

a. nitrogen

b. neon

c. sulfur

d. silver

35. A strontium atom differs from a strontium ion in that the atom has a greater

a. number of electrons

b. number of protons

c. atomic number

d. mass number

36. Which gas is monatomic at STP?

a. chlorine

b. fluorine

c. neon

d. nitrogen


33

37. How many valence electrons does an oxygen atom have?

a. 2 b. 6 c. 8 d. 16

38. The identity of an element is determined by...

a. the number of its protons.

b. the number of its neutrons.

c. the number of its electrons

d. its atomic mass.

39. Which of the following atoms has the largest diameter?

a. F b. Cl c. Br d. I

40. Which of the following elements has the greatest electronegativity?

a. Si b. P c. N d. O

41. Which scientist noted a definite pattern in valence numbers and arranged an early periodic table in order of the

elements atomic mass?

a. Enrico Fermi

b. Dmitri Mendeleev

c. Albert Einstein

d. Madame Curie

42. Which of the following is a noble gas?

a. Sodium

b. Gold

c. Chlorine

d. Neon

43. A gas is called "noble" because

a. it is normally unreactive

b. it is normally inert

c. it has a complete outer energy level

of electrons

d. all of the above

44. Of the following elements, the one that forms cations with varying positive charges is:

a. Fe

b. Na

c. Al

d. Sr

e. N

45. An element having the configuration [Xe]6s1 belongs to the Group:

a. alkali metals

b. halogens

c. alkaline earth

metals

d. None of these

e. noble gases

46. Using the Lewis Dot notation, how many unpaired electrons are there in an atom of tin in its ground state?

a. 4 b. 0 c. 3 d. 2 e. 1

47. Which of the following particles has the greatest atomic radius?

a. Al

b. Si

c. S

d. Al3+

e. P

48. Which of the following forms of electromagnetic radiation has the shortest wavelength?

a. ultraviolet

b. radio waves

c. infrared

d. visible light

e. microwaves

49. For which of the following transitions does the light emitted have the shortest wavelength?

a. n = 4 to n = 2

b. n = 2 to n = 1

c. n = 5 to n = 3

d. n = 4 to n = 3

e. n = 3 to n = 2

50. Researchers at Lawrence Berkeley National Lab have recently formed a new synthetic element with atomic

number 118 and mass number 293. Which of the following elements would have chemical properties most

similar to this new element?

a. Ir

b. Xe

c. Ta

d. Pb

e. S

Unit 4 – Types of Bonding

34

BONDING

As atoms bond with each other, they _____________________ their potential energy, thus creating more stable

arrangements of matter. The force that holds two ________________ together is called a chemical bond. There are

3 types of bonding: ionic, ___________________, and metallic. The number of valence electrons are easily found

by looking up the group number on the periodic table.

Group 1A ___ valence electron








Electron Configurations and Electron Dot Diagrams for Cations

Metals lose electrons to attain noble gas configuration. They make positive ions, ____________.

If we look at an electron configuration, it makes sense. Example: Sodium (Na), 1s22s

22p

63s

1, has _________

valence electron(s). The electron that is removed comes from the ____________ energy level. As a result of the

loss of the electron, the sodium ion (Na+) has the following electron configuration: 1s

22s

22p

6

Calcium has 2 valence electrons. These will come off, forming a positive ion.

Electron Configurations and Electron Dot Diagrams for Anions

Nonmetals gain electrons to attain noble gas configuration. This means they want a(n) ________________ of

electrons, 8 electrons. They make negative ions, ___________________.

If we look at an electron configuration, it makes sense. Example: Sulfur (S), 1s22s

22p

63s

23p

4, has _______ valence

electrons and needs to gain 2 more to have an octet.

The sulfur ion (S-2

) has the same electron configuration as a noble gas: 1s22s

22p

63s

23p

6 Phosphorous has 5 valence

electrons. It will gain _________ electrons to fill the outer shell.

Stable Electron Configurations

All atoms react to achieve __________________ gas configuration. Noble gases, except He, have 2 s electrons and

6 p electrons, totaling 8 valence electrons. They obey the ____________________ rule.

IONIC BONDING

Anions and cations are involved in ionic bonding and are held together by __________________ charges,

electrostatic attraction. The bond is formed through the ______________________ of electrons. Electrons are

transferred to achieve noble gas configuration. Ionic bonds occur between _________________ and nonmetals. All

the electrons must be accounted for! A compound that is composed of _______________ is called an ionic

compound. Note that only the arrangement of electrons has changed. Nothing about the atom’s nucleus has

changed. Ionic compounds have a _______________________ structure, a regular repeating arrangement of ions in


35

the solid. Even though the ions are ___________________ bonded to one another, ionic compounds are

__________________. Strong repulsion breaks crystal apart. The structure is rigid. They have _______________

melting points because of strong forces between ions. They also conduct electricity in the _________________ and

dissolved states. Any compound that conducts electricity when melted or dissolved in water is a(n)

___________________________.

� How many valence electrons must an atom have in its outer energy level in order to be considered stable?

The energy required to separate one mole of the ions of an ionic compound is called ____________________

energy, which is expressed as a negative quantity. The greater (that is, the more negative) the lattice energy is, the

______________________ the force of attraction between the ions. Lattice energy tends to be

__________________________ for more-highly-charged ions (those atoms that have more electrons to give or

those atoms that can take more electrons). Lattice energy also tends to be greater for __________________ ions.

� Between the following ionic compounds, which would be expected to have the higher (more negative)

lattice energy? LiF or KBr

� Between the following ionic compounds, which would be expected to have the higher (more negative)

lattice energy? NaCl or MgS

The electronegativity difference for two elements in an ionic compound is greater than or equal to

_______________.

COVALENT BONDING

A _______________________ is an uncharged group of two or more atoms held together by covalent bonds.

Covalent compounds occur between two ___________________ or a nonmetal and hydrogen. The attraction of two

atoms for a shared _______________ of electrons is called a covalent bond. In a covalent bond, atoms share

electrons and neither atom has an ionic ______________________. Covalent bonds occur between 2

___________________________ because nonmetals hold onto their valence electrons. They can’t give away

electrons to bond, yet, they still want _______________ gas configuration. They get it by sharing valence electrons

with each other. By sharing both atoms get to count the electrons toward noble gas configuration. A

____________________ bond is formed from the sharing of two valence electrons. The electronegativity difference

for two elements in a covalent compound is between _________ and 1.7.

• Do atoms that share a covalent bond have an ionic charge?

Sometimes atoms share more than one pair of valence electrons. A ____________________ bond is when atoms

share two pair of electrons, 4 electrons. A triple bond is when atoms share three pair of electrons, _____ electrons.

Triple bonds are ________________________ and shorter than double bonds. Double bonds are stronger and

shorter than ______________________ bonds.

METALLIC BONDING

The bonding in metals is explained by the _______________________ ____________ model, which proposes that

the atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons that surrounds metallic


36

__________________. These delocalized electrons are not held by any specific atom and can ________________

easily throughout the solid. Metals hold onto their valence electrons very _______________________. Think of

them as positive ions floating in a sea of electrons. Because electrons are free to move through the solid, metals

conduct _______________________. Metals generally have extremely ______________ melting points because it

is difficult to pull metal atoms completely away from the group of cations and attracting electrons. Metals are

________________________ (able to be hammered into sheets). Metals are also ________________________

(able to be drawn into wire) because of the mobility of the particles. Electrons allow atoms to slide by. A mixture

of elements that has metallic properties is called a(n) _____________________.

Homework / Practice

Complete the table by identifying the charge of each of the elements listed and then indicating the formula for ionic

compounds formed between the two substances

Charge O N P S Cl F

Charge 2-

Na 1+ Na2O

Mg

Ca

Al

Li

Zn

What type of bonding is present in the following compounds:

1) SbBr3

2) Ag

3) MgBr2

4) ClO2

5) KCl

6) Fe

7) PbO

8) FeCl2

9) NI3

10) CO2

11) Ni

12) Au

13) LiF

14) Al2O3

15) N2O3

16) Mg3P2

17) CCl4

18) H2O

POLARITY and VSEPR

How each atom fares in a tug-of-war for shared electrons is determined by comparing the

_________________________________ of the two bonded atoms. Recall that electronegativity is the measure of

the ability of an atom in a bond to ________________________ electrons. Atoms with large electronegativity

values, such as fluorine, attract shared valence electrons more __________________ than atoms such as sodium that

have small electronegativities. Electronegativity is a periodic property. With only a few exceptions,

electronegativity values_____________________ as you move from left to right in any period of the periodic table.

Within any group, electronegativity values decrease as you go ___________________ the group. Fluorine has the

highest value of ____________. The greater the difference between the electronegativities of the bonding atoms,

the more _____________________________ the electrons are shared and the more polar the bond.


37

If the electronegativity difference between the two elements in question is:

between 0.0 – 1.7, the bond is ______________________

greater than 1.7, the bond is _____________________

When the electronegativity difference in a bond is 1.7 or greater, the sharing of electrons is so unequal that you can

assume that the electron on the less electronegative atom is ________________________ to the more

electronegative atom. For example, ∆EN for cesium and fluorine is 4.0 − 0.7 = 3.3. Therefore the bond is

_________________.

COVALENT BONDS AND POLARITY

When the atoms in a bond are the same, the electrons are shared ________________________. This results in a

_______________________ covalent bond. _________________________ elements (H2, O2, N2, Cl2, Br2, I2, and

F2) have pure nonpolar covalent bonds. All other covalent bonds are polar. The electron sharing is not equal, but it

is not so unequal that a complete _____________________ of electrons takes place.

Consider hydrogen and chlorine. Hydrogen has an electronegativity of 2.20, and chlorine has an electronegativity of

3.16. The ________ pulls harder on the electrons because its electronegativity is greater. The electrons spend more

time near the Cl. These symbols, __________________ plus (δ+) and delta minus (δ-), represent a partial positive

charge and a partial negative charge.

Polar molecules are molecules with a positive and a negative ______________. This requires two things to be true:

The molecule must contain _______________ bonds. (This can be determined from differences in

electronegativity.) Symmetry cannot ______________________ out the effects of the polar bonds. (Must

determine geometry first.)

• In the following compounds, determine whether the molecule is polar or nonpolar

a. hydrogen fluoride (HF)

b. water (H2O)

c. carbon tetrachloride (CCl4)

d. ammonia (NH3)

e. carbon dioxide (CO2)

VSEPR

VSEPR stands for Valence Shell ______________________ ________________ Repulsion. It predicts three-

dimensional geometry of molecules. The valence shell includes the ______________________ electrons. The

electron pairs try to get as far away as possible to _______________________ repulsion. You can determine the

angles of the bonds. VSEPR is based on the number of pairs of valence electrons, both bonded and unbonded. An

unbonded pair of electrons is referred to as a _______________pair. Calculate the number of bonds and then draw

the dot-dash diagram. The shape of the molecule and bond angle can be determined from this diagram.

LINEAR

Each hydrogen has 1 line attached which represents _______________ electrons. No extra electrons are needed

around hydrogen to have the 2 electrons needed after bonding.


38

A hydrogen molecule is linear. The electrons attempt to maximize their distance from one another by having bond

angle of ____________. Linear compounds are NOT ____________________.

TETRAHEDRAL

Consider CH4. which has __________ bonds! The element you have only one of goes in the

_______________________. The other elements surround it. Connect the elements with a single

_________________ (a single bond). Remember a line represents ___________ electrons. Count your lines for

each element to determine if extra electrons need to be added. Carbon has 4 lines attached which represent

_________ electrons. No extra electrons are needed around carbon. Each hydrogen atom has one line attached

which represents 2 electrons. No extra electrons are needed around hydrogen. Single bonds fill all atoms. There

are _________ bond pairs of electrons pushing away. The electrons can _________________ their distance from

one another by forming a 3-D shape. The furthest they can get away is ___________. This basic shape is a

tetrahedral, a pyramid with a triangular base. The tetrahedral is the shape for everything with 4 bond pairs and

____________ lone pairs around the central atom.

TRIGONAL PYRAMIDAL

Consider phosphorous trichloride (PCl3). How many bonds are in this molecule? ________ . Sketch the dot-dash

diagram for phosphorous trichloride. Please include all electrons. Only the electrons around the

______________________ atom affect the shape. The shape is a basic _______________________________ but

you can’t see the lone pair. The shape is called trigonal pyramidal. The bond angle is ____________ between the

chlorines because the electron pair forces the chlorines closer to each other.

BENT

Consider water (H2O). How many bonds are in this molecule? _____ Sketch the dot-dash diagram for water.

Please include all electrons. Only the electrons around the central atom affect the shape. The shape is still basic

tetrahedral, but you can’t see the _________ lone pairs. The shape is called bent. The bond angle between

hydrogens is ____________.

TRIGONAL PLANAR

Consider H2CO. How many bonds are in this molecule? _____. Sketch the dot-dash diagram for H2CO. Please

include all electrons. (Carbon is the central atom.) The farthest you can get the elements apart is __________. The

shape is flat and called trigonal planar.

• Determine the number of bonds, draw the dot-dash diagram, state the VSEPR shape and provide the bond

angle for the following compounds

a. CO2 b. BCl3 c. SCl2 d. SiF4


39

Homework/Practice

Draw the Lewis structure for each of the following compounds, identify the shape of the molecule, and identify the

polarity of the molecule.

1. CCl4

2. BF3

3. NF3

4. SiO2

5. H2S

INTERMOLECULAR FORCES

Intermolecular forces are forces of _______________________. They are what make solid and liquid molecular

compounds possible. The three intermolecular forces are _________________ bonds, dipole–dipole forces and

London ____________________________ forces.

Hydrogen Bonding

A hydrogen bond is a _________________________________________ attraction that occurs between molecules

containing a hydrogen atom bonded to a small, highly electronegative atom with at least ____________ lone

electron pair. For a hydrogen bond to form, hydrogen must be bonded to a fluorine,

__________________________, or nitrogen atom. F, O, and N are very electronegative so it is a very

_______________________ dipole. Hydrogen bonding is the _________________________ of the intermolecular

forces. Examples include H2O, NH3, and HF.

Dipole-dipole Forces

Polar molecules contain ___________________________ dipoles; that is, some regions of a polar molecule are

always ___________________________ negative and some regions of the molecule are always partially positive.

Attractions between _____________________________ charged regions of polar molecules are called dipole–

dipole forces. Neighboring polar molecules orient themselves so that oppositely charged regions _______________

up. Opposites attract but are not completely hooked as in ionic solids. Dipole-dipole forces depend on the number

of _______________________. Bigger molecules result in more electrons, and more electrons mean

________________________ forces. Dipole–dipole forces are stronger than dispersion forces as long as the

molecules being compared have approximately the same mass. Examples of compounds that exhibit dipole-dipole

forces include CO, HCl, and PH3.

London Dispersion Forces

Dispersion forces are ____________________ forces that result from temporary shifts in the

______________________ of electrons in electron clouds. Remember that the electrons in an electron cloud are in

constant _____________________. When two nonpolar molecules are in close contact, especially when they

collide, the electron cloud of one molecule _______________________ the electron cloud of the other molecule.


40

The electron density around each nucleus is, for a moment, greater in one region of each cloud. Each molecule

forms a __________________________ dipole. When temporary dipoles are close together, a weak dispersion

force exists between oppositely charged regions of the dipoles. Due to the temporary nature of the dipoles,

dispersion forces are the __________________________ intermolecular force. Dispersion forces exist between

____________ gases and compounds that are nonpolar. Examples include Ar, Cl2, Br2, CH4, and CO2. Dispersion

forces ______________________ as the mass of the molecule increases. C2H6 (MW = 30.0 g/mol) has stronger

dispersion forces than CH4 (MW = 16.0 g/mol). This difference in dispersion forces explains why fluorine and

chlorine are gases, bromine is a __________________________, and iodine is a solid at room temperature. The

molecular mass of iodine is greater than that of bromine, and bromine has a greater mass than chlorine.

Intermolecular Forces

To determine what type of intermolecular force a compound has, ask yourself the following questions.

� Does the compound contain hydrogen attached to N, O, or F?

o If yes, the force is hydrogen bonding.

Determine the number of bonds from the Wetter Way and draw the dash-dot diagram.

� Does the central element of the compound contain any lone pairs of electrons?

o If yes, the force is dipole-dipole.

� Does the central element of the compound contain ZERO lone pairs of electrons?

o If yes, the force is dispersion.

Determine the type of intermolecular force in each of the following compounds

1) BCl3 _____________________________ 2) Xe _____________________________

3) NH3 _____________________________ 4) CH4 _____________________________

5) H2 _____________________________ 6) CH3Cl ___________________________

7) HF _____________________________ 8) HBr ____________________________

Types of Bonding Practice Test

1. In a complete sentence, compare and contrast metallic bonds and ionic bonds.

Directions- For each of the following pairs of elements, write the formula for the ionic compound that would form

between them

2. K and Cl

3. Na and N

4. Al and O

5. Calcium and Chlorine

6. Zinc and Sulfur

7. Lithium and Phosphorous

Directions- Draw the Lewis structure , Identify the shape of the molecule, Identify the polarity of the bonds,

Identify the polarity of the molecule, Identify the IMF that would be exhibited

8. CCl4

9. SF2

10. SiO2

11. BI3

12. PCl3

13. N2

14. What does IMF stand for? Which of the three IMF’s is the weakest?


41

15. What type of bond exists between atoms of potassium and chloride in a crystal of potassium chloride?

a. Hydrogen bond

b. Ionic bond

c. Polar covalent bond

d. Nonpolar covalent bond

e. Metallic bond

16. What type of bond exists between atoms in a nitrogen molecule?

a. Hydrogen bond

b. Ionic bond



e. Metallic bond

17. What type of bond exists between atoms of iron in a sample of iron?

a. Hydrogen bond

b. Ionic bond



e. Metallic bond

18. All of the following have covalent bonds except

a. HCl

b. CCl4

c. H2O

d. CsF

e. CO2

19. Which of the following atoms normally forms monatomic molecules?

a. Cl

b. H

c. O

d. N

e. He

20. The complete loss of an electron of one atom to another atom with the consequent formation of electrostatic

charges is said to be

a. Covalent bonding

b. Polar covalent bonding

c. Ionic bonding

d. Coordinate covalent bonding

21. When a metal atom combines with a nonmetal atom, the nonmetal atom will

a. lose electrons and decrease in size

b. lose electrons and increase in size

c. gain electrons and decrease in size

d. gain electrons and increase in size

22. Which formula represents a molecular substance?

a. CaO

b. CO

c. Li2O

d. Al2O3

23. Which combination of atoms can form a polar covalent bond?

a. H and H

b. H and Br

c. N and N

d. Na and Br

24. Fluorine atoms tend to.______.when they form chemical compounds with metals.

a. lose electrons

b. gain electrons

c. neither lose nor gain electrons...they usually share electrons equally with metals.

d. Fluorine atoms do not form compounds with other atoms...fluorine is an inert gas.

25. What is a compound composed of?

a. two or more different elements that are physically combined in a fixed proportion

b. two or more different mixtures that are physically combined in a fixed proportion

c. two or more different elements that are chemically combined in a fixed proportion

d. two or more different elements that are chemically combined in a variable proportion

26. Which of the following compounds is most likely to be ionic?

a. CO2

b. CCl4

c. MgCl2

d. HBr


42

27. How many unshared electron pairs must be included in the Lewis structure for water, H2O?

a. 3

b. 2

c. 1

d. 4

e. 0

28. Which of the following molecules must contain at least one double bond

a. H2O

b. CCl4

c. H2O2

d. CH3I

e. CH3COOH

29. How can a chemical compound be broken?

a. can be broken down by physical means

b. can be broken down by chemical means

c. cannot be broken down

d. can be broken down by physical or chemical means

30. Nitrogen triiodide, NI3, is an unstable molecule that is used as a contact explosive. Its molecular structure is:

a. none of these

b. octahedral

c. square planar

d. tetrahedral

e. pyramidal

31. In which of the following compounds does the bond between the central atom and chlorine have the greatest

ionic character?

a. BCl3

b. FeCl2

c. CCl4

d. HCl

e. CaCl2

32. The Lewis structure for hydrogen cyanide is:

a.

b.

c.

d.

e.

33. In the Lewis structure for CH2Cl2, the number of unshared electron pairs is:

a. 10

b. 8

c. 2

d. 4

e. 6

34. The only intermolecular forces existing between oxygen molecules are:

a. ion-ion attractive forces

b. hydrogen bonding forces

c. permanent dipole forces

d. nuclear forces

e. London dispersion forces

35. Reactions between alkali metals and phosphorous result in compounds with the formula:

a. M3P

b. None of these

c. M2P

d. M2P3

e. MP3

36. A particle X contains 10 electrons, seven neutrons and has a net charge of 3-. The particle is:

a. a nitride ion

b. obviously polyatomic

c. an oxide ion

d. a neon ion

e. none of these are correct

Unit 5 – Nomenclature

43

NAMING COMPOUNDS AND WRITING FORMULAS

A compound is made of two or more ______________________. The name should tell us how many and

what type of atoms. There are two types of compounds: ___________________ compounds and molecular

compounds. The simplest ratio of the ions represented in an ionic compound is called a ______________________

unit. The overall charge of any formula unit is ________________. In order to write a correct formula unit, one

must know the charge of each ion. Atoms are electrically _____________________. They have the same number

of protons and electrons. ________________ are atoms, or groups of atoms, with a charge.

Ions have a different numbers of electrons. An anion is a _____________________ ion. An anion has

gained electrons. Nonmetals can ________________ electrons. The charge is written as a superscript on the right.

F1-

has gained _________ electron. O2-

has gained __________ electrons. A ___________________ is a positive

ion. It is formed by __________________ electrons. There are more _____________________ than electrons.

______________________ form cations. K1+

has lost one electron. Ca2+

has lost __________ electrons. The

charges of monatomic ions, or ions containing only one atom, can often be determined by referring to the periodic

table or table of common ions based on group number. The charge of a monatomic ion is equal to its

_________________________ number. For most of the Group ________ elements, the Periodic Table can tell what

kind of ion they will form from their location. Elements in the same group have similar properties, including the

charge when they are ions.

NAMING CATIONS

We will use the systematic way. For cations, if the charge is always the same (Group A) just write the

_________________ of the metal. Transition metals (as well as tin and lead) can have more than one type of

charge. The charge is indicated with ___________________ numerals in parenthesis. Zinc (Zn2+

) and silver (Ag1+

),

although transition metals, only have __________ possible charge. Roman numerals ARE NOT used for zinc and

silver. Li1+

is called the ____________________ ion. __________ is called the Strontium ion. Fe2+

is called the

iron (II) ion. Iron is a transition metal, so the charge is not always the same. The name of the metal is written, and

the charge is denoted with Roman numerals in parenthesis. Pb2+

is called the lead __________ ion.

Name the following cations.

a) Ca2+

_________________________ b) Al3+

___________________________

c) Sn4+

_________________________ d) Na+ _________________________

e) Fe3+

_________________________ f) Cu+ _________________________

WRITING FORMULAS FOR CATIONS

Write the formula for the metal. If a Roman numeral is in parenthesis use that number for the

_____________________. Indicate the charge with a superscript. If no Roman numeral is given, find the Group A


44

metal on the periodic table and determine the charge from the _____________________ number. The formula for

the nickel (II) ion is Ni2+

. The formula for the gallium (III) ion is ____________.

Write the formulas for the following cations.

a) magnesium ion ________________ b) copper (II) ion ___________________

c) potassium ion ________________ d) silver ion _________________

e) chromium (VI) ion ________________ f) mercury (II) ion ________________

NAMING ANIONS

Naming monatomic anions is always the same. Change the element ending to – ___________. F is the symbol for

fluorine, F1-

is fluoride. Cl1-

is called the chloride ion. _______ is called the oxide ion.

Name the following anions.

a) S2-

_________________________ b) Br1-

___________________________

c) N3-

_________________________ d) As3-

_________________________

e) Te2-

_________________________

WRITING FORMULAS FOR ANIONS

Write the formula for the nonmetal. Find the Group A nonmetal on the periodic table and determine the charge from

the column number.

Write the formulas for the following anions.

a) iodide ion ________________ b) phosphide ion ___________________

c) selenide ion ________________ d) carbide ion _________________

IONIC COMPOUNDS

Oxidation numbers can be used to determine the chemical formulas for ionic compounds. If the oxidation number

of each ion is _________________________ by the number of that ion present in a formula unit, and then the

results are added, the sum must be _______________. In the formula for an ionic compound, the symbol of the

_________________ is written before that of the anion. Subscripts, or small numbers written to the lower

______________________ of the chemical symbols, show the numbers of ions of each type present in a formula

unit.

BINARY IONIC COMPOUNDS

Binary ionic compounds are composed of a metal bonded with a ________________________. Name the metal ion

using a Roman numeral in parenthesis if necessary. Follow this name with the name of the nonmetal ion.

Name the following binary ionic compounds.

a) NaCl _________________

b) Ca3P2 ________________

c) CuO _________________

d) SnBr2 ________________

e) Fe2S3 ________________

f) AlF3 _________________

g) KCl _________________

h) Na3N ________________

i) CrN _________________


45

Write the symbol for the metal. Determine the oxidation number from either the column number or the Roman

numeral and write it as a superscript to the right of the metal’s symbol. To the right of the metal’s symbol, write the

symbol for the nonmetal. Determine the oxidation number from the column number and write it as a superscript to

the right of the nonmetal’s symbol.

� Example: potassium fluoride - K1+

F1-

If the two oxidation numbers add together to get zero, the formula

is a one-to-one ratio of the elements. Answer = KF

� Example: aluminum sulfide - Al3+

S2-

If the two oxidation numbers DO NOT add together to get zero, you

will need to “criss-cross” the superscripts. These numbers now become subscripts. Omit all positive and

negative signs and omit all 1’s. Answer = Al2S3

Write the formulas for the following binary ionic compounds.

a) lithium selenide __________________ b) tin (II) oxide __________________

c) tin (IV) oxide __________________ d) magnesium fluoride ________________

e) copper (II) sulfide __________________ f) iron (II) phosphide _________________

g) gallium (III) nitride __________________ h) iron (III) sulfide __________________

TERNARY IONIC COMPOUNDS

Ternary ionic compounds are composed of at least _________ elements. Name the metal ion, using a Roman

numeral in parenthesis if necessary. Follow this name with the name of the polyatomic ion. Polyatomic ions are

groups of atoms that stay together and have a __________________. Examples are provided on page 7 of the

NCDPI Reference Tables for Chemistry. There is one polyatomic ion with a positive oxidation number (NH4+) that

may come first in a compound. Name the ion. Follow this name with the name of the anion or second polyatomic

ion. Certain polyatomic ions, called ________________________, contain oxygen and another element.

Name the following ternary ionic compounds.

a) LiCN __________________ b) Fe(OH)3 ___________________

c) (NH4)2CO3 __________________ d) NiPO4 __________________

e) NaNO3 __________________ f) CaSO4 __________________

g) (NH4)2O __________________ h) CuSO3 __________________

Write the symbol for the metal or ammonium ion. Write the oxidation number as a superscript to the right of the

metal’s/ammonium ion’s symbol. To the right of the metal’s symbol, write the symbol for the nonmetal or

polyatomic ion. Write the oxidation number as a superscript to the right of the nonmetal’s/polyatomic ion’s symbol.

� Example: potassium nitrate - K1+

NO31-

If the two oxidation numbers add together to get zero, the formula is a

one-to-one ratio of the elements. Answer = KNO3

� Example: aluminum hydrogen sulfate – Al3+

HSO4 1-

If the two oxidation numbers DO NOT add together to

get zero, you will need to “criss-cross” the superscripts. These numbers now become subscripts. Parentheses

are to be placed around polyatomic ions before criss-crossing. Omit all positive and negative signs and omit all

1’s. Answer = Al(HSO4)3


46

Write the formulas for the following ternary ionic compounds.

a) ammonium chloride __________________ b) ammonium sulfide _________________

c) barium nitrate __________________ d) zinc iodate __________________

e) sodium hypochlorite __________________ f) chromium (III) acetate ______________

g) iron (II) dichromate __________________ h) mercury (I) bromate ________________

MOLECULAR COMPOUNDS

Molecular compounds are made of molecules. They are made by joining _______________________ atoms

together into molecules. A molecular compound’s name tells you the number of atoms through the use of

____________________________.

1 mono- 4 tetra- 7 hepta- 10 deca-

2 di- 5 penta- 8 octa-

3 tri- 6 hexa- 9 nona-

The name will consist of two words. Prefix name prefix name –ide One exception is we don’t write mono- if

there is only one of the first element. The following double vowels cannot be used when writing names: (oa) and

(oo).

� Example: NO2 There is one nitrogen. Mononitrogen But, you cannot use mono- on the first element, so drop

the prefix. There are two oxygens. dioxygen You need the suffix –ide. dioxide (Answer: nitrogen dioxide).

� Example: N2O There are two nitrogens. Dinitrogen There is one oxygen. monooxygen You cannot run (oo)

together, so monoxygen. You need the suffix –ide. monoxide (Answer: dinitrogen monoxide).

Name the following molecular compounds.

a) Cl2O7 ____________________________ b) CBr4 ____________________________

c) CO2 ________________________ d) BCl3 ___________________________

When writing a formula of a molecular compound from the name, you will not need to criss-cross oxidation

numbers. Molecular compounds name tells you the number of atoms through the use of prefixes.

� Example: diphosphorus pentoxide The name implies there are 2 phosphorous atoms and 5 oxygens.

Answer: P2O5

� Example: sulfur hexafluoride The name implies there is 1 sulfur atom and 6 fluorines. Answer: SF6

Write the formulas for the following molecules.

a) tetraiodide nonoxide __________________ b) nitrogen trioxide __________________

c) carbon tetrahydride __________________ d) phosphorus trifluoride ______________


47

IONIC MOLECULAR

Smallest Piece Molecule

Types of Elements metal and nonmetal

State of Matter solid

Melting Point Low <300°C

ACIDS

Acids are compounds that give off hydrogen ions (H+) when dissolved in water. Acids will always contain one or

more hydrogen ions next to an anion. The anion determines the name of the acid.

Binary Acids

Binary acids contain hydrogen and an anion whose name ends in –ide. When naming the acid, put the prefix hydro-

and change -ide to -ic acid.

� Example: HCl The acid contains the hydrogen ion and chloride ion. Begin with the prefix hydro-, name

the nonmetallic ion and change -ide to -ic acid. Answer: hydrochloric acid

� Example: H2S The next step is change -ide to -ic acid, but for sulfur the “ur” is added before -ic.

Answer: hydrosulfuric acid

Name the following binary acids.

a) HF ____________________________ b) H3P ____________________________

The prefix hydro- lets you know the acid is binary. Determine whether you need to criss-cross the oxidation

numbers of hydrogen and the nonmetal.

� Example: hydrobromic acid The acid contains the hydrogen ion and the bromide ion. H1+

Br1-

The two

oxidation numbers add together to get zero. Answer: HBr

� Example: hydrotelluric acid The acid contains the hydrogen ion and the telluride ion. H1+

Te2-

The two

oxidation numbers do NOT add together to get zero, so you must criss-cross. Answer: H2Te

Write the formulas for the following binary acids.

a) hydroiodic acid __________________ b) hydroselenic acid __________________

Ternary Acids

The acid is a ternary acid if the anion has oxygen in it. The anion ends in -ate or -ite. Change the suffix -ate to -ic

acid. Change the suffix -ite to -ous acid. The hydro- prefix is NOT used!

� Example: HNO2 The acid contains the hydrogen ion and nitrite ion. Name the polyatomic ion and change

-ite to -ous acid. Answer: nitrous acid

� Example: H3PO4 The acid contains the hydrogen ion and phosphate ion. Name the polyatomic ion and

change -ate to -ic acid. Answer: phosphoric acid


48

Name the following ternary acids.

a) H2CO3 ____________________________ b) H2SO4 __________________________

c) H2CrO4 ________________________ d) HClO2 __________________________

The lack of the prefix hydro- from the name implies the acid is ternary, made of the hydrogen ion and a polyatomic

ion. Determine whether you need to criss-cross the oxidation numbers of hydrogen and the polyatomic ion.

� Example: acetic acid The polyatomic ion must end in –ate since the acid ends in -ic. The acid is made of

H+ and the acetate ion. H

1+ C2H3O2

1- The two charges when added equal zero. Answer: HC2H3O2

� Example: sulfurous acid Again the lack of the prefix hydro- implies the acid is ternary, made of the

hydrogen ion and a polyatomic ion. The polyatomic ion must end in –ite since the acid ends in -ous. The

acid is made of H+ and the sulfite ion. H

1+ SO3

2- The two charges when added do not equal zero, so

you must crisscross the oxidation numbers. Ignore the negative sign and ones are understood. Answer:

H2SO3

Write the formulas for the following binary acids.

a) perchloric acid __________________ b) iodic acid __________________

c) dichromic acid __________________ d) hypochlorous acid ________________

Homework / Practice

Name each of the following compounds

1. CuS

2. CuCl2

3. Ni(C2H3O2)2

4. Co2S3

5. CrBr2

6. AlPO4

7. CaCO3

8. Mg(OH)2

9. Ba(CN)2

10. K2SO4

11. NH4NO3

12. SbBr3

13. Ag

14. KCl

15. PbO

16. FeCl2

17. Al2O3

18. Mg3P2

19. NH4Cl

20. Fe(NO3)3

21. TiBr3

22. Pb(SO4)2

23. P4O6

24. N2O3

25. SiF4

26. P4S10

27. Cl2O3

28. PCl3

Write the formula for each of the following compounds

29. iron (II) chloride

30. mercury (I) bromide

31. chromium (III) oxide

32. manganese (II) nitride

33. cobalt (III) phosphide

34. copper (II) sulfide

35. potassium perchlorate

36. aluminum sulfate

37. silver phosphate

38. cobalt (III) nitrite

39. ammonium sulfite

40. carbon monoxide

41. selenium difluoride

42. calcium sulfide

43. diphosporus trioxide

44. magnesium fluoride

45. zinc oxide

46. arsenic tribromide

47. carbon tetrafluoride

48. sodium sulfide

49. vanadium (IV) carbonate

50. tin (II) nitrite

51. cobalt (III) oxide

52. titanium (II) acetate


49

APPLICATIONS OF THE MOLE

Molar Mass

The molar mass of a compound is the mass of a mole of the _________________________________ particles of the

compound. Because each representative particle is composed of two or more atoms, the molar mass of the

compound is found by adding the molar masses of all of the ________________ in the representative particle. To

determine the molar mass of an element, find the element’s symbol on the periodic table and round the mass so there

is __________ digit beyond the decimal. For example, the molar mass of carbon (C) is ____________ g/mol, of

chlorine (Cl) is ___________ g/mol and of iron (Fe) is _____________ g/mol. In the case of NH3, the molar mass

equals the mass of one mole of nitrogen atoms plus the mass of ___________ moles of hydrogen atoms.

Molar mass of NH3 = molar mass of N + 3 (molar mass of H)

Molar mass of NH3 = 14.0 + 3 (1.0) = 17.0 g/mol

Mole Conversions/Mass Conversions

You can use the molar mass of a compound to convert between mass and moles, just as you used the molar mass of

elements to make these conversions.

1. How many moles are in 56.3 g of Mg?

2. How many moles are in 295 grams of Cr(OH)3?

3. How many grams are in 3.45 moles of NaCl?

4. How many grams are in 1.6 moles of K2CrO4?

Percent Composition

Recall that every chemical compound has a definite composition - a composition that is always the same wherever

that compound is found. The composition of a compound is usually stated as the percent by mass of each element in

the compound. The percent of an element (X) in a compound can be found in the following way.

( )( )ompoundMolarMassC

sXmolarmassXX

'#% =

Determine the percent composition of chlorine in calcium chloride (CaCl2). First, analyze the information available

from the formula. A mole of calcium chloride consists of one mole of calcium ions and ___________ moles of

chloride ions. Next, gather molar mass information from the atomic masses on the periodic table. To the mass of

one mole of CaCl2, a mole of calcium ions contributes ______________ g, and two moles of chloride ions

contribute 2 x 35.5 g = 71.0 g for a total molar mass of ___________ for CaCl2. Finally, use the data to set up a

calculation to determine the percent by mass of an element in the compound.

1. Determine the percent composition of carbon in sodium acetate (NaC2H3O2).

2. Calculate the percent composition aluminum of aluminum oxide (Al2O3).

3. Determine the percent composition of oxygen in magnesium nitrate, which has the formula Mg(NO3)2.

4. Determine the percent composition of sulfur in aluminum sulfate, which has the formula Al2(SO4)3.

5. Determine the percent composition of oxygen in zinc nitrite, which has the formula Zn(NO2)2.


50

Empirical Formula

You can use percent composition data to help identify an unknown compound by determining its empirical formula.

The empirical formula is the ________________________ whole-number ratio of atoms of elements in the

compound. In many cases, the empirical formula is the actual formula for the compound. For example, the simplest

ratio of atoms of sodium to atoms of chlorine in sodium chloride is 1 atom Na : 1 atom Cl. So, the empirical

formula of sodium chloride is Na1Cl1, or NaCl, which is the true formula for the compound. The formula for

glucose is C6H12O6. The coefficients in glucose are all divisible by 6. The empirical formula of glucose is CH2O.

1. Determine the empirical formula for Tl2C4H4O6.

2. Determine the empirical formula for N2O4.

The percent composition of an unknown compound is found to be 38.43% Mn, 16.80% C, and 44.77% O.

Determine the compound’s empirical formula. Because percent means “parts per hundred parts,” assume that you

have ___________ g of the compound. Then calculate the number of moles of each element in the 100 g of

compound. To obtain the simplest whole-number ratio of moles, _________________ each number of moles by the

smallest number of moles. Find the whole number mole ratio for the compound. These numbers become

the____________________________ in the empirical formula.

1. Determine the empirical formula of the following compound: 31.9 g Mg, 27.1 g P

2. The composition of an unknown acid is 40.00% carbon, 6.71% hydrogen, and 53.29% oxygen. Calculate

the empirical formula for the acid.

3. The composition of an unknown ionic compound is 60.7% nickel and 39.3% fluorine. Calculate the

empirical formula for the ionic compound.

4. The composition of a compound is 6.27 g calcium and 1.46 g nitrogen. Calculate the empirical formula for

the compound.

5. Find the empirical formula for a compound consisting of 63.0% Mn and 37.0% O.

Molecular Formula

For many compounds, the empirical formula is not the true formula. A molecular formula tells the

___________________ number of atoms of each element in a molecule or formula unit of a compound. The

molecular formula for a compound is either the same as the empirical formula or a whole-number

_______________________ of the empirical formula. In order to determine the molecular formula for an unknown

compound, you must know the molar mass of the compound in addition to its empirical formula. Then you can

compare the molar mass of the compound with the molar mass represented by the empirical formula.

1. The molecular mass of benzene is 78 g/mol and its empirical formula is CH. What is the molecular

formula for benzene? HINT: Calculate the molar mass represented by the formula CH. Calculate the

whole number multiple, n, and apply it to its empirical formula.

2. The simplest formula for butane is C2H5 and its molecular mass is about 60.0 g/mol. What is the molecular

formula of butane?


51

3. What is its molecular formula of cyanuric chloride, if the empirical formula is CClN and the molecular

mass is 184.5 g/mol?

4. The simplest formula for vitamin C is C3H4O3. Experimental data indicates that the molecular mass of

vitamin C is about 180. What is the molecular formula of vitamin C?

5. The composition of silver oxalate is 71.02% silver, 7.91% carbon, and 21.07% oxygen. If the molar mass

of silver oxalate is 303.8 g/mol, what is its molecular formula?

Homework / Practice

1) A compound is found to have (by mass) 48.38% carbon, 8.12% hydrogen and the rest oxygen. What is its

empirical formula?

2) A compound is known to have an empirical formula of CH and a molar mass of 78.11 g/mol. What is its

molecular formula?

3) Another compound, also with an empirical formula if CH is found to have a molar mass of 26.04 g/mol. What is

its molecular formula?

4) A compound is found to have 1.121 g nitrogen, 0.161 g hydrogen, 0.480 g carbon and 0.640 g oxygen. What

is its empirical formula? (Note that masses are given, NOT percentages.)

5) Phentyfloroform contains 57.54% C, 3.45% H, and 39.01% F. What is its empirical formula?

6) Calculate the empirical formula of the compound containing 81.8% C and 18.2% H.

7) The active ingredient in chocolate is theobromine; a sample was analyzed and determined to be composed of:

147.0 g C 14.0 g H 56.0 g O 98.0 g N

a. Determine the % composition for each element.

b. Determine the empirical formula for theobromine.

c. The molecular weight of theobromine is known to be 180.0 g/mole. What is the molecular

formula?

8) What is the empirical formula of a compound if a 50.0 g sample of it contains 9.1 g Na, 20.6 g Cr, and

22.2 g O?

9) NutraSweet is 57.14% C, 6.16% H, 9.52% N, and 27.18% O. Calculate the empirical formula of NutraSweet

and find the molecular formula. (The molar mass of NutraSweet is 294.30 g/mol)

10) A compound consists of 85% silver and 15% florine by mass. What is the empirical formula?

11) A compound consists of 40% calcium, 12% carbon, and 48% oxygen by mass. What is the empirical formula by

mass?

12) Phentyfloroform contains 57.54% C, 3.45% H, and 39.01% F. What is its empirical formula?

13) Freons are gaseous compounds used in refrigeration. A particular freon contains 9.93% carbon, 56% chlorine,

and 31.4% fluorine. What is the empirical formula?


52

Naming and Formula Math Practice Test

Directions: Write the formula for the compound

1. sodium phosphide

2. iron (II) perchlorate

3. vanadium (V) nitrite

4. nickel (I) oxide

5. magnesium hydroxide

6. cesium nitride

7. nitrogen trichloride

8. hydroxic acid

9. carbon tetrahydride

Directions: Name the compound

10. KCl

11. FeSO4

12. Li2O

13. Cr2S3

14. Ca3N2

15. Fe2S3

16. CuI2

17. PBr3

18. CO2

19. HNO3

20. What is the percent nitrogen in potassium nitrate?

21. What is the ratio of barium ions to Nitrogen ions in a formula unit of barium nitrate?

22. A compound is found to have 46.67% nitrogen, 6.70% hydrogen, 19.98% carbon and 26.65% oxygen. What is

its empirical formula?

23. Determine the empirical formula of N3O6

24. The empirical formula of a compound is CH2. Its molecular mass is 70g/mol. What is its molecular formula?

25. Determine the molecular formula of a compound that is composed of 40.0% carbon, 6.7% hydrogen and

53.5% oxygen. The molecular mass is 120.0g/mol.

26. Which of the following is a binary compound?

a. hydrogen sulfide

b. hydrogen sulfate

c. ammonium sulfide

d. ammonium sulfate

27. Which of the following is a binary compound?

a. potassium chloride

b. ammonium chloride

c. potassium chlorate

d. ammonium chlorate

28. Which is the correct formula for dinitrogen monoxide?

a. NO

b. N2O

c. NO2

d. N2O3

29. Which of the following represents the correct formula for aluminum oxide?

a. AlO b. Al2O3 c. AlO2 d. Al2O

30. Which of the following is the correct name for NaHCO3?

a. sodium hydrogen carbonate

b. sodium acetate

c. nitrogen hydrogen carbonate

d. sodium hydrogen carbon trioxide

31. What is the name of CaCl2?

a. calcium dichloride

b. calcium (II) chloride

c. monocalcium dichloride

d. calcium chloride

32. What is the name of Mg(NO3)2

a. Magnesium nitrate

b. Magnesium (II) nitrate

c. Magnesium dinitrate

d. Magnesium nitrogen oxide

33. What is the name of P2O5?

a. phosphorus oxide

b. phosphorus pentaoxide

c. diphosphorus pentaoxide

d. phosphorus (III) oxide


53

34. What is the formula for sulfur hexachloride?

a. S5Cl b. SHCl c. SCl5 d. SCl6

35. What is the name of the formula Fe(NO3)2?

a. iron nitrate

b. iron (II) nitrate

c. iron dinitrate

d. iron (III) nitrate

36. Which of the following is not a type of chemical formula?

a. Empirical

b. Molecular

c. Structural

d. Parabola

37. What is the approximate percentage oxygen in the formula mass of Ca(NO3)2?

a. 28

b. 42

c. 58

d. 96

e. 164

38. Which formulas could represent the empirical formula and the molecular formula of a given compound?

a. CH2O and C4H6O4

b. CHO and C6H12O6

c. CH4 and C5H12

d. CH2 and C3H6

e. CO and CO2

39. When combining with nonmetallic atoms, metallic atoms generally will

a. lose electrons and form negative

ions

b. lose electrons and form positive ions

c. gain electrons and from negative

ions

d. gain electrons and form positive ions

40. What is the empirical formula of the compound whose molecular formula is P4O10?

a. PO b. PO2 c. P2O5 d. P8O20

41. What is the percent by mass of oxygen in magnesium oxide, MgO?

a. 20% b. 40% c. 50% d. 60%

42. A compound is 86% carbon and 14% hydrogen by mass. What is the empirical formula for this compound?

a. CH b. CH2 c. CH3 d. CH4

43. What is the gram formula mass of (NH4)3PO4?

a. 113 g

b. 121 g

c. 149 g

d. 404 g

44. What is the empirical formula of a compound that contains 85% Ag and 15% F by mass?

a. AgF b. Ag2F c. AgF2 d. Ag2F2

45. What is the molecular formula for a compound that is 46.16% carbon, 5.16% hydrogen and 48.68% fluorine if

the molar mass of this compound is 156.12 g?

a. C3H4F2

b. C5H10F5

c. C6H8F4

d. C6H6F3

46. Which of the following is named incorrectly?

a. H2CO2 : carbonous acid

b. HClO2 : chlorous acid

c. H2SO4 : sulfuric acid

d. HClO : hydrochlorous acid

e. H3PO3 : phosphorous acid

47. A sample of an alcohol is tested and found to contain 52% carbon, 35% oxygen, and 13% hydrogen by mass.

Tests indicate that the molecular weight of the molecule is between 30 and 80. What is the molecular formula of

the alcohol?

a. C2H5OH

b. C3H7OH

c. C5H11OH

d. C4H9OH

e. CH3OH

Unit 6- Chemical Reactions

54

CHEMICAL REACTIONS

All chemical reactions have two parts: (1) A substance that undergoes a reaction is called a

__________________________. In other words, reactants are the substances you start with. (2) When reactants

undergo a chemical change, each new substance formed is called a ___________________________. In other

words, the products are the substances you end up with. The reactants turn into the products.

Reactants → Products. In a chemical reaction, the way atoms are joined is changed. Atoms aren’t

__________________________ or destroyed.

Chemical reactions can be described several ways.

� In a sentence: Copper reacts with chlorine to form copper (II) chloride.

� In a word equation: Copper + chlorine → copper (II) chloride

The arrow separates the reactants from the products. The arrow reads “reacts to ________________.” The plus

sign reads “_____________.” (s) after the formula implies the substance is a ___________________. (g) after the

formula implies the substance is a gas. (l) after the formula implies the substance is a ______________________.

(aq) after the formula implies the substance is aqueous, a solid dissolved in _____________________. __________

used after a product indicates a gas, same as (g). ↓ used after a product indicates a ________________, same as (s).

_____________ indicates a reversible reaction. ________________ or ________________ shows that heat is

supplied to the reaction. ___________________ is used to indicate a catalyst used supplied, in this case, platinum.

A catalyst is a substance that ____________________ ____________ a reaction without being changed by the

reaction. Enzymes are biological or ______________________ catalysts.

There are seven elements that never want to be alone. They form ________________________ molecules. H2 , N2,

O2 , F2 , __________ , Br2 , I2. (1 + 7 pattern on the periodic table)

The following are indications that a chemical reaction has occurred: formation of a

____________________________, evolution of a gas, _____________________ change, and absorption or release

of ________________.

A ________________________ formula uses formulas and symbols to describe a reaction. All chemical equations

are sentences that describe reactions.

� Convert the following sentences to chemical equations.

a) Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form solid iron (II) chloride and

hydrogen sulfide gas.

b) Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon

dioxide gas and sodium nitrate dissolved in water.

Balancing Equations

Atoms can’t be ______________________ or destroyed. All the atoms we start with we must end up with. A

balanced equation has the same number of each element on both _________________ of the equation. Example:

C + O2 → CO This equation is NOT balanced. There is one carbon atom on the left and ________ on the right.


55

There are two oxygen atoms on the left and only one on the right. We need one more oxygen atom in the products.

We can’t change the formula, because it describes what it is. In order to have two oxygen atoms, another CO must

be produced. But where did the other carbon come from? We must have started with two carbon atoms. The

balanced chemical equation is 2 C + O2 → 2 CO

Rules for Balancing

� Write the correct formulas for all the reactants and products.

� Count the number of atoms of each type appearing on both sides.

� Balance the elements one at a time by adding coefficients (the numbers in front).

� Check to make sure it is balanced.

Never change a _________________________ to balance an equation. If you change the formula you are

describing a different reaction. Never put a coefficient in the middle of a formula. 2 NaCl is okay; Na2Cl is not.

� Balance the following reaction. H2 + O2 → H2O

Balance elements in the following order: (1) metals; (2) nonmetals; (3) hydrogen; and (4) oxygen

If an atom appears more than once on a side, balance it last. If you fix everything except one element, and it is even

on one side and odd on the other, double the first number, then move on from there.

� Balance the following equations.

1) _____ CH4 + _____ O2 → _____ CO2 + _____ H2O

2) _____ AgNO3 + _____ Cu → _____ Cu(NO3)2 + ______ Ag

3) _____ Mg + _____ N2 → _____ Mg3N2

4) _____ P + _____ O2 → ______ P4O10

5) _____ Na + _____ H2O → _____ H2 + _____ NaOH

6) _____ Pb(NO3)2 + _____ K2CrO4 → ______ PbCrO4 + ______ KNO3

7) _____ MnO2 + _____ HCl → _____ MnCl2 + ______ H2O + _____ Cl2

8) _____ Ba(CN)2 + _____ H2SO4 → _____ BaSO4 + _____ HCN

9) _____ Zn(OH)2 + _____ H3PO4 → _____ Zn3(PO4)2 + _____ H2O

TYPES OF REACTIONS

Reactions fall into 5 categories. We will recognize the type by the reactants. We will be able to predict the

products. For some we will be able to predict whether they will happen at all.

Synthesis Reactions

Synthesize means to put together. Whenever two or more substances combine to form one single product, the

reaction is called a synthesis reaction. Examples: Ca + O2 → CaO and P2O5 + 3 H2O → 2 H3PO4

We can predict the products if they are two elements. All you need to do is combine the elements, metals first, and

criss-cross oxidation numbers if necessary. After predicting the product, the reaction must be balanced.

� Mg + N2 →

� CaO + H2O →


56

On page 6 of the Chemistry Reference Packet, this reaction is an example of “Metal oxide - water reactions.” The

product listed in the packet is “base.” A base is a metallic hydroxide.

� SO2 + H2O →

On page 6 of the Chemistry Reference Packet, this reaction is an example of “Nonmetal oxide - water reactions.”

The product listed in the packet is “________________.” The acid is a ternary acid. Ternary acids start with

_________ and end in O. The other element goes in the center. This is the only compound for which you can add

the number of elements and use these numbers as subscripts.

� Write and balance the following synthesis reactions.

a) Ca + Cl2 →

b) Fe + O2 →

HINT: Use iron (II).

c) K2O + H2O →

d) Al + O2 →

e) SO3 + H2O →

f) N2O5 + H2O →

Decompositions Reactions

The word decompose implies the compound will “fall apart.” In a decomposition reaction, one compound breaks

down into _____________ or more simple substances.

NaCl → Na + Cl2 CaCO3 → CaO + CO2

We can easily predict the products if it is a binary compound. A binary compound is made up of only two elements.

The compound merely falls apart into its elements.

� H2O →

� HgO →

If the compound has more than two elements, you must consult the Reference Tables, page 6.

� NiCO3 →

NiCO3 is called nickel (II) carbonate. The packet states that a metallic carbonate decomposes to form a MO

(metallic oxide) and CO2. The metallic oxide is nickel (II) oxide.

� Use the Chemistry Reference Tables to write and balance the following decomposition reactions.

a) KClO3 →

b) CaBr2 →

c) Li2CO3 →

d) Cr(OH)2 →

e) NaHCO3 →

Single Replacement

In a single-displacement reaction, one element takes the place of another in a compound. One reactant must be an

element, and the one reactant must be a _______________________. The products will be a different element and a

different compound. F2 + LiCl → LiF + Cl2

Remember zinc, Zn, always forms a ___________ ion doesn’t need parenthesis. ZnCl2 is zinc chloride. In addition,

silver, Ag, always forms a ___________ ion. AgCl is silver chloride.


57

Some single replacement reactions do not occur because some elements are not as ________________ as others. A

more active element _________________________ a less active element. There is a list referred to as the Activity

Series on page 7 of your Chemistry Reference Packet. A higher element on the list replaces lower element. If the

element by itself is lower on the list, the reaction will ___________ occur.

Metals replace metals (and hydrogen)

� K + NaCl →

Potassium wants to replace ________________________. You must check the activity series on page 7 of your

Chemistry Reference Packet to see if this is possible. Because K is higher, potassium can replace sodium. The

potassium will bond with the _____________________ and the sodium will be alone. You must always check to

see if the compound formed needs criss-crossing. Check for balancing.

� Sn + FeCl3 →

Because Sn is NOT higher, tin cannot replace iron. No reaction occurs.

� Write and balance the following single replacement reaction.

a) Rb + AlN →

b) Zn + HCl →

c) Ag + CoBr2 →

Metals replace hydrogen

� Na + H2O (cold) →

Think of water as HOH. Metals high enough on the activity series replace the first ______ and combine with the

OH1-

(hydroxide) according to page 6 of the Reference Tables. Is sodium above hydrogen and higher than the line

marked “Replace hydrogen from cold water” on the activity series? Since the answer is yes, sodium replaces the

first H, bonding with hydroxide.

� Mg + HCl →

Metals higher on the activity series replace the H and combine with the nonmetal according to page 6 of the

Reference Tables. Hydrogen gas is a second product. Is magnesium above hydrogen on the activity series?

� Write and balance the following single replacement reactions.

a) Ag + H2O (steam) →

b) Cu + H2SO4 →

c) Cr + H3PO4 → (HINT: Use Cr3+

)

d) Ca + H2O (steam) →

Nonmetals can replace other ________________________. This is limited to F2 , Cl2 , Br2 and I2 The order of

activity is listed in the Chemistry Reference Packet, page 7. Higher replaces _____________.

� F2 + HCl →

Is fluorine above chlorine in the activity series of halogens? Since the answer is yes, fluorine replaces the chlorine,

bonding with hydrogen.

� Write and balance the following single replacement reactions.

a) Br2 + KCl →

b) Cl2 + KI →


58

Double Replacement

In double-displacement reactions, the positive portions of two ___________________ compounds are interchanged.

The reactants must be two ionic compounds or ______________. Double replacement reactions usually take place

in ________________________ solution.

� NaOH + FeCl3 →

The positive ions change place. You must check to see if you need to criss-cross the products. Now balance. A

double replacement reaction will only happen if one of the products: (1) doesn’t dissolve in water and forms a

__________________, (2) is a _____________ that bubbles out, or (3) is a _________________________

compound usually water.

3NaOH + FeCl3 → Fe(OH)3 + 3NaCl

None of the products are familiar gases. Both products are ionic (not covalent) because they start with metals. We

must consult the Solubility Rules on page 6 of the Chemistry Reference Tables to determine if a solid (a

________________________) is formed. The “Soluble” side of the Solubility Rules states that Group 1 (IA) salts

are soluble; therefore, NaCl is soluble and is NOT the precipitate. The “Insoluble” side of the Solubility Rules states

that all hydroxides except Group 1, Sr, Ba and NH41+

are INSOLUBLE. Therefore, Fe(OH)3 is the precipitate

(solid). In molecular equations, the formulas of the compounds are written as though all species existed as

molecules or whole units. An ionic equation shows dissolved ionic compounds in terms of their free ions. Ions that

are not involved in the overall reaction are called spectator ions. The net ionic equation indicates only the species

that actually take part in the reaction. The following steps are useful for writing ionic and net ionic equations:

1) Write a balanced molecular equation for the reaction.

2) Rewrite the equation to indicate which substances are in ionic form in solution. Remember that all soluble

salts (and other strong electrolytes), are completely dissociated into cations and anions. This procedure

gives us the ionic equation.

3) Lastly, identify and cancel spectator ions on both sides of the equation to arrive at the net ionic equation.

Example: sodium hydroxide + iron (III) chloride yields iron (III) hydroxide + sodium chloride

Balanced Molecular Equation: 3 NaOH + FeCl3 � Fe(OH)3 + 3 NaCl

Complete Ionic Equation:

3Na1+

+ 3OH1-

+ Fe3+

+ 3Cl1-

� Fe(OH)3 + 3Na1+

+ 3Cl1-

Net Ionic Equation: 3OH1-

+ Fe3+

� Fe(OH)3

� Write and balance the following double replacement reaction. Assume the reaction takes place. In addition,

identify the precipitate and write the net ionic equation.

a) CaCl2 + NaOH →

b) CuCl2 + K2S →

c) KOH + Fe(NO3)3 →

d) (NH4)2SO4 + BaF2 →

Combustion

A combustion reaction is one in which a substance rapidly combines with ____________________ to form one or

more oxides. Combustion reactions involve a compound composed of only C and H (and maybe O) that is reacted


59

with oxygen gas. If the combustion is complete, the products will be CO2 and __________________. Combustion

reactions produce heat, and are therefore considered exothermic reactions.

� Complete and balance the following combustion reactions.

a) C4H10 + O2 →

b) C6H12O6 + O2 →

c) C8H8 + O2 →

d) C3H8O3 + O2 →

To determine which type a reaction is, look at the reactants. (E = element and C = compound)

E + E Synthesis

C Decomposition

E + C Single replacement

C + C Double replacement

CH cpd + O2 Combustion

Note: Two other common synthesis reactions include: nonmetallic oxide + water and metallic oxide + water.

� Identify whether the reaction is synthesis, decomposition, single replacement, double replacement or

combustion.

a) H2 + O2 →

b) H2O →

c) Mg(OH)2 + H2SO3 →

d) HgO →

e) KBr + Cl2 →

f) Zn + H2SO4 →

g) AgNO3 + NaCl →

h) C6H6 + O2 →

Homework / Practice

Directions: Balance the following equations.

1. S + O2

→ SO3

2. P + O2 → P

2O

3

3. Na + O2 → Na

2O

4. Al + N2

→ AlN

5. Fe + O2

→ Fe2O

3

6. MgSO4

•7H2O → MgSO

4 + H

2O

7. NH4NO

3 → N

2O + H

2O

8. NaNO3 → NaNO

2 + O

2

9. H2O

2 → H

2O + O

2

10. Al + CuSO4

→ Al2(SO

4)

3 + Cu

11. ZnS + O2

→ ZnO + SO2

12. Na + H2O → NaOH + H

2

13. Al + H2SO

4 → Al

2(SO

4)

3 + H

2

14. Fe(OH)3 + H

2SO

4 → Fe

2(SO

4)

3 + H

2O

15. AgNO3 + K

2CrO

4 → Ag

2CrO

4 + KNO

3

16. MgCl2

+ NaOH → Mg(OH)2

+ NaCl

17. AgNO3 + H

2S → Ag

2S + HNO

3

18. Al(OH)3 + NaOH → NaAlO

2 + H

2O

19. C4H

10 + O

2 → CO

2 + H

2O

20. C3H

6 + O

2 → CO

2 + H

2O

21. C5H

8 + O

2 → CO

2 + H

2O

22. C6H

12O

6 + O

2 → CO

2 + H

2O

Predict the products of the following reactions

23. Na + O2 →

24. K2O + H2O →

25. Al2O3 + H2O →

26. N2O5 + H2O →

27. Ag2O →

28. HNO2 →


60

29. Fe(OH)3 →

30. ZnCO3 →

31. Cs2CO3 →

32. RbClO3 →

33. ZnS + O2 →

34. K + H2O →

35. Al + Pb(NO3)2 --->

36. Cl2 + NaI →

37. Al + CuCl2 →

38. Br2 + CaI2 →

39. Al + HCl →

40. Zn + H2SO4 →

41. Fe + CuSO4 →

42. Ca(OH)2 + H3PO4 →

43. AgNO3 + KCl →

44. Na2CO3 + H2SO4 →

45. Al(OH)3 + HC2H3O2 →

46. Cr2(SO3)3 + H2SO4 →

Write the balanced equation (including states) and identify the type of reaction:

47. Aqueous solutions of ammonium chloride and lead (II) nitrate produce lead (II) chloride precipitate and

aqueous ammonium nitrate.

48. Iron metal reacts with aqueous silver nitrate to produce aqueous iron (III) nitrate and silver metal.

49. Solid potassium nitrate yields solid potassium nitrite and oxygen gas.

50. Calcium metal reacts with chlorine gas to produce solid calcium chloride.

Chemical Reactions and Balancing Practice Test

1. Write the balanced equation of the reaction that occurs when calcium carbonate decomposes to form

calcium oxide and carbon dioxide.

Directions: Balance and identify the type of reaction

2. ____K3PO4 + ____Al(NO3)3 → ____KNO3 + ____AlPO4

3. ____Fe2O3 + ____Al � ____Fe + ____Al2O3

4. ____NaOH � ____Na2O + ____H2O

5. ____HCl + ____Mg � ____MgCl2 + ____H2

6. ____C2H4 + ____O2 � ____CO2 + ____H2O

7. Write the balanced equation of the synthesis reaction that occurs when iron metal and oxygen gas react to

form iron (III) oxide.

8. Write the balanced equation of the combustion reaction that occurs when ethane (C2H6) reacts with oxygen

to form carbon dioxide and water.

9. Write the balanced equation of the reaction that occurs when calcium carbonate decomposes to form

calcium oxide and carbon dioxide.

Directions: Predict the products and balance the reaction

10. _____Na + ______O2 �

11. _____Al + _____Pb(NO3)2 �

12. _____NaF + _____Br2 �

13. In the following unbalanced reaction, what is the coefficient of HOH once the reaction is balanced?

H2CO3 + KOH --> HOH + K2CO3

a. 1 b. 2 c. 3 d. 4


61

14. What are the different types of chemical reactions?

a. synthesis, fusion, combustion, fission, and decomposition

b. single replacement, combustion, and double replacement

c. synthesis, fission, single replacement, combustion, and fusion

d. synthesis, decomposition, combustion, single and double replacement

15. What type of reaction is represented by 2 or more elements forming a compound?

a. Decomposition

b. Synthesis

c. Combustion

d. Single replacement

16. Decomposition is the burning of hydrocarbons in the presence of oxygen.

a. True b. False

17. Which equation represents a double replacement reaction?

a. CaCO3 → CaO + CO2

b. CH4 + 2O2 → CO2 + 2H2O

c. LiOH + HCl → LiCl + H2O

d. C3H8 + 5O2 → 3CO2 + 8H2O

18. MgSO4 + BaCl2 → MgCl2 + BaSO4, is an example of what type of chemical reaction?

a. Single replacement

b. Synthesis

c. Combustion

d. Double replacement

19. Zn + 2 AgNO3 → 2 Ag + Zn(NO3)2, is an example of what type of chemical reaction?

a. Synthesis

b. Single replacement

c. Decomposition

d. Double replacement

20. The cation of one compound replaces the cation in another compound in a double replacement reaction.

a. True b. False

21. Given the unbalanced equation CuS + O2 → CuO + SO2 When it is balanced, what is the sum of the

coefficients?

a. 8 b. 9 c. 10 d. 11

22. Which statement best describes the conservation of atoms in all balanced chemical equations?

a. There is a conservation of mass, number of protons, and charge.

b. There is a conservation of mass, electronegativity, and charge.

c. There is a conservation of only energy, and charge.

d. There is a conservation of mass, energy, and charge.

23. Which one of these chemical reactions is balanced?

a. Na + Cl2 → NaCl

b. H2 + O2 → H2O

c. CuCO3 → CuO + CO2

d. KClO3 → KCl + O2

24. Which is the correct way of setting up a word equation for this balanced chemical equation,

2Na + Cl2 → 2NaCl?

a. Sodium react with chlorine gas to produce sodium chloride.

b. 2 moles of sodium react with 1 mole of chlorine to yield 1 mole of sodium chloride.

c. 2 moles of sodium react with 1 mole of chlorine gas to yield 2 mole of sodium chloride.

d. 2 moles of sodium added with 1 mole of chlorine gas to yield 1 mole of sodium chloride.

25. 2 moles of copper react with 1 mole of oxygen gas to yield 2 moles of copper (ll) oxide. How would you

express this word equation into a balanced chemical equation?

a. Cu + O → CuO

b. Cu + O2 → CuO

c. 2Cu + O → CuO

d. 2Cu + O2 → 2CuO


62

26. When the following reaction is balanced, the sum of all of the coefficients in the equation is:

NaCl + H2SO4 + MnO2 → Na2SO4 + MnSO4 + H2O + Cl2

a. 11

b. 10

c. 6

d. 16

e. 14

27. Which equation represents combustion?

a. 4 Fe + 3 O2 � 2 Fe2O3

b. 2H2O -� 2H2 + O2

c. CH4 + 2O2 � CO2 + 2 H2O

d. Cu + 2 AgNO3 � Cu(NO3)2 + 2 Ag

28. Given the unbalanced equation: Al + O2 �Al2O3 When this equation is completely balanced using the

smallest whole numbers, what is the sum of the coefficients?

a. 9 b. 7 c. 5 d. 4

29. The balanced equation for the complete combustion of benzene, C6H6, is

a. C6H6 + 12 H2O � 6 CO2 + 15 H2

b. 2 C6H6 + 9 O2 � 12 CO + 6 H2O

c. C6H6 + O2 � CO2 + H2O

d. 2 C6H6 + 15 O2 � 12 CO2 + 6 H2O

30. Which equation shows conservation of atoms?

a. H2 + O2 → H2O

b. H2 + O2 → 2H2O

c. 2H2 + O2 → 2H2O

d. 2H2 + 2O2 → 2H2O

31. When ethanol undergoes complete combustion, the products are carbon dioxide and water.

__ C2H5OH + __ O2 � __ CO2+ __ H2O What are the respective coefficients when the equation is

balanced with the smallest whole numbers?

a. 2, 7, 4, 6

b. 1, 3, 2, 3

c. 2, 2, 1, 4

d. 1, 2, 3, 2

e. 2, 4, 6, 4

Unit 7- Stoichiometry

63

STOICHIOMETRY

The word stoichiometry is Greek for “________________________ elements.” The calculations of quantities in

chemical reactions are based on a ______________________ equation. We can interpret balanced chemical

equations several ways. Using the methods of stoichiometry, we can measure the amounts of substances involved in

chemical reactions and relate them to one another. The group or unit of measure used to count numbers of atoms,

molecules, or formula units of substances is the ______________ (abbreviated mol).

Moles in Chemical Reactions

The coefficients tell us how many moles of each kind of element or compound we have.

2 Al2O3 → 4 Al + 3 O2

2 moles of aluminum oxide form 4 moles of aluminum and 3 moles of oxygen gas.

2 H2 + O2 → 2 H2O

___ mole(s) of hydrogen gas and ___ mole of oxygen form ___ mole(s) of water.

2 Na + 2 H2O → 2 NaOH + H2

___ moles of sodium and ___ moles of water form ___ moles of sodium hydroxide and ___ mole of hydrogen gas.

2 Al2O3 → 4 Al + 3 O2

Every time we use 2 moles of Al2O3 we make 3 moles of O2. Every time we use 2 moles of Al2O3 we make 4 moles

of Al.

• Using the balanced equation above, how many moles of O2 are produced when 3.34 moles of Al2O3

decompose?

• 2 C2H2 + 5 O2 → 4 CO2 + 2 H2O

a) If 3.84 moles of C2H2 are burned, how many moles of O2 are needed?

b) How many moles of C2H2 are needed to produce 8.95 moles of H2O?

c) If 2.47 moles of C2H2 are burned, how many moles of CO2 are formed?

• SiCl4 + 2 Mg → 2 MgCl2 + Si

3.74 mol of Mg would make how many moles of Si?

Mass in Chemical Reactions

2 Al2O3 → 4 Al + 3 O2

2 x (102.0) grams of aluminum oxide form 4 x (27.0) grams of aluminum and 3 x (32.0) grams of oxygen. The mass

of Al2O3 was found by adding the masses of 2 aluminums & 3 oxygens.

(2 x 27.0 + 3 x 16.0 = 102.0)

2 H2 + O2 → 2 H2O

2 x (_________) grams of hydrogen and ___ x (32.0) grams of oxygen form ___ x (_______) grams of water.


64

2 Na + 2 H2O → 2 NaOH + H2

___ x (23.0) grams of sodium and 2 x (________) grams of water form ___ x (_________) grams of sodium

hydroxide and ___ x (________) grams of hydrogen gas.

The law of conservation of _____________ applies in chemical reactions. The mass of the reactants equals the mass

of the ________________________in a balanced chemical equation.

• Show that the following equation follows the Law of Conservation of Mass.

2 Al2O3 → 4 Al + 3 O2

Mass – Mole Stoichiometry

The mass of 1 mole of a pure substance is called its _________________ mass. To convert the mass of an element

or compound to the number of moles, use the mass of 1 mol as a conversion factor. We can convert

___________________ to moles using the periodic table. Then we must apply the mole to mole conversion to

change chemicals using the balanced equation. Finally we will turn the moles back to grams using the periodic

table.

• 2 C2H2 + 5 O2 → 4 CO2 + 2 H2O

a) How many moles of C2H2 are needed to produce 8.95 g of H2O?

b) If 2.47 moles of C2H2 are burned, how many grams of CO2 are formed?


How many moles of Mg are needed to make 9.3 g of Si?

• 3 Al + 3 NH4ClO4 → Al2O3 + AlCl3 + 3 NO + 6 H2O

How many moles of water are produced when 32 grams of aluminum are used?

• CO2 + 2 LiOH → Li2CO3 + H2O

What mass of water can be produced from 3.66 moles of lithium hydroxide (LiOH)?

• 2 Al + 3 I2 → 2 AlI3

Calculate the mass of AlI3 (Aluminum Iodide) that can be produced from 3.00 mol of Al.

Mass – Mass Stoichiometry

• 2 Fe + 3 CuSO4 → Fe2(SO4)3 + 3 Cu

If 10.1 g of Fe are added to a solution of copper (II) sulfate, how much solid copper would form?

• 2 Al + 3 I2 → 2 AlI3

Calculate the mass of I2 needed just to react with 35.0 g of Al.


How many grams of MgCl2 are produced along with 9.3 g of silicon?

• 3 Al + 3 NH4ClO4 → Al2O3 + AlCl3 + 3 NO + 6 H2O

a) How many grams of Al must be used to react with 652 g of NH4ClO4?

b) How many grams of NO are produced if 150.0 grams of AlCl3 are also produced?


65

Particles in Chemical Reactions

The number of things in one mole is 6.022 x 1023

. This big number has a short name: the Avogadro constant.

Atom - ______________________

Molecule - Molecular compound (non-metals) or ______________________ (O2 etc.)

Formula unit - _________________ Compounds (Metal and non-metal or metal and a polyatomic ion)

2 Al2O3 → 4 Al + 3 O2

2 x (6.022 x 1023

) formula units of aluminum oxide form 4 x (6.022 x 1023

) atoms of aluminum and

3 x (6.022 x 1023

) molecules of oxygen.

2 H2 + O2 → 2 H2O

2 x (___________) molecules of hydrogen and ___ x (6.022 x 1023

) molecules of oxygen form ___ x

(___________) molecules of water.

2 Na + 2 H2O → 2 NaOH + H2

___ x (6.022 x 1023

) atoms of sodium and ___ x (___________________) molecules of water form ___ x

(__________________) formula units of sodium hydroxide and ___ x (6.022 x 1023

) molecules of hydrogen gas.

Gases and Reactions

In gas conversions, liters of a gas are converted to moles and vice-versa. ____________ stands for standard

temperature and pressure. 0ºC is standard _________________________, and 1 atmosphere is standard

pressure. At STP, ____________ L of a gas = 1 mole

• 2 H2O → 2 H2 + O2

If 6.45 grams of water are decomposed, how many liters of oxygen will be produced at STP?

• CH4 + 2 O2 → CO2 + 2 H2O

How many liters of CH4 at STP are required to completely react with 17.5 L of O2?

• 2 C8H18 + 25 O2 → 16 CO2 + 18 H2O

Octane, C8H18, is one of the hydrocarbons in gasoline. How many liters of oxygen are required, at STP, to

burn 1.00 g of octane?

• 2 C4H10 + 13 O2 → 8 CO2 + 10 H2O

How many liters of CO2 at STP will be produced from the complete combustion of 23.2 g C4H10?

• 2 NiS + 3 O2 → 2 NiO + 2 SO2

What volume of sulfur dioxide is produced from 123 grams of nickel (II) sulfide at STP?

According to Avogadro, equal volumes of gas, at the _____________ temperature and pressure, contain the same

number of particles. _______________ are numbers of particles. We can also change between particles and liters at

STP.

• 2 C4H10 + 13 O2 → 8 CO2 + 10 H2O

a) How many molecules of CO2 at STP will be produced from the complete combustion of 18.2 L C4H10 ?


66

b) How many molecules of O2 at STP are needed to produce 18.2 L of steam?

c) How many liters of CO2 at STP are produced from 3.2 x 1024

molecules of butane, C4H10?

• 4 NH3 + 6 NO → 5 N2 + 6 H2O

Nitrogen monoxide is a pollutant found in smokestack emissions. How many liters of ammonia, NH3, at

STP are needed to produce 1.4 x 1023

molecules of H2O?

Homework / Practice

Solve the following problems. The reactions may not be balanced.

1. If 20.0 g of magnesium react with excess hydrochloric acid, how many grams of magnesium chloride are

produced? Mg + HCl → MgCl2 + H

2

2. How many moles of oxygen gas are produced in the decomposition of 5.00 g of potassium chlorate?

KClO3

→ KCl + O2

3. If excess ammonium sulfate reacts with 20.0 g of calcium hydroxide, how many grams of ammonia (NH3)

are produced? (NH4)

2SO

4 + Ca(OH)

2 → CaSO

4 + NH

3 + H

2O

4. If excess sulfuric acid reacts with 0.2564 moles of sodium chloride, how many grams of hydrochloric acid

are produced? H2SO

4 + NaCl → HCl + Na

2SO

4

5. How many grams of silver phosphate are produced if 10.0 g of silver acetate react with excess sodium

phosphate? AgC2H

3O

2 + Na

3PO

4 → Ag

3PO

4 + NaC

2H

3O

2

6. What volume of chlorine gas, measured at STP, is needed to produce 10.0 g of potassium permanganate

(KMnO4)? K

2MnO

4 + Cl

2 → KMnO

4 + KCl

7. Suppose that you could decompose 0.250 mol of Ag2S into its elements.

a. How many moles of silver would form?

b. How many moles of sulfur would form from 38.8 g of silver sulfide?

c.

8. Ammonia (NH3) is made industrially by reacting nitrogen gas and hydrogen gas under pressure, at high

temperature and in the presence of a catalyst. If 4.0 mol of hydrogen react, how many moles of ammonia

will be produced?

9. How many liters of Cl2 can be produced from 5.60 mole HCl at STP? 4 HCl + O2 → 2 Cl

2 + 2 H

2O

10. Given the equation Al4C

3 + 12 H

2O → 4 Al(OH)

3 + 3 CH

4 How many moles of water are needed

to react with 100 g Al4C

3?

11. How many grams of zinc phosphate are formed when 10.0 g of Zn are reacted with the phosphoric acid?

The other product is hydrogen gas.

12. Given the equation C2H

4 + 3 O

2 → 2 CO

2 + 2 H

2O

a. If 6.0 mol of CO2 are produced, how many moles of O

2 were reacted?

b. How many liters of O2 are required for the complete reaction of 45 g of C

2H

4 at STP?

c. If 18.0 g of CO2 are produced, how many grams of H

2O are produced?


67

Balance the following equations to use with questions 16 – 21:

13. ____ Al + ____ O2 → ____ Al2O3

14. ____ Cu + ____ AgNO3 → ____ Ag + ____ Cu(NO3)2

15. ____ Zn + ____ HCl → ____ ZnCl2 + ____ H2

Perform the following calculations:

16. Zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen. How many moles of HCl are

required to produce 7.50 moles of ZnCl2?

17. Copper metal reacts with silver nitrate to form silver and copper (II) nitrate. How many grams of copper

are required to form 250 g of silver?

18. When aluminum is burned in excess oxygen, aluminum oxide is produced. How many grams of oxygen

are required to produce 0.75 moles of Al2O3?

19. Copper metal reacts with silver nitrate to form silver and copper (II) nitrate. How many moles of silver

will be produced from 3.65 moles of silver nitrate?

20. When 9.34 g of zinc react with excess hydrochloric acid how many grams of zinc chloride will be

produced?

21. How many liters of oxygen gas at STP are required to react with 65.3 g of aluminum in the production of

aluminum oxide?

22. Given the following equation: Na2O + H2O ---> 2 NaOH How many grams of Na2O are required to

produce 1.60 x 102 grams of NaOH?

23. Given the following equation: 8 Fe + S8 → 8 FeS

a. What mass of iron is needed to react with 16.0 grams of sulfur?

b. How many grams of FeS are produced?

24. Given the following equation: Cu + 2 AgNO3 ---> Cu(NO3)2 + 2 Ag

a. How many moles of Cu are needed to react with 3.50 moles of AgNO3?

b. If 89.5 grams of Ag were produced, how many grams of Cu reacted?

25. The average human requires 120.0 grams of glucose (C6H12O6) per day. How many grams of CO2 (in the

photosynthesis reaction) are required for this amount of glucose? The photosynthetic reaction is:

6 CO2 + 6 H2O ---> C6H12O6 + 6 O2

Stoichiometry Practice Test

Directions: Solve the following problems, showing all work.

1. Balance the equation ______NaOH � ______Na2O + _____H2O

2. How many moles of water are produced from 4 moles of sodium hydroxide?


68

3. Balance the equation ____KCl + _____O2 � _____KClO3

4. How many moles of potassium chlorate are produced from 9 moles of oxygen?

5. Fe2O3 + 2Al � 2Fe + Al2O3 What mass of aluminum oxide is produced when 4 moles of aluminum

react?

6. P4O10 + 6H2O � 4H3PO4 How many grams of phosphoric acid are produced by the reaction of 12.5 g of

water?

7. Using the reaction in question 6, how many grams of P4O10 must react if the reaction produces 25 moles

H3PO4?

8. Balance the equation _____HNO3 + ______Cu � _______Cu(NO3)2 + ______H2

9. How many moles of nitric acid must react in order to form 83 g of copper nitrate?

10. Balance the equation ______NH3 + ________O2 � _______NO + ______H2O

11. What mass of nitrogen monoxide will be formed when 7.2 g nitrogen trihydride react?

12. C2H4 + 3O2 � 2CO2 + 2H2O Determine the mass of water produced if 50 g C2H4 and 50 g O2 react.

13. Consider the balanced equation Zn + 2HCl → ZnCl2 + H2 How many moles of ZnCl2 will be produced if 7

moles of HCl are used?

a. 2 moles

b. 2.5 moles

c. 3.5 moles

d. 4 moles

14. Given : C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(g) Is this chemical equation balanced?

a. True b. False

15. In the reaction below, how many moles of oxygen gas is produced by the decomposition of 4 moles of

mercury (II) oxide? 2HgO → 2Hg + O2

a. 1 mole

b. 2 moles

c. 3 moles

d. 4 moles

16. True or False, 6 moles of H2 is needed to completely react with 2 moles of N2 in the balanced chemical

reaction N2 + 3H2 → 2NH3

a. True b. False

17. If 18.0 grams of carbon are burned in 55.0 grams of oxygen, how many grams of carbon dioxide are

formed?

a. 44.01 grams CO2

b. 75.6 grams CO2

c. 151 grams CO2

d. 66.0 grams CO2

18. How many moles of Al2O3 are formed when a mixture of 0.36 moles Al and 0.36 moles O2 is ignited?

a. 0.12

b. 0.18

c. 0.28

d. 0.46

e. 0.72

19. A mass of 21.5 grams of calcium hydroxide reacts with an excess of phosphoric acid. What mass of

calcium phosphate could be recovered from solution?

a. 284 grams

b. 186 grams

c. 94.7 grams

d. 31.6 grams

e. 326 grams


69

20. If one mole of the rocket fuel ammonium perchlorate, NH4ClO4 (s) is allowed to react with excess Al so all

of the NH4ClO4 is consumed, how many molecules of water will be produced?

3NH4ClO4 (s) + 3Al (s) � Al2O3 (s) + AlCl3 (s) + 3NO (g) + 6H2O (g)

a. 3.61 x 1023

b. 1.0 x 1023

c. 6.02 x 1023

d. 1.20 x 1024

e. 3.01 x 1024

21. How many grams of phosphorous trichloide, PCl3, is produced from 93.0 grams of P4 (s) and 213 g of Cl2

(g), assuming the reaction goes to completion? The balanced equation for the reaction is:

P4 (s) + 6Cl2 (g) � 4PCl3 (g)

a. 277 g

b. 416 g

c. 213 g

d. 104 g

e. 69.3 g

22. In the oxidation of ethane: 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O how many moles of O2 are required to react

with 1 mole of ethane?

a. 7 moles b. 2 moles c. 3.5 moles

23. In the reaction: 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O how many moles of CO2 are formed when 1 mole of O2

is consumed?

a. 7 moles b. 1.75 moles c. 0.57 moles

24. In the reaction: 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O how many moles of CO2 are formed when 5moles of

ethane are consumed?

a. 10 moles b. 4 moles c. 2 moles

25. How many liters of H2 at STP are required to react with 2.3 g of Fe3O4? H2 + Fe3O4 � FeO + H2O

a. 0.22 L b. 0.44 L c. 0.56 L

26. When 0.05 mole H2 is mixed with 0.05 mole CO, what is the maximum number of moles of methanol

(CH3OH) that can be obtained? H2 + CO � CH3OH

a. 0.10 mole b. 0.05 mole c. 0.025 mole

Unit 8- Gas Laws

70

THE GAS LAWS

The gas laws describe how gases behave. They can be predicted by theory and the amount of change can be

calculated with mathematical equations. One ____________________________ is equal to 760 mm Hg, 760 torr,

or _______________ kPa (kilopascals).

• Perform the following pressure conversions.

a) 144 kPa = ______________ atm b) 795 mm Hg = ______________ atm

c) 669 torr = ______________ kPa d) 1.05 atm = ______________ mm Hg

Air pressure at higher altitudes, such as on a mountaintop, is slightly ______________________ than air pressure at

sea level. Air pressure is measured using a ________________________. More molecules mean more

____________________ between the gas molecules themselves and more collisions between the gas molecules and

the walls of the container. Number of molecules is ________________________ proportional to pressure.

Doubling the number of gas particles in a basketball _____________________ the pressure. Gases naturally move

from areas of high pressure to ____________ pressure because there is empty space to move in. If you double the

number of molecules, you _____________________ the pressure. As you remove molecules from a container, the

pressure ________________________ until the pressure inside equals the pressure outside. In a smaller container,

molecules have less room to move. The molecules hit the sides of the container _________________ often, striking

a smaller area with the same force. As volume decreases, pressure increases. Volume and pressure are

______________________ proportional. As the pressure on a gas increases, the volume decreases. Raising the

temperature of a gas increases the _______________________ if the volume is held constant. At higher

temperatures, the particles in a gas have greater ________________________ energy. They move faster and collide

with the walls of the container more often and with greater ___________________, so the pressure rises. If you

start with 1 liter of gas at 1 atm pressure and 300 K and heat it to 600 K, one of 2 things happens. Either the

volume will increase to 2 liters at ______ atm, or the pressure will increase to ______ atm while the volume

remains constant.

Ideal Gases and the Kinetic Molecular Theory

In this unit we will assume the gases behave ideally. _____________________ gases do not really exist, but this

makes the math easier and is a close approximation. Gas particles are much smaller than the spaces between them.

The particles have negligible _____________________. There are no attractive or repulsive

___________________ between gas molecules. Gas particles are in constant, _________________________

motion. Until they bump into something (another particle or the side of a container), particles move in a straight line.

No kinetic energy is ____________________ when gas particles collide with each other or with the walls of their

container. All gases have the same ______________________ energy at a given temperature. Temperature is a

measure of the average kinetic energy of the particles in a sample of matter. There are no gases for which this is

true. Real gases behave more ideally at ________________ temperature and _________________ pressure. At low

temperature, the gas molecules move more _______________________, so attractive forces are no longer

Unit 8- Gas Laws

71

negligible. As the pressure on a gas increases, the molecules are forced closer together and

_________________________ forces are no longer negligible. Therefore, real gases behave more ideally at high

temperature and low pressure.

Avogadro’s Law

Avogadro’s law states that equal volumes of different gases (at the same temperature and pressure) contain equal

numbers of ________________ or molecules. 2 liters of helium has the same number of particles as ______ liters

of oxygen. The molar volume for a gas is the volume that one mole occupies at 0.00ºC and 1.00 atm.

1 mole = 22.4 L at STP (standard temperature and pressure). As a result, the volume of gaseous reactants and

products can be expressed as small whole numbers in reactions.

• How many moles are in 45.0 L of a gas at STP?

• How many liters are in 0.636 moles of a gas at STP?

The volume of a gas is directly proportional to the number of moles.

2

2

1

1

n

V

n

V=

• Consider two samples of nitrogen gas. Sample 1 contains 1.5 mol and has a volume of 36.7 L. Sample 2

has a volume of 16.5 L at the same temperature and pressure. Calculate the number of moles of nitrogen in

sample 2.

• If 0.214 mol of argon gas occupies a volume of 652 mL at a particular temperature and pressure, what

volume would 0.375 mol of argon occupy under the same conditions?

• If 46.2 g of oxygen gas occupies a volume of 100. L at a particular temperature and pressure, what volume

would 5.00 g of oxygen gas occupy under the same conditions?

Boyle’s Law

At Boyle’s law states that the pressure and volume of a gas at constant temperature are inversely

proportional. Inversely proportional means as one goes up the other goes ________________.

P1 V1 = P2 V2

• Sketch the PV graph that represents Boyle’s law.

• A balloon is filled with 25 L of air at 1.0 atm pressure. If the pressure is changed to 1.5 atm, what is the

new volume? (Make sure the pressure and volume units in the question match.)

• A balloon is filled with 73 L of air at 1.3 atm pressure. What pressure is needed to change the volume to

43 L?

• A gas is collected in a 242 cm3 container. The pressure of the gas in the container is measured and

determined to be 87.6 kPa. What is the volume of this gas at standard pressure?

• A gas is collected in a 24.2 L container. The pressure of the gas in the container is determined to be 756

mm Hg. What is the pressure of this gas if the volume increases to 30.0 L?

Unit 8- Gas Laws

72

Charles’ Law

The volume of a gas is directly proportional to the Kelvin temperature if the pressure is held constant.

2

2

1

1

T

V

T

V=

K = °C + 273

• Sketch the PV graph that represents Charles’ law.

The V-T graph for Charles’ law results in a _____________________________ _________________ because

pressure and volume are directly proportional.

• What is the temperature of a gas that is expanded from 2.5 L at 25 ºC to 4.1 L at constant pressure? (Make

sure the volume units in the question match and make sure to convert degrees Celsius to Kelvin.)

• What is the final volume of a gas that starts at 8.3 L and 17 ºC and is heated to 96 ºC?

• A 225 cm3 volume of gas is collected at 57 ºC. What volume would this sample of gas occupy at standard

temperature?

• A 225 cm3 volume of gas is collected at 42 ºC. If the volume is decreased to 115 cm

3, what is the new

temperature?

Gay-Lussac’s Law

The temperature and the pressure of a gas are directly related at constant volume.

2

2

1

1

T

P

T

P=

• Sketch the PT graph that represents Gay-Lussac’s law.

• What is the pressure inside a 0.250 L can of deodorant that starts at 25 ºC and 1.2 atm if the temperature is

raised to 100 ºC? Volume remains constant. (Make sure the pressure units in the question match and make

sure to convert degrees Celsius to Kelvin.)

• A can of deodorant starts at 43 ºC and 1.2 atm. If the volume remains constant, at what temperature will

the can have a pressure of 2.2 atm?

• A can of shaving cream starts at 25 ºC and 1.30 atm. If the temperature increases to 37 ºC and the volume

stays constant, what is the pressure of the can?

• A 12 ounce can of a soft drink starts at STP. If the volume remains constant, at what temperature will the

can have a pressure of 2.20 atm?

The Combined Gas Law

The gas laws may be combined into a single law, called the combined gas law, which relates two sets of conditions

of pressure, volume, and temperature by the following equation.

2

22

1

11

T

VP

T

VP=

Unit 8- Gas Laws

73

• A 15 L cylinder of gas at 4.8 atm pressure at 25 ºC is heated to 75 ºC and compressed to 17 atm. What is

the new volume?

• If 6.2 L of gas at 723 mm Hg at 21 ºC is compressed to 2.2 L at 4117 mm Hg, what is the temperature of

the gas?

• A sample of nitrogen monoxide has a volume of 72.6 mL at a temperature of 16 °C and a pressure of

104.1 kPa. What volume will the sample occupy at 24 °C and 99.3 kPa?

• A hot air balloon rises to an altitude of 7000 m. At that height the atmospheric pressure drops to

300 mm Hg and the temperature cools to -33 °C. Suppose on the hot air balloon there was a small balloon

filled to 1.00 L at sea level and a temperature of 27 °C. What would its volume ultimately be when it

reached the height of 7000 m?

Dalton’s Law of Partial Pressures

Dalton’s law of partial pressures states that the _________________ pressure of a mixture of gases is equal to the

sum of the pressures of all the gases in the mixture, as shown below.

Pt = P1 + P2 + P3 + … Pt = total pressure

The partial pressure is the contribution by that gas.

• What is the total pressure in a balloon filled with air if the pressure of the oxygen is 170 mm Hg and the

pressure of nitrogen is 620 mm Hg?

• In a second balloon the total pressure is 1.30 atm. What is the pressure of oxygen (in mm Hg) if the

pressure of nitrogen is 720 mm Hg?

• A container has a total pressure of 846 torr and contains carbon dioxide gas and nitrogen gas. What is the

pressure of carbon dioxide (in kPa) if the pressure of nitrogen is 50 kPa?

• When a container is filled with 3 moles of H2, 2 moles of O2 and 4 moles of N2, the pressure in the

container is 8.7 atm. The partial pressure of H2 is _____.

It is common to synthesize gases and collect them by displacing a volume of ________________.

• Hydrogen was collected over water at 21°C on a day when the atmospheric pressure is 748 torr. The

volume of the gas sample collected was 300 mL. The vapor pressure of water at 21°C is 18.65 torr.

Determine the partial pressure of the dry gas.

• A sample of oxygen gas is saturated with water vapor at 27ºC. The total pressure of the mixture is

772 mm Hg and the vapor pressure of water is 26.7 mm Hg at 27ºC. What is the partial pressure of the

oxygen gas?

The Ideal Gas Law

Remember ideal gases do not exist. Molecules do take up ______________________. There are

_________________________ forces; otherwise, there would be no liquids.

P V = n R T

Unit 8- Gas Laws

74

Pressure times volume equals the number of ___________________ (n) times the ideal gas constant (R) times the

temperature in Kelvin.

� R = 0.0821 (L atm)/(mol K) or R = 8.314 (L kPa)/(mol K) or R = 62.4 (L mm Hg)/(mol K)

The one you choose depends on the unit for pressure!

• How many moles of air are there in a 2.0 L bottle at 19 ºC and 747 mm Hg?

• What is the pressure in atm exerted by 1.8 g of H2 gas exerted in a 4.3 L balloon at 27 ºC?

• Sulfur hexafluoride (SF6) is a colorless, odorless and very unreactive gas. Calculate the pressure (in atm)

exerted by 1.82 moles of the gas in a steel vessel of volume 5.43 L at 69.5 ºC.

• Calculate the volume (in liters) occupied by 7.40 g of CO2 at STP.

• A sample of nitrogen gas kept in a container of volume 2.30 L and at a temperature of 32 ºC exerts a

pressure of 476 kPa. Calculate the number of moles of gas present.

• A 1.30 L sample of a gas has a mass of 1.82 g at STP. What is the molar mass of the gas?

• Calculate the mass of nitrogen gas that can occupy 1.00 L at STP.

Homework / Practice

1. Identify whether the descriptions below describe an ideal gas or a real gas.

a) Gas particles move in straight lines until they collide with other particles or the walls of their

container.

b) Individual gas particles have a measurable volume.

c) The gas will not condense even when compressed or cooled.

d) Collisions between molecules are perfectly elastic.

e) Gas particles passing close to one another exert an attraction on each other.

2. Explain the following using the kinetic-molecular theory:

a) As a gas is heated, its rate of effusion through a small hole increases if all other factors remain

constant.

b) A strong-smelling gas released from a container in the middle of a room is soon detected in all

areas of the room.

3. Does atmospheric pressure increase or decrease as altitude above sea level increases?

4. Convert the following:

a. 0.200 atm = _____ mm Hg

b. 790 mm Hg = _____ Pa

c. 123 kPa = _____ atm

d. 0.935 atm = ______ torr

5. A 24 L sample of a gas (at fixed mass and constant temperature) exerts a pressure of 3.0 atm. What

pressure will the gas exert if the volume is changed to 16 L?

6. A common laboratory system to study Boyle’s law uses a gas trapped in a syringe. The pressure in the

system is changed by adding or removing identical weights on the plunger. The original gas volume is 50.0

mL when two weights are present. Predict the new gas volume when 4 more weights are added.

7. Helium gas in a balloon occupies 2.40 L at 400. K. What volume will it occupy at 300 K?

8. If 26.5 g of oxygen gas occupies a volume of 100. L at a particular temperature and pressure, how many

moles of oxygen gas will there be in 350. L under the same conditions?

Unit 8- Gas Laws

75

9. A bicycle tire is inflated to 55 lb/in2 at 15 °C. Assume that the volume of the tires does not change

appreciably once it is inflated.

i. The tire and the air inside it are heated to 30 °C by road wear, does the pressure in the tire

increase or decrease?

ii. Because the temperature has doubled, does the pressure double to 110 psi? Why or why

not?

10. At one point in the cycle of a piston in an automobile engine, the volume of the trapped fuel mixture is

400 cm3 at a pressure of 1.0 atm and a temperature of 27 °C. In the compression of the piston, the

temperature reaches 77 °C and the volume decreases to 50.0 cm3. What is the new pressure?

11. A gas storage tank has a volume of 3.5 x10 5 m3 when the temperature is 27 °C and the pressure is 1.0 atm.

What is the new volume of the tank if the temperature drops to - 10.0°C and the pressure drops to 0.95

atm?

12. A container holds three gases: oxygen, carbon dioxide, and helium. The partial pressures of the three gases

are 2.00 atm, 3.00 atm, and 4.00 atm, respectively. What is the total pressure inside the container?

13. A gas occupies 11.2 liters at 0.860 atm. What is the pressure if the volume becomes 15.0 L?

14. How much will the volume of 75.0 mL of neon change if the pressure is lowered from 50.0 torr to 8.00

torr?

15. Calculate the decrease in temperature when 2.00 L at 20.0 °C is compressed to 1.00 L.

16. What change in volume results if 60.0 mL of gas is cooled from 33.0 °C to 5.00 °C?

17. A gas occupies 1.00 L at standard temperature. What is the volume at 333.0 °C?

18. If a gas is cooled from 323.0 K to 273.15 K and the volume is kept constant what final pressure would

result if the original pressure was 750.0 mm Hg?

19. If a gas in a closed container is pressurized from 15.0 atmospheres to 16.0 atmospheres and its original

temperature was 25.0 °C, what would the final temperature of the gas be?

20. At conditions of 785.0 torr of pressure and 15.0 °C temperature, a gas occupies a volume of 45.5 mL. What

will be the volume of the same gas at 745.0 torr and 30.0 °C?

21. A gas sample occupies 3.25 liters at 24.5 °C and 1825 mm Hg. Determine the temperature at which the gas

will occupy 4250 mL at 1.50 atm.

22. If 10.0 liters of oxygen at STP are heated to 512 °C, what will be the new volume of gas if the pressure is

also increased to 1520.0 mm Hg?

23. How many moles of gas are contained in 890.0 mL at 21.0 °C and 750.0 mm Hg pressure?

24. Calculate the volume 3.00 moles of a gas will occupy at 24.0 °C and 762.4 mm Hg.

25. How many moles of a gas would be present in a gas trapped within a 37.0 liter vessel at 80.00°C at a

pressure of 2.50 atm?

26. Find the volume of 2.40 mol of gas whose temperature is 50.0 °C and whose pressure is 2.00 atm.

27. Determine the number of moles of Krypton contained in a 3.25 liter gas tank at 5.80 atm and 25.5 °C. If the

gas is Oxygen instead of Krypton, will the answer be the same? Why or why not?

Unit 8- Gas Laws

76

Gas Laws Practice Test

1. Explain how the temperature is related to the kinetic energy and motion of gas particles.

2. If the volume of a gas contained within balloon were to be tripled, what would be the impact upon the pressure

if Kelvin temperature is maintained as constant?

Directions: Solve the following problems. Show all your work, including units

3. In a mixture of carbon dioxide, oxygen gas, sulfur dioxide and carbon monoxide, the pressure of the carbon

dioxide is 0.3 atm, oxygen gas is 0.5 atm, sulfur dioxide is 0.6 atm, and the pressure of the carbon monoxide is

0.1 atm. What is the total pressure in the container?

4. A high-altitude balloon contains 4 Liters of helium gas at 1.35 atm. What is the volume when the balloon rises

to an altitude where the pressure is only 1.20 atm? (Assume that the temperature remains constant.)

5. If a sample of gas occupies 27 Liters at 12 Celsius, what will be its volume at 112 Celsius if the pressure does

not change?

6. A gas has a pressure of 122 kPa at -6 Celsius (negative 6). What will be the pressure at 85 Celsius if the

volume does not change?

7. A gas at 10 kPa and 45 Celsius occupies a container with an initial volume of 4 Liters. By changing the

volume, the pressure of the gas increases to 25 kPa as the temperature is raised to 190 Celsius. What is the new

volume?

8. You fill a rigid steel cylinder that has a volume of 840 milliliters with oxygen gas to a final pressure of

1.1 atmospheres at 145 Celsius. How many moles of nitrogen gas does the cylinder contain?

9. What is the temperature when 4 moles of carbon dioxide occupies a 2 L container and exerts a pressure of

745 torr?

10. What pressure, in atm, will be exerted by 1.25 moles of a gas at 39 Kelvin if it is contained in a 5 Liter vessel?

11. What volume will 29 grams of nitrogen gas occupy at 10 Celsius and a pressure of 620 torr?

12. A 35mL sample of hydrogen gas is collected over water at a temperature of 24 oC, the vapor pressure of the

water at that temperature is 2.99 kPa, and the atmospheric pressure is 765.5 torr. What is the pressure of the dry

hydrogen gas?


13. As the pressure of a gas at 2 atm is changed to 1 atm at constant temperature, the volume of the gas

a. decreases b. increases c. remains the same

14. According to the kinetic molecular theory, molecules increase in kinetic energy when they

a. Are mixed with other molecules at lower temperature

b. Are frozen into a solid

c. Are condensed into a liquid

d. Are heated to a higher temperature

15. Collide with each other in a container at lower temperature At STP, 32.0 liters of O2 contain the same number

of molecules as

a. 22.4 L Ar

b. 28.0 L of N2

c. 32. 0 L of H2

d. 44.8 L of He

Unit 8- Gas Laws

77

16. An 8.25 L sample of oxygen is collected at 25°C and 1.022 atm pressure. What volume will the gas occupy

0.940 atm and -15°C?

a. 1.78 L

b. 5.00 L

c. 10.4 L

d. 8.76 L

e. 7.77 L

17. A motorist fills his car tires to 32 lb/in2 pressure at a temperature of 30°C. Assuming no change in volume,

what will be the pressure in the tires when the motorist drives across Death Valley, with a pavement

temperature of 78°C?

a. 12 lb/in2

b. 28 lb/in2

c. 37 lb/in2

d. 4.8 lb/in2

e. 83 lb/in2

18. The mass of 2.37 liters of a gas is 8.91 grams. What is the density of the gas?

a. 3.76 g/L

b. 6.54 g/L

c. None of these

d. 0.266 g/L

e. 21.1 g/L

19. If temperature is constant, the relationship between pressure and volume is

a. Direct b. inverse

20. A 268 cm3 sample of an ideal gas at 18°C and 748 torr pressure is placed in an evacuated container of volume

648cm3. To what centigrade temperature must the assembly be heated so that the gas will fill the whole

chamber at 748 torr?

a. 431°C

b. 120°C

c. 704°C

d. 597°C

e. 324°C

21. How big a volume of dry oxygen gas at STP would you need to take to get the same number of oxygen

molecules as there are hydrogen molecules in 25.0 liters at 0.850 atm and 35°C

a. 18.8 L

b. 0.068 L

c. 0.656 L

d. 4.2 L

e. 32.3 L

22. Nitrogen has a molar mass of 28.02 g/mol. What is the density of nitrogen at 1.05 atm and 37°C?

a. None of these

b. 2.82 g/L

c. 0.89 g/L

d. 1.25 g/L

e. 4.72 g/cm3

23. How many moles of gas would it take to fill an average man's lungs, total capacity of which is about 4.5 liters?

Assume 1.00 atm pressure and 37.0°C.

a. 37.0 mol

b. 1.24 mol

c. 0.75 mol

d. 0.18 mole

e. 11.2 mol

24. Which flask contains the greatest number of molecules?

a. Flask 3 (O2)

b. Flask 1 (NH3)

c. Flask 2 (CH4)

d. Flasks 2 and 3

e. All are the same

25. If pressure is constant, the relationship between temperature and volume is

a. Direct b. Inverse

26. If pressure of a gas is increased and its volume remains constant, what will happen to its temperature?

a. Increase b. Decrease c. Stay the same

Unit 8- Gas Laws

78

27. One way to increase pressure on a gas is to

a. decrease temperature

b. increase volume

c. increase the number of gas particles

d. lower the kinetic energy of the gas molecules

28. If a gases volume is decreased and pressure is constant, its temperature will


29. How do gas particles respond to an increase in volume?

a. increase in kinetic energy and decrease in temperature

b. decrease in kinetic energy and decrease in pressure

c. increase in temperature and increase in pressure

d. increase in kinetic energy and increase in temperature

30. If the temperature of a gas remains constant but pressure is decreased, the volume will


31. Convert 2.3 atm into mmHg

a. 2300 mmHg

b. 1750 mmHg

c. 2.3 mmHg

d. 0.0030 mmHg

32. The pressure of a gas is 750.0 torr when its volume is 400.0 mL. Calculate the pressure (in atm) if the gas is

allowed to expand to 600.0 mL at constant temperature.

a. 0.660 atm

b. 1.48 atm

c. 500.0 atm

d. 1125 atm

33. The volume of a gas is increased from 150.0 mL to 350.0 mL by heating it. If the original temperature of the

gas was 25.0 °C, what will its final temperature be (in °C)?

a. - 146°C

b. 10.7°C

c. 58.3°C

d. 422°C

e. 695°C

34. Standard temperature and pressure (STP) refers to which conditions?

a. 0 oC and 1 kPa

b. 0 oC and 1 mm Hg

c. 0 K and 1 kPa

d. 0 K and 1 atm

e. 273 K and 1 atm

35. If 4 moles of a gas are added to a container that already holds 1 mole of gas, how will the pressure change

within the container? (Assume volume and temperature are constant.)

a. The pressure will be 5 times as great.

b. The pressure will be 2 times as great.

c. The pressure will be 4 times as great.

d. The pressure will not change.

e. None of the above are correct.

36. A 4.0 L sample of hydrogen gas at 700 mm Hg would occupy what volume at 250 mm Hg? (Assume

temperature and number of particles stays constant.)

a. 1.4 x 10 -7

L

b. 1.4 L

c. 11.2 L

d. 2.4 L

e. 7.0 x 10 5 L

37. A 25 L tank of oxygen under a pressure of 80. atm would require what pressure to decrease the volume to

1.0 L? (Assume temperature and number of particles stays constant.)

a. 0.31 atm b. 3.2 atm c. 2000 atm

d. There is not enough information to answer the question.

e. None of these is correct.

Unit 8- Gas Laws

79

38. A balloon containing 2.50 L of gas at 1 atm would be what volume at a pressure of 300 kPa? (Assume

temperature and number of particles stays constant.)

a. 6.33 L

b. 8.11 L

c. 0.844 L

d. 120. L

e. 000833 L

39. A syringe containing 75.0 mL of air is at 298 K. What will the volume of the syringe be if it is placed in a

boiling water bath (373 K). Assume pressure and the number of particles are held constant.

a. 59.9 mL b. 188 mL c. 300. mL

d. 8.34 x 106 mL e. None of the above are correct.

40. A gas occupies 40.0 mL at 127 oC. What volume will it occupy at -73

oC? (Assume pressure and number of

particles is constant.)

a. 182 mL b. 8.80 mL c. 80.0 mL

d. 20.0 mL e. None of these is correct

41. If 88.0 grams of solid carbon dioxide evaporates, how many liters of CO2 gas will be formed at a temperature of

300 K and 2.00 atmospheres of pressure?

a. 98.5 liters

b. 2170 liters

c. 24.6 liters

d. 1080 liters

42. Which of the following equations correctly combines Boyle's and Charles' Laws?

a. 2211 VPVP =

b. 2211 VTVT =

c.

2

22

1

11

T

VP

T

VP=

d.

2

22

1

11

V

TP

V

TP=

e.

2

22

1

11

P

VT

P

VT=

43. A 50.0 mL sample of a gas is at 3.00 atm of pressure and a temperature of 298 K. What volume would the gas

occupy at STP?

a. 0.00728 mL

b. 15.3 mL

c. 18.2 mL

d. 137 mL

e. None of these is

correct.

44. A syringe contains 60.0 mL of air at 740 mm Hg pressure and 20 oC. What would be the temperature at which

the syringe would contain 30.0 mL at a pressure of 370 mm Hg? (Assume no gas could leak in or out of the

syringe.)

a. -200 oC

b. 0.0137 oC

c. 5 oC

d. 73.3 oC

e. None of these is

correct

45. A sealed container contains 1.0 mol of hydrogen and 2.0 moles of nitrogen gas. If the total pressure in the

container is 1.5 atm, what is the amount of pressure exerted by each gas?

a. H2 = 1.0 atm and N2 = 0.50 atm

b. H2 = 0.50 atm and N2 = 1.0 atm

c. H2 = 1.0 atm and N2 = 2.0 atm

d. H2 = 2.0 atm and N2 = 1.0 atm

e. There is not enough information given to answer the question.

46. A sample of gas is collected by water displacement. The atmospheric pressure in the room is 757 mm Hg and

the vapor pressure of water is 17 mm Hg. What is the partial pressure of hydrogen under these conditions?

a. 17 mm Hg

b. 740 mm Hg

c. 757 mm Hg

d. 774 mm Hg

e. You cannot answer this question because you do not know the temperature.

Unit 9- Solids, Liquids and Phase Changes

80

SOLIDS AND LIQUIDS

States of Matter

There are ______ states of matter. A solid is a form of matter that has its own definite _____________ and volume.

A solid cannot _________________. The particles can vibrate but cannot move around. The particles of matter in a

solid are very tightly ____________________; when heated, a solid expands, but only slightly. A liquid is a form of

matter that flows, has ____________________ (definite) volume, and takes the _________________ of its

container. The particles in a liquid are not rigidly held in place and are _______________ closely packed than are

the particles in a solid; liquid particles are able to move past each other. A liquid is not very

__________________________. Like solids, liquids tend to expand when heated. A gas is a form of matter that

flows to conform to the ____________________ of its container and fills the entire _______________________ of

its container. Compared to solids and liquids, the particles of gases are very far apart. Because of the significant

amount of space between particles, gases are easily compressed. _____________________ is composed of

electrons and positive ions at temperatures greater than ____________ °C. The sun and other stars are examples of

plasma.

• Identify the following as a property of a solid, liquid or gas. The answer may include more that one state of

matter.

1. flows and takes the shape of a container

2. compressible

3. made of particles held in a specific arrangement

4. has definite volume

5. always occupies the entire space of its container

6. has a definite volume but flows

The word_____________________ refers to the gaseous state of a substance that is a solid or a liquid at room

temperature. For example, steam is a vapor because at room temperature water exists as a liquid. Some substances

are described as _______________________, which means that they change to a gas easily at room temperature.

Alcohol and gasoline are ______________ volatile than water. Kinetic-molecular theory predicts the constant

motion of the liquid particles. Individual liquid molecules do not have fixed positions in the liquid. However,

forces of ________________________ between liquid particles limit their range of motion so that the particles

remain closely packed in a fixed volume. These attractive forces are called ___________________________

forces. Inter = between. Molecular = molecules. A liquid diffuses more _______________________ than a gas at

the same temperature, however, because intermolecular attractions interfere with the flow.

__________________________ is a measure of the resistance of a liquid to flow. Viscosity decreases with

________________________ temperature. Particles in the middle of the liquid can be attracted to particles above

them, below them, and to either side. For particles at the surface of the liquid, there are no attractions from above to

balance the attractions from _______________. Thus, there is a net attractive force pulling down on particles at the

surface. _____________________ ____________________ is a measure of the inward pull by particles in the


81

interior. Soaps and detergents decrease the surface tension of water by disrupting the _______________________

bonds between water molecules. For a substance to be a solid rather than a liquid at a given temperature, there must

be strong attractive forces acting between particles in the solid. These forces limit the motion of the particles to

__________________________ around fixed locations in the solid. Thus, there is more order in a solid than in a

liquid. The particles can only vibrate and revolve in place. Because of this order, solids are much less

_________________ than liquids and gases. In fact, solids are not classified as fluids. Most solids are more

_________________ than most liquids. A crystalline solid is a solid whose atoms, ions, or molecules are arranged

in an orderly, geometric, three-dimensional structure. Most solids are _____________________. Amorphous solids

lack an orderly internal structure. Think of them as __________________________ liquids. Examples of

amorphous solids include ____________________ and glass.

Phase Changes

If a substance is usually a liquid at room temperature (as water is), the gas phase is called a _________________.

Vaporization is the process by which a liquid changes into a gas or vapor. Vaporization is an endothermic process -

it requires _______________. When vaporization occurs only at the _____________________ of an uncontained

liquid (no lid on the container), the process is called evaporation. Molecules at the surface break away and become

gas. Only those with enough _____________________ energy (KE) escape. Evaporation is a

_______________________ process. It requires heat, which is endothermic. __________________ pressure is the

pressure exerted by a vapor over a liquid. As temperature increases, water molecules gain kinetic energy and vapor

pressure ______________________. When the vapor pressure of a liquid equals atmospheric pressure, the liquid

has reached its boiling point, which is 100°C for water at sea level. Recall that standard atmospheric pressure equals

______ atm. At this point, molecules throughout the liquid have the energy to enter the gas or vapor phase. The

temperature of a liquid can never ______________ above its boiling point. Boiling is an

__________________________ process. It requires the addition of heat. As you go up into the mountains (increase

in elevation), atmospheric pressure ______________. Lower external pressure requires ______________________

vapor pressure. Lower vapor pressure means lower ______________________ point. As a result, spaghetti cooks

slower in the mountains than at the beach. When you use a pressure cooker to can vegetables, the external pressure

around the mason jars rises. This raises the vapor pressure needed in order for water to boil. In turn, the boiling

point is raised so the food cooks ______________________.

Some phase changes release energy into their surroundings. For example, when a vapor loses energy, it may change

into a __________________. Condensation is the process by which a gas or vapor becomes a liquid. It is the

___________________ of vaporization. In a closed system, the rate of vaporization can equal the rate of

condensation. When first sealed, the molecules gradually _________________ the surface of the liquid. As the

molecules build up above the liquid, some condense back to a liquid. Equilibrium is reached when the rate of

vaporization __________________ the rate of condensation. Molecules are constantly changing phase. The total

amount of liquid and vapor remains _______________________.


82

The melting point of a solid is the temperature at which the ____________________ holding the particles together

are broken and the solid becomes a liquid. When heated the particles vibrate more _____________________ until

they shake themselves free of each other. The freezing point is the temperature at which a liquid becomes a

_________________________ solid. The freezing point is the _______________ as the melting point. The process

by which a solid changes directly into a gas without first becoming a liquid is called _______________________.

Solid air fresheners and dry ice are examples of solids that sublime. When a substance changes from a gas or vapor

directly into a solid without first becoming a liquid, the process is called _________________________. Deposition

is the reverse of sublimation. _______________ is an example of water deposition.

• Classify the following phase changes.

1. dry ice (solid carbon dioxide) to carbon dioxide gas ____________________________

2. ice to liquid water ________________________________

3. liquid water to ice ________________________________

4. water vapor to liquid water ________________________________

Phase Diagrams

Temperature and _____________________ control the phase of a substance. A phase diagram is a graph of

pressure versus temperature that shows in which phase a substance exists under different conditions of temperature

and pressure. A phase diagram typically has ______ regions, each representing a different phase and three curves

that ________________________ each phase.

The points on the curves (lines) indicate conditions under which two phases coexist. The critical point indicates the

critical pressure and the critical temperature above which a substance cannot exist as a ____________________.

The triple point is the point on a phase diagram that represents the temperature and pressure at which three phases of

a substance can __________________________. The __________________________ slope of the solid-liquid line

in the phase diagram for water indicates that the solid floats on its liquid.

• What happens to solid CO2 at -100 ºC and 1 atm pressure as it is heated to room temperature?

0.0098 Temperature (°C)


83

• What happens to water at 1 atm as the temperature rises from -15°C to 60°C?

• What state of matter is water at 50°C and 20 atm?

• At what temperature does the triple point occur for water?

• At what temperature does the critical point occur for carbon dioxide?

• At standard pressure and -78°C, what phase change occurs for carbon dioxide?

• What state of matter is carbon dioxide at -80°C and 2 atm?

Solids and Liquids Practice Test

Directions: Identify the proper sections by indicating the interval between 2 letters.

1. Which section

represents the gas

being warmed?

_____ to _______

2. Which section

represents a phase

change from solid to

liquid?

_____ to _______

3. Define viscosity-

4. The temperature at

which the vapor

pressure of a liquid

equals the external or atmospheric pressure is known as the________________________

Directions: Using the phase diagram below,

answer questions 5-7:

5. What does letter T represent?

_______________

6. At 0.75 atm, what is the melting point

and boiling point of the substance?

________and _________

7. At 0.25 atm, what is the freezing point?

________

Temperature (oC)

Pressure

(atm) C

0.25

0.75

125 175 250

T

A

B

C

D E

F

G

H

200º C

60º C

Heating Curve

Energy


84

Multiple Choice

8. Under the same conditions of temperature and pressure, a liquid differs from a gas because the particles of

the liquid

a. are in constant straight-line motion

b. take the shape of the container they occupy

c. have no regular arrangement

d. have stronger forces of attraction between them

9. The phase change represented by the equation I2 (s) → I2 (g) is called

a. sublimation

b. condensation

c. melting

d. boiling

10. Which of the following terms represents the temperature and pressure at which three states of a compound

can coexist

a. Law of definite composition

b. Van der Waals forces

c. Graham’s Law of Diffusion

d. Triple point

e. Critical point

11. What is the smallest portion of a crystal lattice that reveals the 3-dimensional pattern?

a. unit cell

b. crystal structure

c. coordinate system

d. crystalline symmetry

12. What forces hold nonpolar particles together?

a. magic

b. hydrogen bonding

c. London dispersion

d. dipole-dipole

13. Compared with the particles in a solid, the particles in a liquid usually are

a. higher in energy

b. closer together

c. more massive

d. less fluid

14. What is the process of a substance changing from a vapor to a solid without passing through the liquid

phase?

a. condensation

b. deposition

c. sublimation

d. evaporation

15. A liquid forms when the average energy of a solid substance's particles

a. increases

b. changes form

c. creates an orderly arrangement

d. decreases

16. Which of the following is an NOT an amorphous solid?

a. silly putty

b. play dough

c. ice

d. glass

17. Which term best describes the process by which particles escape from both the surface of a liquid and from

within the liquid itself and enter the gas phase?

a. boiling

b. evaporation

c. aeration

d. surface tension

18. The attractive forces in a solid are

a. too weak to prevent the particles from changing positions

b. strong enough to hold the particles in fixed positions

c. less effective than those in a liquid

d. weaker than those of a liquid particles


85

19. When electrons in a covalent bond spend more time around on nucleus of the compound than the other, the

molecule is considered

a. weak

b. polar

c. ionic

d. nonpolar

20. Which of the following phase changes results in an overall increase in randomness of particles over the

course of the change?

a. deposition

b. condensation

c. melting

d. freezing

21. What type of crystals are like giant molecules?

a. covalent network

b. covalent molecular

c. metallic

d. ionic

22. The difference between crystalline and amorphous solids is determined by

a. temperature changes

b. pressure when the substances are formed

c. strength of molecular forces

d. the particle arrangement

23. Which of the following statements is false?

a. Condensed states have much higher densities than gases.

b. Molecules are very far apart in gases and closer together in liquids and solids.

c. Gases completely fill any container they occupy and are easily compressed.

d. Vapor refers to a gas formed by evaporation of a liquid or sublimation of a solid.

e. Solid water (ice), unlike most substances, is denser than its liquid form (water).

24. Which physical state/ property is incorrectly matched?

a. liquids and solids - rigid shape

b. gases - easily compressed

c. gases and liquids – flow

d. solids - higher density than gases

e. liquids – incompressible

25. Which one of the following statements does not describe the general properties of liquids accurately?

a. Liquids have characteristic volumes that do not change greatly with changes in temperature.

(Assuming that the liquid is not vaporized.)

b. Liquids diffuse only very slowly when compared to solids.

c. The liquid state is highly disordered compared to the solid state.

d. Liquids have high densities compared to gases.

26. For which of the following would permanent dipole-dipole interactions play an important role in

determining physical properties in the liquid state?

a. BF3

b. ClF

c. BeCl2

d. F2

e. CCl4

27. Identify which property liquids do not have in common with solids.

a. rigid shape

b. volumes do not change significantly with pressure

c. hydrogen bonding forces can be significant

d. practically incompressible

e. volumes do not change significantly with temperature

28. Which of the following interactions are the strongest?

a. hydrogen bonding force

b. ion-ion interactions

c. dipole- dipole force

d. London-dispersion force


86

29. Which one of the following statements does not describe the general properties of solids accurately?

a. Solids have characteristic volumes that do not change greatly with changes in temperature.

b. Solids have characteristic volumes that do not change greatly with changes in pressure.

c. Solids diffuse only very slowly when compared to liquids and gases.

d. Solids are not fluid.

e. Most solids have high vapor pressures at room temperature.

30. For which of the following would dispersion forces be the most important factor in determining physical

properties in the liquid state?

a. H2O

b. NaCl

c. F2

d. HF

e. NH4Cl

31. For which of the following would hydrogen bonding not be an important factor in determining physical

properties in the liquid state?

a. HI

b. H2O

c. HF

d. NH3

e. H2O2

32. Which technique listed below separates a mixture of liquids on the basis of their boiling points?

a. Distillation

b. Extraction

c. Filtration

d. Reflux

e. None of the

above

33. The melting point of a solid is the same as the ____ of its liquid.

a. Boiling point

b. Freezing point

c. Sublimation point

d. Condensation point

e. Critical point

34. Which one of the following statements does not describe the general properties of liquids accurately?

a. In the liquid state the close spacing of molecules leads to large intermolecular forces that are

strongly dependent on the nature of the molecules involved.

b. Liquids are practically incompressible.

c. As the temperature of a liquid is increased, the vapor pressure of the liquid decreases.

d. The normal boiling point of a liquid is the temperature at which the vapor pressure of the liquid

becomes equal to exactly 760 torr.

e. Vapor pressures of liquids at a given temperature differ greatly, and these differences in vapor

pressure are due to the nature of the molecules in different liquids.

35. Some solids can be converted directly to the vapor phase by heating. The process is called ____.

a. Fusion

b. Sublimation

c. Vaporization

d. Condensation

e. Distillation

36. Which of the images shown here depicts a phase that has definite volume but not definite shape?

a. The one on

the left

b. The one in the

middle

c. The one on

the right

37. Ice floats in water because:

a. Water is denser than ice

b. Ice is colder than water

c. Water has a substantial surface

tension

d. Ice is denser than water


87

38. Which phase depicted here has both a definite shape and a definite volume?

a. The one in the middle

b. The one on the right

c. The one on the left

d. The one in the middle and the one

on the right

39. Which of the phases depicted here can be easily compressed?

a. The one in the middle

b. The one on the right

c. The one on the left

d. The one in the middle and the one

on the right

40. Which phase of matter is depicted here?

a. Liquid

b. Gas

c. Plasma

d. Solid

41. Which phase(s) depicted here have the ability to flow?

a. The one on the right

b. The one on the left

c. The ones on the right and the left

d. The one in the middle and the one on the right

e. The one in the middle

42. During the phase change from liquid to solid:

a. energy must be removed

b. energy must be absorbed

c. there is no change in energy

43. Definite shape, definite volume, and a low rate of diffusion are characteristics of:

a. Fluids

b. Liquids

c. Gases

d. Solids

44. Which phase of matter is depicted here?

a. Solid

b. Gas

c. Liquid

d. Plasma

Unit 10- Solutions and Solubility

88

SOLUTIONS

A solution is made up of a solute and a _______________________________. The solvent does the

________________________________. The solute is the substance that is dissolved. If a solution is made of two

liquids, the one in ______________________ quantity is the solute. _________________________ is the universal

solvent. Water is a versatile solvent because of its attraction to other molecules and its

___________________________. Most of the water on the Earth is not pure, but rather is present in solutions.

Table salt (NaCl), like a great many ionic compounds, is _________________________ in water. The salt solution

is also an excellent ___________________________ of electricity. This high level of electrical conductivity is

always observed when ionic compounds dissolve to a significant extent in water. The process by which the charged

particles in an ionic solid separate from one another is called _____________________________. You can represent

the process of dissolving and dissociation in shorthand fashion by the following equation.

________________________________________ Water is not only good at dissolving ionic substances. It also is a

good solvent for many _________________________________ compounds. Consider the covalent substance

sucrose, commonly known as table sugar, as an example. Although water dissolves an enormous variety of

substances, both ionic and covalent, it does not dissolve everything. The phrase that scientists often use when

predicting solubility is “________________ dissolves like.” The expression means that dissolving occurs when

similarities exist between the solvent and the solute. A salt dissolves faster if it is _________________________ or

shaken, if the particles are made ________________________ and if the temperature is ______________________.

In order to dissolve the solvent molecules must come in ______________________________ with the solute.

Stirring moves fresh _____________________________ next to the solute. The solvent touches the surface of the

solute. __________________________________ pieces increase the amount of surface of the solute. For solids in

liquids, as the temperature goes up the solubility goes ______________________. A higher temperature makes the

molecules of the solvent move around ______________________________ and contact the solute harder and

________________________ often. It speeds up dissolving. Higher temperature usually increases the

_______________________________ that will dissolve.

For gases in a liquid, as the temperature goes up the solubility goes _______________________. For gases in a

liquid, as the pressure goes up the solubility goes ______________________.

Solubility is the ________________________________ amount of substance that will dissolve at that temperature

(usually measured in grams/liter). If the amount of solute dissolved is less than the maximum that could be

dissolved, the solution is called a(n) ___________________________ solution. A solution which holds the

maximum amount of solute per amount of the solution under the given conditions is called a(n)

_____________________________ solution. A(n) _________________________________ solution contains more

solute than the usual maximum amount and are unstable. They cannot permanently hold the excess solute in solution

and may release it suddenly. A(n) __________________ crystal will make the extra come out. Generally, a


89

supersaturated solution is formed by dissolving a solute in the solution at an elevated temperature, at which

solubility is _______________________ than at room temperature, and then slowly cooling the solution.

• How many grams of potassium chlorate (KClO3) will dissolve in 100 g of water at 30ºC?

• How many grams of potassium nitrate (KNO3) will dissolve in 100 g of water at 50ºC?

• At what temperature will 90 grams of Pb(NO3)2 dissolve in 100 g of water?

• At what temperature will 50 grams of KCl dissolve in 100 g of water?

• If 45 g of KCl is dissolved in 100 g of water at 60ºC, is the solution unsaturated, saturated or

supersaturated?

• If 90 g of Pb(NO3)2 is dissolved in 100 g of water at 40ºC, is the solution unsaturated, saturated or

supersaturated?

• If 30 g of KNO3 is dissolved in 100 g of water at 20ºC, is the solution unsaturated, saturated or

supersaturated?

• If 10 g of KClO3 is dissolved in 100 g of water at 50ºC, is the solution unsaturated, saturated or

supersaturated?

___________________________ means that two liquids can dissolve in each other.

________________________________ means they cannot. Oil and ______________________ are immiscible.

Measuring Solutions

Chemists never apply the terms strong and weak to solution concentrations. Instead, use the terms concentrated and

_________________________. Concentration is a measure of the amount of solute dissolved in a certain amount of

solvent. A concentrated solution has a _________________________ amount of solute. A dilute solution has a

__________________________ amount of solute. For chemistry applications, the concentration term molarity is


90

generally the most useful. Molarity is the number of moles of _______________________ in 1 Liter of the

solution.

Liters

MolesMolarity =

Note that the volume is the total solution volume that results, not the volume of solvent alone. Suppose you need

1.0 Liter of a 1 M copper (II) sulfate solution.

STEP 1: Measure a mole of copper (II) sulfate.

STEP 2: Place the CuSO4 in a volumetric flask.

STEP 3: Add some water to dissolve the CuSO4 and then add enough additional water to bring the total volume

of the solution to 1.0 L.

• What is the molarity of a solution with 2.0 moles of NaCl in 4.0 Liters of solution?

• What is the molarity of a solution with 3.0 moles dissolved in 250 mL of solution?

• How many moles of NaCl are needed to make 6.0 L of a 0.75 M NaCl solution?

• 0.200 moles of NaOH are dissolved in a small amount of water then diluted to 500. mL. What is the

molarity?

• How many moles are in 1500 mL of a 3.2 M solution of nitric acid (HNO3)?

• 80.6 g of KCl are dissolved in a small amount of water then diluted to 500. mL. What is the molarity?

• 125 g of NaC2H3O2 are dissolved in a small amount of water then diluted to 750. mL. What is the molarity?

• How many grams of CaCl2 are needed to make 625 mL of a 2.00 M solution?

• How many grams of aluminum nitrate are needed to make 600. mL of a 0.500 M Al(NO3)2 solution?

Refer to the Figure on page 89 to answer the following questions:

• What is the molarity of a KNO3 solution at 10ºC? (100 g of water = 100 mL of water)

• What is the molarity of a NaNO3 solution at 10ºC?

• What is the molarity of a KClO3 solution at 70ºC?

Dilution

The number of moles of solute doesn’t change if you add more solvent.

M1 x V1 = M2 x V2

M1 and V1 represent the starting concentration and volume. M2 and V2 represent the ______________ concentration

and volume.

• 2.0 L of a 0.88 M solution are diluted to 3.8 L. What is the new molarity?

• 6.0 L of a 0.55 M solution are diluted to 8.8 L. What is the new molarity?

• You have 150 mL of 6.0 M HCl. What volume of 1.3 M HCl can you make?

• 6.0 liters of a 0.55 M solution are diluted to a 0.35 M solution. What is the final volume?


91

• You need 450 mL of 0.15 M NaOH. All you have available is a 2.0 M stock solution of NaOH. How do

you make the required solution?

Compounds in Aqueous Solution and Double Replacement Reactions

The _________________________________ of ions when an ionic compound dissolves in water is called

dissociation. Although no compound is completely insoluble, compounds of very low solubility can be considered

insoluble.

• Using the solubility rules printed on page 6 of the NCDPI Reference Tables for Chemistry, determine

whether the following salts are soluble in water.

a) sodium chloride

b) mercury (I) acetate

c) potassium nitrate

d) nickel carbonate

e) barium sulfate

f) ammonium bromide

In a double-replacement reaction, two compounds exchange partners with each other to produce two different

compounds. The general form of the equation is

AB + CD → AD + CB

Signs that a double-replacement reaction has taken place include a color change, the release or absorption of energy,

evolution of a gas, and formation of a _______________________________.

Homework / Practice

1. Suppose you had 2.00 moles of solute dissolved into 1.00 L of solution. What's the molarity?

2. Calculate the molarity of 25.0 grams of KBr dissolved in 750.0 mL.

3. 80.0 grams of glucose (C6H12O6, mol. wt = 180. g/mol) is dissolved in enough water to make 1.00 L of

solution. What is its molarity?

4. What is the molarity when 0.75 mol is dissolved in 2.50 L of solution

5. What is the molarity of 245.0 g of H2SO4 dissolved in 1.00 L of solution?

6. What is the molarity of 5.00 g of NaOH in 750.0 mL of solution?

7. How many moles of Na2CO3 are there in 10.0 L of 2.0 M soluton?

8. How many moles of Na2CO3 are in 10.0 mL of a 2.0 M solution?

9. How many grams of Ca(OH)2 are needed to make 100.0 mL of 0.250 M solution?

10. What is the molarity of a solution made by dissolving 20.0 g of H3PO4 in 50.0 mL of solution?

11. What weight (in grams) of KCl is there in 2.50 liters of 0.50 M KCl solution?

12. What is the molarity of a solution containing 12.0 g of NaOH in 250.0 mL of solution?

Solutions Practice Test

Directions: For credit, show all steps in your calculations and include units.

1. What is the molarity of a solution of NaOH if 12 liters of the solution contains 3 moles of NaOH?

2. You have a 3.5 L solution that contains 20 grams of NaCl. What is the molarity of the solution?


92

Directions: Using the solubility curve on page 88, answer the questions three through five.

3. Which is most soluble at 20ºC? _______

4. How many grams of KClO3 can be dissolved in 100g H2O at 90ºC? _____

5. At 40ºC, how much KCl can be dissolved in 300 g. H2O? _________

6. In Koolaid, describe what is the solute and what is the solvent.

Determine whether, according to the solubility rules, the mixing of these

substances will make a solution or a precipitate.

solution precipitate

7. Water and Mg(OH)2

8. Water and Na2CO3

Determine how the following conditions can affect the rate of dissolving KCl

in water.

Faster Slower

9. Decrease the temperature of the water

10. Agitate the mixture

12. Which of these compounds are soluble in water?

a. CaBr2

b. PbCl2

c. SrS

d. CaCO3

13. Iron (III) sulfide is soluble in water.

a. True b. False

14. LiBr is

a. Soluble

b. Insoluble

c. can't tell the solubility

d. a covalent compound

15. NH4OH is insoluble.

a. True b. False

16. Which of these compounds is soluble?

a. Pb(OH)4

b. NaHCO3

c. BaCrO4

d. Mg3(PO4)2

17. Powdered NaCl will dissolve slower then NaCl crystals because there is less surface area for the reaction to

take place.

a. True b. False

18. Which term indicates that there is a large quantity of solute, compared to the amount of solvent in a

solution

a. Dilute

b. Concentrated

c. Unsaturated

d. Saturated

19. Ten grams of sodium hydroxide is dissolved in enough water to make 1L of solution. What is the molarity

of the solution?

a. 0.25 M

b. 0.5 M

c. 1 M

d. 1.5 M

20. Which solution is the most concentrated?

a. 1 mole of solute dissolved in 1 liter of solution?

b. 2 moles of solute dissolved in 3 liters of solution?

c. 6 moles of solute dissolved in 4 liters of solution?

d. 4 moles of solute dissolved in 8 liters of solution?


93

21. What is the total number of moles of H2SO4 needed to prepare 5.0 liters of a 2.0 M solution of H2SO4?

a. 2.5

b. 5.0

c. 10

d. 20

22. What is the molarity of a KF (aq) solution containing 116 grams of KF in 1.00 liter of solution?

a. 1.00 M

b. 2.00 M

c. 3.00 M

d. 4.00 M

23. The solubility of a gas will ___ when a solution containing the gas is heated and the solubility of a gas in a

solution will ___ when the pressure over the solution is decreased.

a. decrease...decrease

b. decrease...increase

c. increase...decrease

d. increase...increase

24. How many grams of potassium nitrate are required to prepare 3.00 x 102 mL of 0.750 M solution?

a. 2.28 x 104 g

b. 84.5 g

c. 22.8 g

d. 2.4 g

e. 0.00223 g

25. How many grams of sodium chloride are dissolved in 50.0 mL of 1.50 M solution?

a. 0.00324 g

b. 117 g

c. 4.38 x 103 g

d. 23.4 g

e. 4.38 g

26. A 500 mL sample of a 0.350 M solution is left open on a lab counter for two weeks, after which the

concentration of the solution is 0.955 M. What is the new volume of the solution?

a. 183 L

b. 223 mL

c. 0.183 L

d. 1.83 mL

e. 0.605 L

27. A chemist makes a stock solution of potassium chromate solution by dissolving 97.1 grams of the

compound in 1.00 liter of solution. What volume of the solution must be diluted with water in order to

prepare 200.0 mL of 0.200 M solution?

a. 80.0 mL

b. 0.150 L

c. 750. mL

d. 0.0800 mL

e. 120. mL

28. A 25.0-g sample of sodium hydroxide is dissolved in 400. mL of water. What is the concentration of the

solution?

a. 0.10 M

b. 62.5 M

c. 1.56 M

d. 100. M

e. 1.56 x 10-3

M

29. How many milliliters of 6.0 M HNO3 are needed to prepare 500 mL of 0.50 M HNO3?

a. 0.25 mL

b. 300 mL

c. 15.76 mL

d. 40 mL

e. None of these are correct

30. How many grams of calcium chloride are needed to prepare 300 mL of a 0.250 M solution?

a. 832 g

b. 5.66 g

c. 566 g

d. 8.32 g

e. 112 g

31. In a solution of sugar and water, the solvent is the:

a. sugar b. water

32. Gases dissolve best in liquids when:

a. the pressure is high and the temperature is low

b. the pressure is low and the temperature is low

c. the pressure is low and the temperature is high

d. the pressure is high and the temperature is high


94

33. The solubility of potassium nitrate in water at 35 °C is about 60 grams KNO3 per 100 grams of water. How

many grams of KNO3 should dissolve in 300 grams of water at 35 °C?

a. 180 grams b. 335 grams c. 20 grams

34. Breaking up a solid speeds dissolving in a liquid by:

a. decreasing the pressure

b. slowing hydration

c. raising the temperature

d. increasing surface area

35. Most salts become more soluble in water as the:

a. temperature is decreased

b. pressure is decreased

c. pressure is increased

d. temperature is increased

36. Calculate the molarity of a solution made when 80 grams of NaOH is dissolved in 2 L of solution

a. 40 M NaOH

b. 82 M NaOH

c. 1 M NaOH

d. 160 M NaOH

37. Which procedure will increases the solubility of KCl in water?

a. stirring the solute and solvent mixture

b. increasing the surface area of the solute

c. raising the temperature of the solvent

d. increasing the pressure on the surface of the solvent

38. Under which conditions are gases most soluble in water?

a. high pressure and high temperature

b. high pressure and low temperature

c. low pressure and high temperature

d. low pressure and low temperature

Unit 11- Acids and Bases

95

Acids and Bases

Properties of Acids and Bases

Acids taste _________________. Lemon juice and _____________________________, for example, are both

aqueous solutions of acids. Acids conduct electricity; they are ________________________. Some are strong

electrolytes, while others are _________________ electrolytes. An acetic acid solution, which is a weak electrolyte,

contains only a few ions and does not conduct as much current as a strong electrolyte. The bulb is only

_____________________ lit. Acids cause certain colored dyes (_________________________) to change color.

(Litmus paper turns _______________.) Acids react with metals to form ______________________________ gas.

This property explains why acids corrode most metals. Acids react with hydroxides (bases) to form water and a

___________________. Bases taste _________________________ and feel ______________________________.

Bases can be strong or weak electrolytes.

Naming Acids

Acids are compounds that give off _________________________ ions (H+) when dissolved in water. Acids will

always contain one or more hydrogen ions next to an __________________________. The anion determines the

name of the acid.

Naming Binary Acids

Binary acids contain hydrogen and an anion whose name ends in –ide. When naming the acid, put the prefix

______________________- and change -ide to -ic acid.

Example: HCl The acid contains the hydrogen ion and chloride ion. Begin with the prefix hydro-, name the

nonmetallic ion and change -ide to -ic acid.

Example: H2S The acid contains the hydrogen ion and sulfide ion. Begin with the prefix hydro- and name the

nonmetallic ion. The next step is change -ide to -ic acid, but for sulfur the “ur” is added before -ic.

• Name the following binary acids.

a) HF ___________________________________________

b) H3P __________________________________________

Writing the Formulas for Binary Acids

The prefix hydro- lets you know the acid is binary. Determine whether you need to criss-cross the oxidation

numbers of hydrogen and the nonmetal.

Example: Hydrobromic acid The acid contains the hydrogen ion and the bromide ion. The two oxidation

numbers add together to get zero. The prefix hydro- lets you know the acid is binary.

Example: Hydrotelluric acid The acid contains the hydrogen ion and the telluride ion. The two oxidation

numbers do NOT add together to get zero, so you must criss-cross.

• Write the formulas for the following binary acids.

a) Hydrocarbonic acid _______________ b) Hydroselenic acid _______________


96

Naming Ternary Acids

The acid is a ternary acid if the anion has oxygen in it. The anion ends in -ate or -ite. Change the suffix -ate to -

_______ acid. Change the suffix -ite to -ous acid The hydro- prefix is NOT used!

Example: HNO3 The acid contains the hydrogen ion and nitrate ion. Name the polyatomic ion and change -ate

to -ic acid.

Example: HNO2 The acid contains the hydrogen ion and nitrite ion. Name the polyatomic ion and change -ite to

-ous acid.

Example: H3PO4 The acid contains the hydrogen ion and phosphate ion. Name the polyatomic ion and change -

ate to -ic acid.

• Name the following ternary acids.

a) H2CO3 ___________________________________________________

b) H2SO4 ___________________________________________________

c) H2CrO4 ___________________________________________________

d) HClO2 ___________________________________________________

Writing the Formulas for Ternary Acids

The lack of the prefix hydro- from the name implies the acid is ternary, made of the hydrogen ion and a polyatomic

ion. Determine whether you need to criss-cross the oxidation numbers of hydrogen and the polyatomic ion.

Example: Acetic acid The polyatomic ion must end in –ate since the acid ends in -ic. The acid is made of H+ and

the acetate ion. The two charges when added equal zero.

Example: Sulfurous acid Again the lack of the prefix hydro- implies the acid is ternary, made of the hydrogen ion

and a polyatomic ion. The polyatomic ion must end in –ite since the acid ends in -ous. The acid is made of H+ and

the sulfite ion. The two charges when added do not equal zero, so you must crisscross the oxidation numbers.

• Write the formulas for the following ternary acids.

a) perchloric acid ______________________ b) iodic acid _____________________

c) nitrous acid _______________________ d) bromic acid ___________________

Types of Acids and Bases

Arrhenius Definitions - The simplest definition is that an acid is a substance that produces _____________________

ions when it dissolves in water. A hydronium ion, H3O+, consists of a hydrogen ion attached to a

__________________ molecule. A hydronium ion, H3O+, is equivalent to H

+. HCl and H3PO4 are acids according

to Arrhenius. A base is a substance that produces ________________________ ions, OH–, when it dissolves in

water. Ca(OH)2 and NaOH are Arrhenius bases. NH3, ammonia, could not be an Arrhenius

___________________.

Monoprotic acids have only ____________ ionizable hydrogen. Some acids have more than one ionizable hydrogen

and are called ______________________________ acids.


97

Bronsted-Lowry Definitions - An Bronsted-Lowry acid is a ________________________ (H+) donor. HBr and

H2SO4 are Bronsted-Lowry acids. When a Bronsted-Lowry acid dissolves in water it gives its proton to water. HCl

(g) + H2O (l) ↔ H3O+ + Cl

- A Bronsted-Lowry base is a proton acceptor. B + H2O ↔ BH

+ + OH

- A

Brønsted-Lowry base does not need to contain OH-.

Consider HCl(aq) + H2O(l) → H3O+(aq) + Cl

-(aq) HCl donates a proton to water. Therefore, HCl is an

_______________. H2O accepts a proton from HCl. Therefore, H2O is a ______________.

• Identify the acid and base in the following reactions.

a) H2SO3 + H2O ↔ HSO3- + H3O

+

Acid _____________________________ base _________________________

b) NH3 + H2SO4 ↔ NH4+ + HSO4

-

Acid _____________________________ base _________________________

Molarity and Dilution

The concentration of a solution is the amount of solute present in a given quantity of solution.

_________________________ is the number of moles of solute in 1 liter of solution.

The procedure for preparing a less concentrated solution from a more concentrated one is called a

___________________________.

M1 V1 = M2 V2

PRACTICE:

• What is the molarity of an acetic acid (HC2H3O2) solution with 4.0 moles dissolved in 250 mL of solution?

• 3.25 moles of the base potassium hydroxide (KOH) are dissolved in a small amount of water then diluted to

725 mL. What is the concentration?

• How many moles are in 1250 mL of a 3.60 M solution of nitric acid (HNO3)?

• 6.0 L of a 1.55 M LiOH solution are diluted to 8.8 L. What is the new molarity of the lithium hydroxide

solution?

• You have 250 mL of 6.0 M HCl. How many milliliters of 1.2 M HCl can you make?

• 4.0 liters of a 0.75 M solution of sulfuric acid (H2SO4) are diluted to a 0.30 M solution. What is the final

volume?

• You need 350 mL of 0.25 M NaOH. All you have available is a 2.0 M stock solution of NaOH. How do you

make the required solution?

Strength of Acids and Bases

The strength of a base is based on the percent of units ___________________________________, not the number

of OH– ions produced. The strength of a base does NOT depend on the _____________________________. 1A

and _______ hydroxides, excluding __________, are strong bases. Some bases, such as Mg(OH)2, are not very


98

soluble in water, and they don’t produce a large number of OH– ions. However, they are still considered to be

strong bases because all of the base that does dissolve completely dissociates. The strength of an acid is based on

the percent of units dissociated, not the number of ____________ ions produced. The strength of an acid does NOT

depend on the _______________________________. There are 6 strong acids: HCl, HBr, HI, HClO4, HNO3, and

H2SO4. Strong acids and bases are strong __________________________________ because they dissociate

completely. Electrolytes conduct ______________________________. Weak acids and bases don’t completely

ionize, so they are weak electrolytes. Although the terms weak and strong are used to compare the

_____________________________ of acids and bases, dilute and concentrated are terms used to describe the

_____________________________ of solutions.

pH Scale

Water ionizes; it falls apart into _________________. H2O � H+ + OH

- The preceding reaction is called the

_____________________________________ of water. [H+ ] = [OH-] = 1 x 10

-7 M When [H+ ] = [OH-], the

solution is _________________________. At 25°C, Kw = [H+] [OH

-] = 1 x 10

-14 Kw is called the ion-product

constant. If [H+] > 10

-7 then [OH

-] < 10

-7. The solution is ______________________ when [H

+] > [OH

-]. If [H

+]

< 10-7

then [OH-] > 10

-7. The solution is __________________________ when [OH

-] > [H

+]. In most applications,

the observed range of possible hydronium or hydroxide ion concentrations spans 10–14

M to ______M. To make

this range of possible concentrations easier to work with, the pH scale was developed. pH is a mathematical scale in

which the concentration of hydronium ions in a solution is expressed as a number from _________ to __________.

pH meters are instruments that measure the exact pH of a solution. Indicators register different colors at different

pH’s. In neutral solution, pH = 7. In an acidic solution, pH < 7. In a basic solution, pH > 7. As the pH drops from

7, the solution becomes more acidic. As pH increases from 7, the solution becomes more basic.

The pH of a solution equals the negative logarithm of the hydrogen ion concentration.

pH = - log [H+]

Chemists have also defined a pOH scale to express the basicity of a solution.

pOH = - log [OH-]

If either pH or pOH is known, the other may be determined by using the following relationship.

pH + pOH = 14.00

• Find the pH of the following solutions.

a) The hydronium ion concentration equals: 10–2

M. pH = _________________

b) The hydronium ion concentration equals: 1 x 10–6

M. pH = _________________

c) The hydroxide ion concentration equals: 10–5

M. pH = _________________

d) The hydroxide ion concentration equals: 10–3

M. pH = _________________

• If a certain carbonated soft drink has a hydrogen ion concentration of 1.0 x 10–4

M, what are the pH and

pOH of the soft drink?


99

Calculating Ion Concentrations From pH

If either pH or pOH is known, the hydrogen ion or hydroxide ion can be found.

[H+] =10

-pH [OH

-] =10

-pOH

• Find the [H+] of a solution that has a pH equal to 6.

• Find the [H+] of a solution that has a pH equal to 5.

• Find the [H+] of a solution that has a pOH equal to 6.



• Find the [OH-] of a solution that has a pH equal to 10.

Calculating Ion Concentration From Ion Concentration

If either [H+] or [OH

-] is known, the hydrogen ion or hydroxide ion can be found.

[H+] [OH

-] = 1 x 10

-14

• Find the hydrogen ion concentration if the hydroxide ion concentration equals: 1 x 10–8

M.

• Find the hydroxide ion concentration if the hydrogen ion concentration equals: 1 x 10–9

M.

Indicators

Chemical _____________________ whose colors are affected by acidic and basic solutions are called indicators.

Many indicators do not have a sharp color change as a function of ____________. Most indicators tend to be

__________________ in more acidic solutions.

• Which indicator is best to show an equivalence point pH of 4?



Neutralization Reactions

The reaction of an acid and a base is called a neutralization reaction. Acid + base � salt + water A salt is an

___________________ compound.

• Predict the products of and balance the following neutralization reactions. (Remember to check the

oxidation numbers of the ions in the salt produced.)

a) HNO3 + KOH �

The salt is composed of the ________________ ion and the _________________ ion.

b) HCl + Mg(OH)2 �

c) H2SO4 + NaOH �

Titration

The known reactant molarity is used to find the unknown _________________________ of the other solution.

Solutions of known molarity that are used in this fashion are called _________________________ solutions. In a

titration, the molarity of one of the reactants, acid or base, is known, but the other is unknown.

• A 15.0-mL sample of a solution of H2SO4 with an unknown molarity is titrated with 32.4 mL of 0.145M

NaOH to the bromothymol blue endpoint. Based upon this titration, what is the molarity of the sulfuric acid

solution?

First find the number of moles of the solution for which you know the molarity and volume. Next, use the mole-

mole ratio to determine the moles of the unknown. Finally, determine the molarity of the unknown solution.


100

• If it takes 45 mL of a 1.0 M NaOH solution to neutralize 57 mL of HCl, what is the concentration of the

HCl ?

• If it takes 67.0 mL of 0.500 M H2SO4 to neutralize 15.0 mL of Al(OH)3 what was the concentration of the

Al(OH)3 ?

• How many milliliters of 0.275 M HCl will be needed to neutralize 25.0 mL of 0.154 M NaOH?

Titration Curves

A plot of ___________ versus volume of acid (or base) added is called a titration curve.

Strong Base-Strong Acid Titration Curve

Consider adding a strong base (e.g. NaOH) to a solution of a strong acid (e.g. HCl). Before any base is added, the

pH is given by the strong _________________ solution. Therefore, pH ____ 7. When base is added, before the

equivalence point, the pH is given by the amount of strong acid in _________________________. Therefore, pH <

7. At ________________________________ point, the amount of base added is stoichiometrically equivalent to

the amount of acid originally present. Therefore, pH =_________. To detect the equivalent point, we use an

indicator that changes ____________________somewhere near 7.00. Past the equivalence point all acid has been

consumed. Thus one need only account for excess __________________. Therefore, pH ______ 7.

Homework / Practice

1. Name the following compounds as acids.

a. H2SO4

b. H2SO3

c. H2S

d. HClO4

2. Write formulas for the following acids.

a. nitrous acid

b. hydrobromic acid

c. phosphoric acid

3. Use an activity series to identify two metals that will not generate hydrogen gas when treated with an acid.

4. Write the equation that represents the following reaction: the ionization of HClO3 in water

5. Explain how strong acid solutions conduct an electric current.

6. Write balanced molecular equations for the reactions of acids and bases.

a. aluminum metal with dilute nitric acid

b. calcium hydroxide solution with acetic acid


101

7. CaCO3 (s) + HCl (aq) � CaCl2 (aq) + H2O (l) + CO2 (g)

a. Balance the above equation.

b. How many liters of CO2 form at STP if 5.0 g of calcium carbonate are treated with excess

hydrochloric acid?

8. Consider the following reaction: NH4+ (aq) + CO3

-2 (aq) ↔ NH3 (aq) + HCO3-1 (aq)

a. What reactant serves as the base?

b. What reactant serves as the acid?

9. Given the following reaction: HCO3-1 (aq) + OH-1 (aq) ↔ CO3

-2 (aq) + H2O (l)

a. What reactant serves as the base?

b. What reactant serves as the acid?

10. Write the formula for the salt formed in each of the following neutralization reactions.

a. hydrobromic acid combined with barium hydroxide

b. lithium hydroxide combined with sulfuric acid

11. H2SO4 (aq) + NaOH (aq) � Na2SO4 (aq) + H2O (l)

a. Balance the above neutralization equation

b. In order to completely consume all reactants, what should be the mole ratio of acid to base?

12. Consider the reaction represented by the following incomplete equation:

Ba(OH)2 (aq) + H2SO4 (aq) �

a. Predict the products of this reaction, and write the balanced equation.

b. Use the solubility rules to determine the solubility of the salt produced in the reaction.

c. If 0.030 mol of Ba(OH)2 is consumed, how many grams of water are produced?

13. Perform the following calculations.

a. If the hydronium concentration is 1 x 10-6

M for a solution, calculate the hydroxide concentration.

b. If the hydroxide concentration is 1 x 10-12

M for a solution, calculate the hydronium concentration.

c. If the pOH = 4.00 for a solution, calculate the pH. Is the solution acidic or basic?

d. If the hydronium concentration is 1.00 x 10-3

M, calculate the pOH.

e. If the pOH = 5.0 for a solution, calculate the hydroxide concentration.

f. If the pH = 12.0 for a solution, calculate the hydronium concentration.


102

g. If the pH = 3.00 for a solution, calculate the hydroxide concentration.

h. If the hydronium concentration = 1.0 x 10-8

M for a solution, calculate the hydroxide

concentration.

14. Compare and Contrast Arrhenius acids/bases and Brønsted-Lowry acids/bases.

15. Label the acid (A), base (B) in each of the following reactions.

a. H2SO4 + NH3 → HSO4 + NH4

b. HC2H3O2 + H2O → H3O+ + C2H3O2

c. NaHCO3 + HCl → NaCl + H2CO3

16. Find [OH] for 1.0 × 10

-12 M HClO4.

17. What is the pH of 1.0 × 10-4

M HCl?

18. What is the pH of 1.5 × 10-3

M NaOH?

19. A solution of HNO3 has a pH of 4.0. What is the molarity of HNO3?

20. What is the molarity of KOH in a solution that has a pH of 10.0?

Acids and Bases Practice Test

Directions: Give the names or formulas for the following acids, bases, and salts:

1. KOH________________________

2. HNO3 _______________________

3. Sulfuric Acid__________________

4. Magnesium hydroxide ___________

Directions:.

5. In complete sentences, define an acid according to the Arrhenius theory.

Directions: Label (according to Bronsted-Lowry) the Bronsted-Lowry acid, Bronsted-Lowry base, conjugate acid,

and conjugate base in each of the equations below:

6. H₂O + HC2H3O2 ⇆ H₃O⁺ + C2H3O2⁻

7. CN-- + H₃O⁺ ⇆ H₂O + HCN

Directions Identify the following as an acid or a base, strong or weak.

a.Acid or base b.Strong or weak

8. 2 M KOH __________ _____________

9. 7 M H₂SO₄ __________ _____________

10. 0.12 M H2S _________ _____________

Directions: Complete and balance the following neutralization reactions.

11. NaOH + HCl → _______________ + __________________


103

12. H₂SO₄ + KOH → _____________ + __________________

13. Determine the pOH for a solution of HNO3 that has a concentration of 0.01 M.

14. Determine the pH for a solution of CuOH that has an [OH-] of 0.000001 M.

Directions: Complete the following chart.

[H₃₃₃₃O⁺⁺⁺⁺]or [H+] pH [OH⁻⁻⁻⁻] pOH Acidic or basic

15. 1 x 10⁻⁻⁻⁻6

16.

2

17.

4

18. According to the Arrhenius theory, a base yields

a. H+ as the only positive ion in an aqueous solution

b. OH+ as the only positive ion in an aqueous solution

c. OH- as the only negative ion in an aqueous solution

d. H- as the only negative ion in an aqueous solution

19. In the reaction H2SO4(aq) --> 2H+

(aq) + SO4-2

(aq) H2SO4 is a(n)

a. Arrhenius acid

b. Arrhenius base

c. Salt

20. Arrhenius acids yield

a. OH- as the only negative ion in an aqueous solution

b. H- as the only negative ion in an aqueous solution

c. H3O+ as the only positive ion in an aqueous solution

d. OH+ as the only positive ion in an aqueous solution

21. A substance that conducts an electrical current when dissolved in water is called

a. an acid

b. an electrolyte

c. an ionic compound

d. a nonelectrolyte

22. Which of the following can conduct an electric current?

a. Mg(OH)2 (s)

b. H2O (s)

c. NaOH (aq)

d. NH4Cl (s)

23. In the process of neutralization a salt and a base react to yield water and an acid.

a. True b. False

24. A student dissolved NaCl (s) in water, and tested with a battery, wire, and a light blub to see if it conducted

an electric current. The solution conducted an electric current. This is because NaCl (s) is

a. a salt and an electrolyte

b. a salt and a nonelectrolyte

c. a Arrhenius acid and an electrolyte

d. a Arrhenius base and an electrolyte

25. In the process of neutralization, an Arrhenius acid and an Arrhenius base react to form which of the

following?

a. Water only

b. Salt and Carbon dioxide

c. Water and Carbon dioxide

d. Water and Salt

26. What type of chemical reaction is neutralization?

a. single replacement

b. double replacement

c. synthesis

d. decomposition


104

27. A titration reaction is a complete neutralization reaction where the moles of H+ equal the moles of OH

-.

a. True b. False

28. Which of the following reactants will represent a neutralization?

a. BaCl2 + CaSO4

b. HCl + F

c. Ca(OH)2 + H2SO4

d. NaCl + H2O

29. Titration is a process in which

a. which a volume of solution of unknown concentration is used to determine the concentration of

another solution.

b. which a volume of solution of known concentration is used to determine the volume of another

solution.

c. which a volume of solution of known concentration is used to determine the concentration of

another solution.

d. which a volume of solution of known concentration is used to determine the curve of another

solution.

30. What is the molarity of HCl (aq) if 25 mL of 8.0 M NaOH (aq) neutralizes exactly 20.0 mL of HCl (aq)?

a. 5M

b. 10M

c. 15M

d. 20M

31. At the end point of titration, what is the relationship between moles of H+ and OH

-?

a. the moles of H+ are greater than OH

-

b. the moles of OH- are greater than H

+

c. the moles of H+ are equal to moles of OH

-

d. there is no relationship between moles of H+ and OH

-

32. What is the molarity of NaOH if 5 milliliters of 4M HCl (aq) neutralizes exactly 10 mL of NaOH (aq)?

a. .5M

b. 1M

c. 1.5M

d. 2M

33. If 10.milliliters of a 0.40 M HBr solution is required to neutralize exactly 0.2 M of NaOH, what is the

volume of the base?

a. 10ml

b. 20ml

c. 30ml

d. 40ml

34. The molarity of an acid can be calculated if a base of known concentration (standard base) is added, drop

by drop, to a specific volume of the acid until the indicator changes color.

a. True b. False

35. One acid-base theory states that an acid is an H+

a. Acceptor

b. Eliminator

c. Dissolver

d. Donor

36. According to the Bronsted-Lowry acid-base theory, a base is a substance that can

a. donate an electron

b. accept a proton

c. donate a proton

d. accept a electron

37. The acidity or alkalinity of a solution can be measured by its pH value.

a. True b. False

38. In the following reaction, NH3 + HCl --> NH4+ + Cl

- NH3 acts as a(n)

a. base in the reverse reaction.

b. acid in the forward reaction.

c. base in the forward reaction.

d. acid in the reverse reaction.


105

39. Which of these 1 M solutions will have the highest pH?

a. H3PO4

b. HCl

c. NaCl

d. NaOH

40. Which pH indicates an acidic solution?

a. 1 b. 7 c. 9 d. 12

41. Which of these pH numbers indicates the lowest level of acidity?

a. 1 b. 3 c. 8 d. 12

42. Which formula represents a salt?

a. KOH

b. KCl

c. CH3OH

d. CH3COOH

43. Which substance can be classified as an Arrhenius acid?

a. HCl

b. NaCl

c. LiOH

d. KOH

44. An acidic solution could have a pH of

a. 7 b. 10 c. 3 d. 14

45. What is the pH of a 0.00001 molar HCl solution?

a. 1 b. 9 c. 5 d. 4

46. What is the pH of a solution with a hydronium ion concentration of 0.01 moles per liter?

a. 1 b. 2 c. 10 d. 14

47. Given the equation: H+ + OH

- ↔ H2O Which type of reaction does the equation represent?

a. esterification

b. decomposition

c. hydrolysis

d. neutralization

48. As the hydrogen ion concentration of an aqueous solution increases, the hydroxide ion concentration of this

solution will

a. decrease

b. increase

c. remain the same

49. A student wishes to prepare approximately 100 milliliters of an aqueous solution of 6 M HCl using

12 M HCl. Which procedure is correct?

a. adding 50 mL of 12 M HCl to 50 mL of water while stirring the mixture steadily.

b. adding 50 mL of 12 M HCl to 50 mL of water and then stirring the mixture steadily.

c. adding 50 mL of water to 50 mL of 12 M HCl while stirring the mixture steadily.

d. adding 50 mL of water to 50 mL of 12 M HCl and then stirring the mixture steadily.

50. What is the pH of an acetic acid solution if the [H3O+] = 1 x 10

-4 mol/L?

a. 1

b. 2

c. 3

d. 4

e. 5

Unit 12- Kinetics and Thermochemistry

106

REACTION KINETICS

Energy Diagrams

Reactants always start a reaction so they are on the _________________ side of the diagram. Products are on the

right. The exothermic reaction gives off ___________________ because the products are at a lower energy level

than the reactants. In an exothermic graph, the reactants have _____________________ energy than the products.

The change in energy is a _________________________ value.

The endothermic reaction absorbs heat because the products are at a _________________________ energy level

than the reactants. In an endothermic graph, the products have _____________________ energy than the reactants.

The change in energy is a _______________________ value.

Scientists have observed that the energy released in the formation of a compound from its elements is always

identical to the energy required to ______________________ that compound into its elements.

_____________________________ energy is the minimum amount of energy that reacting particles must have to

form the activated complex. The activated complex is a short-lived, _________________ arrangement of atoms that

may break apart and re-form the reactants or may form products. To calculate the activation energy,

______________________ the energy of the reactants from the energy at the top of the peak. The enthalpy or heat

of reaction (∆H) is the amount of ___________________ released or absorbed in the reaction. To determine ∆H,

take the energy of the products and _______________________ the energy of the reactants.

• The heat content of the reactants of the forward reaction is about ________ kilojoules.

• The heat content of the products of the forward reaction is

about ________ kilojoules.

• The heat content of the activated complex of the forward

reaction is about _______ kilojoules.

• The activation energy of the forward reaction is about

_______ kilojoules.

• The heat of reaction (∆H) of the forward reaction is about

_______ kilojoules.

• The forward reaction is (endothermic or exothermic).


107

The activation energy can be lowered by adding a __________________________. The catalyst

_________________ the activation energy by providing an alternate pathway for the reaction to occur.

Expressing Reaction Rates

We generally define the average _____________________ of an action or process to be the change in a given

quantity during a specific period of time. Reaction rates cannot be calculated from balanced equations as

stoichiometric amounts can. Reaction rates are determined experimentally by measuring the

_____________________________ of reactants and/or products in an actual chemical reaction.

Collision Theory

According to the collision theory, atoms, ions, and molecules must collide with each other in order to react. The

following three statements summarize the collision theory.

1. Particles must ________________________ in order to react.

2. The particles must collide with the correct _______________________________.

3. The particles must collide with enough ________________________ to form an unstable activated

complex, also called a _____________________________ state, which is an intermediate particle made

up of the joined reactants.

The _____________________________ amount of energy that colliding particles must have in order to form an

activated complex is called the activation energy of the reaction. Particles that collide with energy less than the

activation energy ____________________________ form an activated complex. In an exothermic reaction,

molecules collide with enough energy to overcome the activation energy barrier, form an activated complex, then

__________________________ energy and form products at a lower energy level. In the reverse endothermic

reaction, the reactant molecules lying at a _______________ energy level must absorb energy to overcome the

activation energy barrier and form high-energy products.

Factors Affecting Reaction Rates

The reaction rate for almost any chemical reaction can be modified by varying the conditions of the reaction.

1) ___________________________________. As you know, some substances react more readily than others.

The more reactive a substance is, the _____________________________ the reaction rate.

2) ____________________________________. The rate of a reaction is _________________ when the

concentrations of reacting particles are increased. Increasing the number of reactant particles increases

probability of collisions. The rate of gaseous reactions can be ___________ by pumping more gas into the

reaction container.

3) ______________________ the surface area of reactants provides more opportunity for collisions with other

reactants, thereby ______________________ the reaction rate.


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4) Generally, increasing the ____________________________ at which a reaction occurs

_____________________ the reaction rate. Raising the temperature ___________________ both the

collision frequency and the collision energy.

5) Adding a ___________________ affects the rate of a chemical reaction. A catalyst is a substance that

increases the rate of a chemical reaction without itself being _____________________in the reaction.

6) ________________________ gases increases the rate of a chemical reaction.

REACTION ENERGY

Energy is the ability to do___________________ or produce heat. It exists in two basic forms, potential energy and

______________________ energy. Potential energy is energy due to the ______________________ or position of

an object. Kinetic energy is energy of ________________. The potential energy of the dammed water is converted

to kinetic energy as the dam gates are opened and the water flows out. Chemical systems contain

_________________ kinetic energy and potential energy. As temperature increases, the motion of submicroscopic

particles ______________________, so its kinetic energy __________________________. The potential energy of

a substance depends upon its composition: the type of atoms in the substance, the number and type of chemical

bonds joining the atoms, and the particular way the atoms are arranged.

Law of Conservation of Energy and Heat

The law of conservation of energy states that in any chemical reaction or physical process, energy can be converted

from one form to another, but it is neither created nor ________________________.

Heat, which is represented by the symbol ____, is energy that is in the process of flowing from a

_____________________ object to a cooler object. The SI unit of heat and energy is the joule (J). Heat involves a

transfer of energy between 2 objects due to a ________________________ difference. When the warmer object

loses heat, its temperature decreases and q is _________________________. When the cooler object absorbs heat,

its temperature ________________ and q is positive.

The specific heat of any substance is the amount of heat required to raise the temperature of ____ gram of that

substance by one degree Celsius. Because different substances have different compositions, each substance has its

own specific heat.

q = m Cp ∆T

q = heat (J); m = mass (g); Cp = specific heat (J/(g.°C); ∆T = change in temperature = Tf – Ti (°C)

Exothermic: Heat flows _________ of the system (to the surroundings). The value of ‘q’ is negative. Endothermic:

Heat flows _________ the system (from the surroundings). The value of ‘q’ is positive.

• The temperature of a sample of iron with a mass of 10.0 g changed from 50.4°C to 25.0°C with the release

of 114 J heat. What is the specific heat of iron?


109

• A piece of metal absorbs 256 J of heat when its temperature increases by 182°C. If the specific heat of the

metal is 0.301 J/g.°C, determine the mass of the metal.

• If 335 g water at 65.5°C loses 9750 J of heat, what is the final temperature of the water? The specific heat

of water is 4.18 J/g.°C.

• As 335 g of aluminum at 65.5°C gains heat, its final temperature is 300.°C. The specific heat of aluminum

is 0.897 J/g.°C. Determine the energy gained by the aluminum.

Heat changes that occur during chemical and physical processes can be measured accurately and precisely using a

___________________________. A calorimeter is an insulated device used for measuring the amount of heat

absorbed or released during a chemical or physical process. A coffee-cup calorimeter made of ________ Styrofoam

cups.

• Suppose you put 125 g of water into a foam-cup calorimeter and find that its initial temperature is 25.6°C.

Then, you heat a 50.0 g sample of the unknown metal to a temperature of 115.0°C and put the metal sample

into the water. Both water and metal have attained a final temperature of 29.3°C. Heat flows from the hot

metal to the cooler water and the temperature of the water rises. The flow of heat stops only when the

temperature of the metal and the water is equal. Assuming no heat is lost to the surroundings, the heat

gained by the water is equal to the heat lost by the metal. Determine the specific heat of the metal.

• You put 352 g of water into a foam-cup calorimeter and find that its initial temperature is 22.0°C. What

mass of 134°C lead, Clead = 0.129 J/g°C, can be placed in the water so that the equilibrium temperature is

26.5°C?

• You put water into a foam-cup calorimeter and find that its initial temperature is 25.0°C. What is the mass

of the water if 14.0 grams of 125°C nickel, CNi = 0.444 J/g°C, can be placed in the water so that the

equilibrium temperature is 27.5°C?

Phase Changes Review

Solid → liquid ________________________ Liquid → solid ___________________

Liquid → gas ________________________ Gas → liquid _____________________

Solid → gas ________________________ Gas → solid ______________________

Energy and Phase Changes

q = m Hf q = m Hv

Hf = latent heat of fusion (J/g) ; Hv = latent heat of vaporization (J/g)

Heat of vaporization (Hv) is the energy required to change one gram of a substance from ________________ to gas.

Heat of fusion (Hf) is the energy required to change one gram of a substance from __________________ to liquid.


110

� How much heat does it take to melt 12.0 g of ice at 0 °C? Hf for water is 334 J/g.

� How much heat must be removed to condense 5.00 g of steam at 100 °C? Hv = 2260 J/g.

Three equations can be used in calculating energy.

q = m Cp ∆T q = m Hf q = m Hv

Solving Problems

The total heat equals the sum of all the heats you have to use.

Go in the following order when energy is being added to the system.

1) Heat ice q = m Cice ∆T 2) Melt ice q = m Hf

3) Heat water q = m Cwater ∆T 4) Boil water q = m Hv

5) Heat steam q = m Csteam ∆T

Numbers Needed For Energy Problems Involving Water (look up in reference tables)

For ice, specific heat = __________ For water, specific heat = _________

For steam, specific heat = _____________ Heat of vaporization = ___________

Heat of fusion = _____________

� How much heat does it take to heat 12 g of ice at – 6 °C to 25 °C water? Round to a whole number.

� How much heat does it take to heat 35 g of ice at 0 °C to steam at 150 °C? Round to a whole number.

� How much heat does it take to convert 16.0 g of ice to water at 0 °C?

� How much heat does it take to heat 21.0 g of water at 12.0°C to water at 75.0°C?

� How much heat does it take to heat 14.0 g of water at 12.0°C to steam at 122.0°C?

Entropy

Entropy (S) is a measure of the ______________________ or randomness of the particles that make up a system.

Spontaneous processes always result in a(n) _____________________ in the entropy of the universe. Entropy of a

solid < Entropy of a liquid << Entropy of a gas A solid has an orderly arrangement. A liquid has the molecules

next to each other. A gas has molecules moving all over the place.

Several factors affect the change in entropy of a system.

1. Changes of state. Entropy ___________________ when a solid changes to a liquid and when a liquid

changes to a gas because these changes of state result in freer movement of the particles.

2. Dissolving of a gas in a solvent. When a gas is dissolved in a liquid or solid solvent, the motion and

randomness of the particles are limited and the entropy of the gas _____________.

3. Change in the number of gaseous particles. When the number of gaseous particles increases, the entropy

of the system usually ___________________ because more random arrangements are possible.


111

4. Dissolving of a solid or liquid to form a solution. When solute particles become dispersed in a solvent, the

disorder of the particles and the entropy of the system usually ________________.

5. Change in temperature. A temperature increase results in increased disorder of the particles and a(n)

____________________ in entropy.

� Predict the sign of ∆Ssystem for: O2 (g) � O2 (aq)

� Predict the sign of ∆Ssystem for: C6H6 (s) � C6H6 (l)

� Predict the sign of ∆Ssystem for: C (s) + CO2 (g) � 2 CO (g)

Homework / Practice

1. If 200. g of water at 20.0 °C absorbs 41840 J of heat, what will its final temperature be?

2. Aluminum has a specific heat of 0.900 J/g.°C. How much energy is needed to raise the temperature of a

625 g block of aluminum from 30.7 °C to 82.1 °C?

3. If a reaction is exothermic, are the products or the reactants more stable?

4. List 4 factors that can speed up a chemical reaction.

5. For each of the following examples, state whether the change in entropy is positive, negative or remains the

same.

(a) HCl (l) → HCl (g)

(b) C6H12O6 (aq) → C6H12O6 (s)

(c) 2 NH3 (g) → N2 (g) + 3 H2 (g)

(d) 3 C2H4 (g) → C6H12 (l)

6. Consider the following equilibrium equation: H2O (g) + C (s) → H2 (g) + CO (g) + heat energy. Will

the reaction rate increase or decrease when each of the following occurs?

(a) a catalyst is introduced

(b) the temperature of the system is lowered

7. Convert from one unit to the other:

(a) 1.69 Joules to calories

(b) 820.1 J to kilocalories

(c) 423 calories to kilocalories

(d) 20.0 calories to Joules

(e) 252 cal to J

(f) 2.45 kilocalories to calories

8. When 15.0 g of steam drops in temperature from 275.0 °C to 250.0 °C, how much heat energy is released?

9. How much heat (in kJ) is given out when 85.0 g of lead cools from 200.0 °C to 10.0 °C?

10. A certain mass of water was heated with 41,840 Joules, raising its temperature from 22.0 °C to 28.5 °C.

Find the mass of water.


112

11. Determine the energy required (in kilojoules) when cooling 456.2 grams of water at 89.2 °C to a final

temperature of 5.9 °C.

12. Determine the specific heat of a 150.0 gram object that requires 62.0 cal of energy to raise its temperature

12.0 °C.

13. Determine the energy required to raise the temperature of 46.2 grams of aluminum from 35.8 °C to 78.1 °C.

14. Determine the energy required to:

(a) melt 5.62 moles of ice at 0 °C.

(b) boil 43.89 grams of water at 100.0 °C.

(c) Convert 16.2 grams of ice to liquid water.

(d) Convert 5.8 grams of water to steam

(e) Convert 98.2 grams of water to ice.

(f) Convert 52.6 grams of steam to water

(g) Convert 125.0 grams of ice at 0.0 °C to steam at 100.0 °C.

(h) Convert 25.9 grams of steam at 100.0 °C.to ice at 0.0 °C.

15. How many degrees of temperature rise will occur when a 25.0-g block of aluminum absorbs 10.0 kJ of

heat?

Kinetics and Thermochemistry Practice Test

Directions: Match the terms below with their correct definitions. (1-4)

a. calorimeter

b. thermochemistry

c. enthalpy

d. system

1. The study of heat changes that accompany chemical reactions and phase changes

2. The specific part of the universe that contains the reaction or process you wish to study

3. The heat content of a system at constant pressure

4. An insulated device used to measure the amount of heat absorbed or released during a chemical or physical

process

5. Predict the change in entropy (∆S) for the following reaction (will it increase, decrease, or stay the

same).CH4(g) + 2O2(g) � 2H2O(l) + CO2(g)

6. Which of the following species has the highest entropy at 25°C? Explain your answer.

a) CH3OH(l)

b) CO(g)

c) MgCO3(s)

d) H2O(l)

e) Ni(s)

7. If the temperature of a 25-g sample of liquid water is raised 40°C, how much heat is absorbed by the

water?

8. Copper metal has a specific heat of 0.385 J/g·°C , calculate the amount of heat required to raise the

temperature of 38 g of copper from 20.0°C to 575°C.


113

9. When a 50.0-g nugget of pure gold is heated from 35.0°C to 50.0°C, it absorbed 5200.0 J of energy.

Find the specific heat of gold.

10. When 80.0 grams of a certain metal at 90.0 °C was mixed with 100.0 grams of water at 30.0 °C, the

final equilibrium temperature of the mixture was 36.0 °C. What is the specific heat of the metal?

Directions: Use the energy diagram for the rearrangement

reaction of methyl isonitrile to acetonitrile to answer the

following questions. (11-13)

11. What kind of reaction is represented by this diagram,

endothermic or exothermic?

12. What does the symbol E represent?

13. How does a catalyst speed a reaction?

14. As ice cools from 273 K to 263 K, the average kinetic energy of its molecules will

a) decrease

b) increase

c) remain the same

15. The heat of fusion is defined as the energy required at constant temperature to change 1 unit mass of a

a) gas to a liquid

b) gas to a solid

c) solid to a gas

d) solid to a liquid

16. What is the total number of joules of heat energy absorbed by 15 grams of water when it is heated from

30°C to 40°C?

a) 10

b) 63

c) 150

d) 630

17. How many joules of heat are absorbed when 70.0 grams of water is completely vaporized at its boiling

point?

a) 23, 352

b) 7, 000

c) 15, 813

d) 158, 130

18. What occurs as potassium nitrate is dissolved in a beaker of water, indicating that the process is

endothermic?

a) The temperature of the solution decreases.

b) The temperature of the solution increases.

c) The solution changes color.

d) The solution gives off a gas.

19. Given the change of phase: CO2(g) changes to CO2(s), the entropy of the system

a) decreases

b) increases

c) remains the same

20. A solid is dissolved in a beaker of water. Which observation suggests that the process is endothermic?

a) The solution gives off a gas.

b) The solution changes color.

c) The temperature of the solution decreases.

d) The temperature of the solution increases.


114

21. When a catalyst is added to a system at equilibrium, a decrease occurs in the

a) activation energy

b) potential energy of the

reactants

c) heat of reaction

d) potential energy of the

products

22. Which statement explains why the speed of some chemical reactions is increased when the surface area

of the reactant is increased?

a) This change increases the density of the reactant particles.

b) This change increases the concentration of the reactant.

c) This change exposes more reactant particles to a possible collision.

d) This change alters the electrical conductivity of the reactant particles.

23. Which conditions will increase the rate of chemical reaction?

a) decreased temperature and decreased concentration of reactants?

b) decreased temperature and increased concentration of reactants?

c) increased temperature and decreased concentration of reactants?

d) increased temperature and increased concentration of reactants?

24. In a chemical reaction, a catalyst changes the

a) potential energy of the

products

b) heat of reaction

c) potential energy of the

reactants

d) activation energy

25. Catalysts increase the rate of reaction by being consumed.

a) True b) False

26. The following reaction coordinate diagram represents...

a) an endothermic reaction

b) an exothermic reaction

c) a reaction that is neither

endothermic nor exothermic

d) a reaction in which a

catalyst is used

27. If 1.45 J of heat are added to a 2.00 g sample of aluminum metal and the temperature of the metal

increases by 0.798 oC, what is the specific heat of aluminum?

a) 0.579 J/g deg

b) 0.909 J/g deg

c) 1.68 J/g deg

d) 3.63 J/g deg

28. Water has a specific heat of 4.184 J/g deg while glass (Pyrex) has a specific heat of 0.780 J/g deg. If

10.0 J of heat is added to 1.00 g of each of these, which will experience the larger increase of

temperature?

a) glass b) water

c) They both will experience the same change in temperature since only the amount of a

substance relates to the increase in temperature.

29. The collision theory states that a reaction is most likely to occur if reactant particles collide with the

proper

a) energy and concentration

b) energy and orientation

c) concentration and orientation

d) pressure and orientation


115

30. What will happen to the rate of reaction when temperature increases?

a) Increase

b) Decrease

c) remains the same

d) increase then decrease

31. Given the following reaction H+

(aq) + OH-(aq) --> H2O(l) if the concentration of the reactants is increased,

the rate of reaction will

a) Increase

b) Decrease

c) remains the same

d) increase then decrease

32. Entropy is a measure of the randomness or disorder of a system.

a) True b) False

33. What does a catalyst decrease when introduced to a reaction?

a) the rate of reaction

b) the energy released

during the reaction

c) the activation energy

d) the kinetic energy

34. Which process is accompanied by a increase in entropy?

a) melting of ice

b) freezing of water

c) condensing of water

vapor

35. Systems in nature tend to undergo changes toward what kind of energy and entropy?

a) lower energy and lower entropy

b) lower energy and higher entropy

c) higher energy and higher entropy

d) higher energy and lower entropy

36. What phase change represents a decrease in entropy?

a) solid to liquid

b) liquid to gas

c) gas to liquid

d) solid to gas

37. H2O(g) → H2O(l) The entropy in this equation

a) Increases

b) Decreases

c) remains the same

38. Which sample has the lowest entropy?

a) 1 mole of KNO3(l)

b) 1 mole of KNO3(g)

c) 1 mole of H2O(g)

d) 1 mole of KNO3(s)

39. In a chemical reaction, if the products have more entropy than the reactants, the change in entropy is

negative.

a) True b) False

40. What is the change in entropy in the following reaction C + O2 -> CO2

a) Increases

b) Decreases

c) remains the same

d) not enough information

41. What is the change in entropy in the following reaction 4Al(s) + 3O2(g) -> 2Al2O3(s)

a) Increase

b) Decrease

c) remains the same


42. What is the change in entropy in the following reaction N2(g) + O2(g) -> 2NO(g)

a) Increase

b) Decrease

c) remains the same



116

43. A 47.5 gram sample of a metal at a temperature of 425°C is placed in 1.00 liters (1000 g) of water

which had an initial temperature of 18°C. What is the specific heat capacity of the metal if the final

temperature of the metal and water at equilibrium is 21°C? (The specific heat capacity of water is

4.18 J/°C·g.

a) 0.03 J/°C·g

b) 1.47 J/°C·g

c) -0.75 J/°C·g

d) 0.65 J/°C·g

e) 12.54 J/°C·g

44. 20.0 mL of pure water at 285 K is mixed with 48 mL of water at 315 K. What is the final temperature

of the mixture in kelvins?

a) 306 K

b) 290 K

c) 318 K

d) None of these

e) 275 K

45. As a result of an exothermic reaction,

a) the energy of the system is increased and the energy of the surroundings are decreased.

b) the energy of the system and the energy of the surroundings are decreased.

c) the energy of the system is decreased and the energy of the surroundings are increased.

d) the energy of the system and the energy of the surroundings are increased.

e) None of these are accurate

46. How much energy is needed to convert 50.0 g of ice at -5.00°C to water at 25°C?

a) 22.4 kJ

b) 37.3 kJ

c) 21.9 kJ

d) 18.0 J

e) 175 J

47. The specific heat of iron is 0.450 J/(g·°C). How much heat is required to raise the temperature of a

5.00 gram sample of iron from 22°C to 53°C?

a) -43 J

b) 155 J

c) 344 J

d) 18 J

e) 69.8 J

48. Which of the following is not an endothermic process?

a) combustion

b) melting

c) crystallization

d) vaporization

e) sublimation

east meck chemistry -...

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