final module 2.2007.sep - weber state university … · ... has mass and occupies space. ... of any...

2.1 INORGANIC CHEMISTRY

2 Inorganic Chemistry

Objective 1

Defi ne and give examples of the following terms:

Reading Assignment: Read Text pages 27-28.

Weight: Downward force exerted on an object as a result of gravity.

Weight is a function of the distance of an object from the earth.

Mass: Amount of matter in an object.

Mass is NOT a function of distance from the earth.

It remains constant regardless of location.

Density: Mass per unit volume. Density = Mass/Volume

The amount of matter in a particular volume.

The “compactness" of an object.

Examples of density units: g/dL, g/cm2, mg/L, mg/dL

Matter: Has mass and occupies space.

States of Matter: Solid Liquid GasAbility to completely fi ll a container

none some completely

Proximity of molecules close intermediate distant

Amount of molecular motion low intermediate high

Compressibility minimal minimal to none very compressible

Density very high high low

Element: Building blocks of matter. For example: carbon, calcium, oxygen, iron, etc.

Atom: Smallest component of an element that still has properties of that element.

INORGANIC CHEMISTRY 2.2

Objective 2

Given the atomic symbol of any of the following elements, identify the name of the element.

Reading Assignment: Refer to Table 2.1, 8.

hydrogen H nitrogen N iodine I

carbon C chlorine Cl lithium Li

oxygen O fl uorine F silver Ag

calcium Ca sulfur S phosphorous P

potassium K magnesium Mg lead Pb

sodium Na iron Fe copper Cu

mercury Hg

Objective 3

Identify the major function(s) performed in living organisms by the following elements:

Reading Assignment: Refer to Table 2.1, Text, Page 28.

Calcium (Ca): Structure of bones and teeth, muscle contraction

Phosphorus (P): ATP, protein, DNA

Chlorine (Cl): NaCl (salt), water balance

Sulfur(S): Protein s, disulfi de bridges

Potassium (K): Nerve transmission, muscle contraction

Sodium (Na): NaCl, water balance, membrane depolarization, nerve transmission

Magnesium (Mg): Enzymes

Iodine (I): Thyroid function

Iron (Fe): Blood hemoglobin

Oxygen (O): Element in water molecule (H2O), aerobic cellular respiration

Carbon (C): Structural framework


Objective 4

Defi ne the following terms.

Reading Assignment: Read Text Pages 27-29.

Proton: (p+) Positively-charged particle in the nucleus of an atom, rela-tively heavy. The number of protons defi nes the atom's identity.

Neutron: (no) Neutral or uncharged (i.e., no charge) particle in the nucleus, relatively heavy.

Electron: (e-) Negatively-charged particle outside the nucleus, relatively light.

Neutral Atom:

An atom with no net electrical charge. For a neutral atom, the number of electrons = the number of protons.

Atomic Number:

Equal to the number of protons in an atom's nucleus. It is also equal to the number of (e-) in an UNCHARGED atom.

Atomic Weight:

The mass of an average atom of an element

Mass Number:

The number of protons plus number of neutrons or round off the atomic weight.

Isotopes: Atoms with the same number of p+, diff erent number of no, and diff erent mass number. Isotopes of the same element form a set of nuclides.

Objective 5

Explain why the list of elements is referred to as a "periodic" table/chart. Given a representative section from the Periodic Table of the Elements determine the following for the average atom of any element: a. The atomic number b. The atomic weight and mass number. c. The number of electrons in a neutral atom. d. The number of protons and number of neutrons in the nucleus.

Reading Assignment: Read Text, Pages 27-29. Refer to the periodic table on the next page.Appendix III Page A-39

Columns (IA, IIA, etc.) are called “groups”. Rows are called “periods.”When elements are arranged in order of increasing atomic number, there is a periodic repetition of properties. We will only consider the fi rst 20 elements for detailed analysis.


Arrangement of the Periodic Table (chart) of the Elements:

Metals: Left two-thirds of periodic chart

Non-metals: Right one-third of periodic chart

Metalloids: Elements that form a narrow diagonal band in the periodic table between the metals and non metals; have properties similar to both metals and minerals

1H

1.008

Atomic number = number of protonsAtomic symbolAtomic weight (round off to get the mass number of most common iso-tope)

The number of neutrons = Atomic weight (rounded-off to whole number) minus the atomic number.

* This periodic chart is abridged for teaching purposes.

*


Objective 6

Defi ne electron shell. Given a representative section from the periodic table, predict the number of electrons in the outer shell.

Reading Assignment: Read Text, Pages 29-30.Refer to Figure 2.2 Page 29

A shell is a location and enegry of elections around a nucleusThe electron distribution and confi guration around the nucleus infl uences how an atom will interact to form a molecule(s)—outer shell electrons are very important.For example, the electrons in calcium (Ca) have the following arrangement:

1st shell has 2e-

2nd shell has 8e-

3rd shell has 8e-

4th shell has 2e-

Atoms of elements 3-20 are most stable when they have 8 electrons in their outer shells.Elements 1 & 2 only need 2 in their outer shell.

Group numbers indicate how many electrons are found in the outer shell.

If atoms are not stable, they will "try" to become stable by either giving away or gaining electrons.Elements near the left side of the periodic table generally give away electrons, whereas elements near the right side tend to gain electrons.

Objective 7

Defi ne the following terms associated with gaining or losing electrons:

Reading Assignment: Read Text, Page 30-31.Refer to Figure 2.3 page 31

Ion: Charged atoms or charged particles. Examples: K+, Na+, Mg++, Ca++, Cl-

Cation: Positively-charged ion

Anion: Negatively-charged ion


Objective 8

Given a representative section of the periodic table, determine whether an atom generally loses or gains electrons when undergoing a chemical reaction. Considering the number of electrons donated or accepted during a chemical reaction, predict the charge on the resulting ion(s).

Reading Assignment: Read Text, Page 30-31.

Elements on the left side of the periodic table tend to donate electrons, leaving them with a net positive charge.

Elements on the right side of the periodic table tend to accept electrons, leaving them with a net negative charge.

One electron gained by an atom results in a charge of -1 on that atom, two electrons gained results in a charge of -2, etc.

One electron donated results in a charge of +1, two electrons donated results in a charge of +2, etc.

Objective 9

Describe the characteristic and common elements involved in the forming of ionic, covalent, and hydrogen bonds.

Reading Assignment: Read Text, Pages 30-33; Refer to Figures 2.3, 2.4, 2.5, 2.6.

Ionic: An ionic bond occurs between a metal and a nonmetal. It involves donating and accepting electrons (not sharing).It results when a cation and an anion attract each other.

Covalent: A covalent bond occurs between two nonmetals.It involves sharing of electrons.It results when neutral atoms share electrons.

Polar Covalent: A polar covalent bond exists when the sharing of electrons between two nonmetal atoms is unequal.

Nonpolar Covalent: A nonpolar covalent bond exists when two nonmetal atoms share electrons equally.

Hydrogen: A hydrogen bond usually occurs between a hydrogen atom and oxygen or nitrogen. It involves unequal sharing of electrons, which leaves hydrogen with a partial positive charge and the oxygen or nitrogen with a partial negative charge.It results from a relatively weak attraction between this partially positively-charged H atom and another partially negatively-charged atom.


Objective 10

Defi ne the following terms:

Reading Assignment: Read Text, Pages 27, 30 & 33.

Compound: A substance formed from the chemical combination of 2 or more elements.

Molecule: Smallest component of any compound or of a pure substance.

Molecular weight: 1. Mass of a molecule.2. Round atomic weights of each atom to the nearest whole number. 3. Add them together.

Objective 11

Given the chemical formula of a compound, identify the constituent atoms and the number of each in one molecule of the compound.

Reading Assignment: Refer to Figure 2.4 Page 31; Refer to Table 2.2 page 34.

Examples:

H2O

H 1.008 1CO

2

C 12.011 12H 1.008 1 O 15.99 16O 15.99 16 O 15.99 16

Molecular weight of H2O: 18 Molecular weight of CO

2: 44

1. Sulfuric acid: H

2SO

4 = 2 atoms hydrogen, 1 atom sulfur, 4 atoms oxygen

MW = (2 x 1) + (1 x 32) + (4 x 16) = 98 2. Calcium phosphate = Ca

3(PO

4)

2 A subscript following a parenthesis means that every

element in the parenthesis is multiplied by that number. Therefore, in a molecule of calcium phosphate, there are:3 calcium atoms; (2 x 1) or 2 phosphorus atoms; (2 x 4) or 8 oxygen atoms

MW = (3 x 40) + (2 x 31) + (8 x 16)

120 + 62 + 128 = 310


Objective 12

Defi ne the term polyatomic ion. Given the formula for a polyatomic ion, be able to provide a name for the ion.

Reading Assignment: No additional sources.

Polyatomic IonsAs we learned earlier, if an atom becomes either negatively-charged (by accepting one or more extra electrons), or positively-charged (by donating one or more electrons) it is called an ion.

Sometimes we observe a group of two or more diff erent kinds of atoms whose atomic structures favor one another so much that they typically appear together. These groups are called polyatomic molecules (“many atoms”). If there is either an excess or defi cit in the total electrons (i.e., a positive or negative net charge) in the group, then it is called a polyatomic ion. For example, in sodium hydroxide (NaOH) we fi nd a common polyatomic ion ( the hydroxide ion--OH-) forming an ionic bond with a sodium cation (Na+). The hydroxide ion is composed of two diff erent kinds of atoms: oxygen and hydrogen. The hydroxide ion contains a total of 10 electrons: 8 electrons (from oxygen) plus 1 electron (from hydrogen) plus 1 extra from another atom. This is a polyatomic ion, because it has one extra electron and has a net charge of -1. The name, formula, and charge of some common polyatomic ions are given in the table below. You should know the ion by its symbol (e.g., HCO

3-1 is

bicarbonate).

Common Polyatomic IonsName of Ion Formula Name of Ion Formula

Bicarbonate HCO3

- Sulfate SO4

- 2

Carbonate CO3-

-2 Sulfi te SO3

- 2

Phosphate PO4

- 3 Hydroxide OH-

Nitrate NO3

- Ammonium NH4

+

Nitrite NO2

-

Objective 13

Defi ne the terms solution, colloid and suspension. Distinguish the relative size of the particles present in solutions, colloids, and suspensions.

Reading Assignment: Read Text, Pages 37-40.


Particle sizes: Solutions: Very small (crystalloids).

Colloids: Medium”size (large molecules).

Suspensions: Large (sand).

Defi nitions of Terms:

Solution: A homogenous mixture of two or more substances.

Solute: The substance that is put into the liquid. Technically, the substance(s) present in a solution in the smaller amount.

Solvent: Generally speaking, it is the liquid fraction of a mixture. However, it can also be technically defi ned as the material present in a solution in the largest amount.

SOLUTIONS, COLLOIDS AND SUSPENSIONS

Protoplasm, or the living matter in a cell, is a solution. What is a solution? Solutions are homogeneous (uniform) mixtures of two or more substances. The substance which goes into solution is called a solute, while the substance which dissolved the solute is called the solvent. For example, if one places a sugar cube in a glass of water, the sugar soon disap-pears. The sugar has dissolved in the water. In this case, sugar is the solute and water is the solvent. In the solution, there are still molecules of sugar and molecules of water, but they have evenly dispersed themselves, due to the fact that molecules are in constant motion.

In protoplasm, the solvent is water. There are thousands of solutes dissolved in this water to make the solution called protoplasm.

A solution is said to be dilute if there is only a small amount of solute per unit volume of solution. Concentrated solutions have much more solute per unit volume of solution. Saturated solutions have as much solute present as the solvent is capable of dissolving. There are several ways to express concentrations in a more precise manner.

Percent solutions express concentration as the part of solute per 100 parts of solution. For example, a 5% solution of glucose has fi ve grams of glucose dissolved in enough water to make 100 ml. of glucose solution (5 g/100ml or 5 g/dl). It is important to recognize that one ml of distilled water weighs one gram for conversions of volume to weight.

Besides having some understanding of the concentrations of solutions, one needs to under-stand something about the size of the particles in solutions versus the size of particles in other mixtures.

SOLUTIONS: Solutions contain very small particles such as ions and small molecules. These particles (those that are small enough to form solutions) are also referred to as


crystalloids. Crystalloids remain evenly dispersed in the solution---they do not “settle out”. They are small enough to pass through both the relatively large pores of fi lter paper and through the smaller pores of a semipermeable membrane (semipermeable membranes are used for dialysis in artifi cial kidneys).

COLLOIDS: Colloids consist of larger molecules or groups of molecules such as proteins. The particles do not settle out, largely because they are similarly-charged and thus repel each other. Colloids will pass through fi lter paper, but are too large to pass through a semipermeable membrane.

SUSPENSIONS: Suspensions include large particles such as sand and clay. Muddy water is an example of a suspension. If you allow it to stand, the particles of mud settle to the bottom. If you have seen a bottle or bag of blood in a blood bank refrigerator, you will notice that the red blood cells have settled to the bottom of the container. Cells are in suspension in the blood.

Our blood, therefore, is a suspension of cells. It is also a crystalloid solution (dissolved salts, etc.) as well as a colloid, because proteins are “dissolved” in it. Protoplasm also has both crystalloids and colloids in it.

The cytoplasm, which is the cellular protoplasm bounded by the cellular and nuclear mem-brane, is a water solution of colloids and crystalloids. Suspended in this solution are the cellular organelles such as mitochondria, lysosomes, and free ribosomes.

Objective 14

Distinguish the properties of solutions and suspensions..

Solutions: Suspensions:

Clear Translucent or opaque

Not fi lterable Filterable

Do not settle out Settle out

Small crystalloid solutes Large particles

Objective 15

Identify an important diff erence between crystalloids and colloids.

Reading Assignment: Refer to Objective 14. Also, Read Text, Page 40.

Dialyzing membranes are semipermeable membranes with very small pores (just large enough, for example, to allow water, ions, and small molecules to pass through).Filter paper has larger pores.


Crystalloids: Small enough to pass through fi lter paper and small enough to pass through a dialyzing membrane.

Colloids: Small enough to pass through fi lter paper but not small enough to pass through a dialyzing membrane.

Objective 16

Defi ne dilute solution, concentrated solution, and specify the procedure for preparing a percent solution.

Reading Assignment: Refer to Objective 14. Also, Read Text, Page 40.

Solute: The material that is dissolved (present in smaller amounts).

Solvent: The liquid (present in largest amount).

Preparing a % solution:

Use % desired (in grams) of solute make it up to 100 ml with solvent

Preparing a .9% NaCl solution:Use 0.9 grams of NaCl (salt) add enough H

20 to make 100 ml

Prepare a 5% glucose solution: Use 5.0 grams of sugar add enough H

2O to make 100 ml

Objective 17

Defi ne acid, base, and salt. Given the ions formed when a compound is placed in water, classify it as an acid, a base, or a salt.

Reading Assignment: Read Text, Pages 40-41; Refer to Figure 2.9.

Acid: Substances that dissociate in water and produce hydrogen ions (H+)

Base: Substances that dissociate in water and produce hydroxide ions (OH-)

Salt: Substances that dissociate in water and produce an anion and a cation, neither of which is H+ or OH-

HCl H+ + Cl- H2SO

42H+ + SO

4–2

acid acid base

NaOH Na+ + OH- H2CO

3H+ + HCO

3-

base acid base


NaCl Na+ + Cl-

salt

Objective 18

Defi ne pH.

Reading Assignment: Read Text, Pages 40-41. Refer to Figure 2.9.

pH: A measure of acidity/alkalinity or a measure of the concentration of H+ in a solution.

More H+ means more acidic.Less H+ means more alkaline (basic).

Objective 19

Given the pH of a solution, categorize it as acidic, basic, or neutral.

Reading Assignment: Read Text, Pages 40-41; Refer to Figure 2.9.

A pH of 7.0 is neutral.A pH less than 7.0 is acidic.A pH greater than 7.0 is alkaline (basic).

Objective 20

Given the pH of two solutions, compare their relative hydrogen ion concentrations.


If the pH values diff er by 1 unit, that means they diff er by 101.If the pH values diff er by 2 units, that means they diff er by 102.If the pH values diff er by 3 units, that means they diff er by 103.


Objective 21

Distinguish between “neutral” pH and the “average” pH range of the blood.

Reading Assignment: Refer to Figure 2.9, Page 40-41.

Average pH of the blood: 7.35-7.45Neutral pH: 7.00 (neutral pH is defi ned very precisely)

Objective 22

Defi ne weak acid, strong acid, and contrast their disruptive eff ects on human physiology.


Strong acids dissociate completely. Weak acids do not. Weak acids are buff ers.

Strong acid: H2SO

4 2H+ + SO

4

sulfuric acid sulfate

Weak acid: H2CO

3H+ + HCO

3-

carbonic acid bicarbonate

Notice the arrows in the above reactions.Strong acids have a more disruptive eff ect on physiology (assuming similar concentration). The more hydrogen ions released, the more acidic the solution.

Objective 23

Explain the function of buff er systems.

Reading Assignment: Read Text, Pages 41.

Chemical compounds that can convert strong acids or bases into weak acids or bases are called buff ers.

HCO3 - (weak base) and H2CO3 (weak acid) are buff ers in our bodies

H+ + HCO3- H2CO3 H

2O + CO

2

pH buffers: strong acids weak acids strong bases weak bases


Objective 24

Distinguish physical processes versus chemical reactions.

Reading Assignment: No additional sources

Physical processes include dissolving, fi ltering, freezing, melting and boiling. These are pro-cesses that do not change the basic composition of the substance (no new substances are formed). Physical processes usually change the appearance of the substance. Chemi cal processes (i.e., chemical reactions), such as burning a piece of wood and rust formation, are reactions that form new substances. Burning changes wood and oxygen into smoke, ashes, etc. Rust formation changes metal and oxygen into rust (metal oxides).

Physical Processes: Example: freezing of water.1. No new substances are formed.2. No major bonds are changed.

Chemical Processes: Example: rust forming.1. New substances are formed.2. Major bonds are changed.

Objective 25

Describe oxidation/reduction (Redox) reactions.

Reading Assignment: No additional information.

Oxidation: Removal of an electron (e-) during a chemical reaction loses e-, energy, and H+

Reduction: Gain of an electron (e-) during a chemical reaction gains e-, energy, and H+

Whenever an oxidation occurs, a reduction must also occur...this pair of oxidation/reduc-tion reactions is called a Redox reaction. You may fi nd the mnemonic "OIL RIG" helpful:

Oxidation Is Loss

Reduction Is Gain


Objective 26

Given a chemical reaction, classify it as a synthesis, decomposition, exchange, or reversible reaction.


Decomposition:catabolism, exergenic (energy out)

2 H2O 2 H

2+ O

2

Synthesis:anabolism, endergonic (energy in)

2 H2

+ O2

2 H2O

Exchange (Replacement): HClacid

+ NaOHbase

H2O

water+ NaCl

salt

Reversible: 2 H2O 2 H

2+ O

2

evaporation

rain

NaCl Na+ + Cl–

Objective 27

Identify the four most abundant elements in living systems.

Reading Assignment: Refer to Table 2.1 Page 28.

The four most abundant elements in living systems are carbon, hydrogen, oxygen and nitrogen. Their percentages in living systems are as follows:

Most abundent elements: C H O N

Percent of living systemsBy number of 9.5 63 25.5 1.4

By mass 18 10 65 3


Objective 28

Distinguish inorganic from organic compounds.

Reading Assignment: Read Text, Pages 37, 42.

Inorganic Compounds

Do not (usually) contain carbon.Are small in size. Include mainly ionic bonds.Are not very fl ammable.Are usually H

2O soluble (hydrophilic)

Examples: H2O, CO, CO

2,

minerals, simple salts such as NaCl, simple acids/bases such as HCl & NaOH

Organic Compounds

Contain carbon.Are frequently large in size.Include mainly covalent bonds. Usually fl ammable.May not be H

2O soluble (hydrophobic).

Examples: fat, carbohy-drates, proteins, nucleic acids, vitamins, complex acids/bases

Objective 29

Given any of the properties of water, provide an example of how this property is utilized in the human body.

Reading Assignment: Read Text, Pages 37-39; Refer to Figure 2.8, Page 38.

Properties of water: Solvent/suspending mediumParticipates in chemical reactionsHeat buff erEvaporative coolingLubricant

Water can absorb a large amount of heat (i.e., high heat capacity). As water evaporates from the skin it reduces body temperature (energy loss or removal due to heat of vaporization). Water is composed of polar covalent bonds, which makes it a versatile solvent. Water participates in chemical reactions. In hydrolysis chemical reactions, adding water breaks or "lyses" chemical bonds. In dehydration synthesis reactions, water is removed, and new molecules are formed. Water has a very high surface tension and is a great lubricant.


Answers to Self-Test are at the end of this module.

Self-Test #1

1. Anything that has mass and occupies space is called _________________.

2. All matter is composed of basic building blocks called ________________.

3. There are _________________ diff erent kinds of elements.

4. Water is composed of the two elements called _________________ and _________________.

5. If you took a bar of pure lead (lead is an element) and cut it into pieces so small that they could not be broken up any more and still be lead, you would have _________________ of lead.

6. _________________ are the smallest part of elements.

7. Since there are at least 112 diff erent elements, there are also at least 112 diff erent kinds of _________________.

Self-Test #2

Replace each boldface item with the corresponding name or symbol.

1. I _________________ is important to the functioning of the thyroid gland.

2. Fe _________________ is an important part of hemoglobin.

3. Bridges are built from iron ________.

4. Healthy bones need plenty of calcium ________.

5. C _________________is found in every organic molecule. 6. Milk and dairy products are rich in Ca _________________.7. Na _________________ is needed to maintain your body’s water balance.

8. Carbon ________ is an important part of organic molecules.

9. Babies used to get Pb _________________ poisoning by chewing the paint off of their cribs.

10. Sodium ________ and chloride make up table salt.


11. Cl _________________ is added to swimming pools to kill bacteria.

12. Without potassium ________ your nerves and muscles cannot function.

13. H _________________ and O _________________make up water.

14. Rotten eggs smell like sulfur ________.

15. Newborn babies have drops of silver ________ nitrate put in their eyes.

Self-Test #3

1. Atoms are composed of three kinds of subatomic particles called _________________, _________________, and _________________.

2. Particles in the nucleus which have a neutral charge are called _________________.

3. Particles in the nucleus which have a positive charge are called _________________.

4. Particles orbiting around the nucleus with a negative charge are called ________________.

5. Protons have a _________________ charge and are in the nucleus.

6. Neutrons are found in the _________________ and have a _________________ charge.

7. Electrons are (inside/outside) the nucleus and carry a _________________ charge.

8. The notation for proton is _________________. 6 p+ means _________________.

9. The designation or notation for ten electrons is written _________________.

10. The designation for neutron is _________________.

Self-Test #4

Use the Periodic Chart and fi ll in the following blanks.

1. Hydrogen has _________________ proton(s).

2. Non-ionized hydrogen has _________________ electron(s).

3. Lithium has _________________ proton(s).

4. The atomic number of carbon is _________________.

5. Carbon has _________________ proton(s).


6. Non-ionized potassium has _________________ electron(s).

7. Sodium has _________________ proton(s).

8. The atomic number for sodium is _________________.

9. Non-ionized sodium has _________________ electron(s).

Self-Test #5

Fill in the chart below for the three isotopes for Mg. Remember, the # of protons is the same for all isotopes of a given element.

1 2 3

atomic number

number of protons

number of neutrons 14 no

mass number 24 25

atomic weight for Mg is _______ . (Hint: look at Periodic Table).

Fill in the chart for the four isotopes of iron.

1 2 3 4

atomic number

number of protons

number of neutrons 30 no 32 no

mass number 54 57

Average atomic weight for all Fe atoms is ______.

Self-Test #6

1. Anything that has mass and occupies space is _______________.

2. Matter is composed of building blocks called _______________.

3. The smallest part of an element that still has the characteristics of that element is called an/a _______________.

4. There are _______________diff erent elements.


Find the correct symbol for each element.

____ 5. calcium a. Ca____ 6. potassium b. Mg____ 7. iron c. K____ 8. carbon d. Fe____ 9. magnesium e. C

____ 10. silver a. Ag____ 11. lead b. Pb____ 12. phosphorus c. F____ 13. sodium d. P____ 14. fl uorine e. Na____ 15. sulfur f. S

16. Atoms are composed of three particles called _______________, _______________, and _______________.

17. Subatomic particles within the nucleus and with a positive charge are called _______________.

18. no designates a particle within the nucleus that has a _______________ charge.

19. A particle orbiting the nucleus and designated e- is called an/a _______________.

20. The particle in item #19 has a _______________ charge.

21. The atomic number tells you how many _______________ an atom contains.

22. You can also tell how many electrons an unreacted atom has because the number of

electrons equals the number of _______________.

How many protons do each of the following have?

Which elements have the following atomic number?

23. Li ____ 28. 9 _______________

24. Na ____ 29. 1 _______________

25. C ____ 30. 47 _______________

26. O ____ 31. 19 _______________

27. Ca ____ 32. 15 _______________

33. The mass number is equal to the sum of __________ and ________ .

34. Which element has an atomic weight of 22.9898? _______________ .


35. How many no does it have? _______________ .

36. How many e- does it have? ________________ .

37. The reason some of the atomic weights are not whole numbers is because many ele ments have _________________ .

38. ______________ are atoms of an element which have the same number of protons, but diff erent numbers of neutrons.

Self-Test #7

Electron Confi guration

1. The fi rst energy level or electron shell can hold _______________ electrons.

2. The second energy level can hold _______________ electrons.

3. The way an atom reacts with another atom depends on the number of electrons in the _______________ shell.

4. Group I elements have _______________ electron(s) in their outer shell. They will want to (donate/accept) electrons.

5. If an atom gives away one electron, it will have a net charge of _______________ .

6. Group VII elements have _______________ electron(s) in their outer shell. They will want to (donate/accept) electron(s).

7. If an atom accepts one electron, it will have a net charge of _______________ .

8. An atom that has donated or accepted an electron is called an _______________ .

9. When two or more atoms are held together in a molecule due to attraction of opposite charges, it is called an _______________ . 10. A neutral phosphorous atom has _______________ protons and _______________ electrons. It has _______________ electrons in its outer shell. It wants to (donate/ accept) _______________ electrons. The resulting ion of phosphorous will then have a net charge of _______________ .

11. A neutral calcium atom has _______________ electrons in its outer shell. It wants to (donate/accept) _______________ electrons. An ion of calcium has a net charge of _______________ .


12. A positive ion is called an/a _______________ .

13. A negative ion is called an/a _______________ .

Self-Test #8

For each compound listed below, list by name and symbol the elements of which it is composed. After each element write the number of atoms of that element that are found in the compound and the molecular weight of the compound. The answer to the fi rst problem is provided.

1. NaHCO3 (sodium bicarbonate):

sodium (Na), 1 ; carbon (C), 1; hydrogen (H), 1; oxygen (O), 3. MW = 84

2. HI (hydrogen iodide):

3. Na2SO

4 (sodium sulfate):

4. C2H

6 (ethane gas):

5. KOH (potassium hydroxide):

6. MgCl2 (magnesium chloride):

7. HNO3 (nitric acid):

8. (NH4)

2 CO

3 (ammonium carbonate):

Self-Test #9

Find the polyatomic ion in each of the molecules in the table below. Write its symbol and its name in the proper column. The fi rst exercise has been completed for you. If you need help in remembering, check the table of common polyatomic ions in Objective 12.


Molecule Polyatomic ion Name of ion

1. Na3PO

4 PO

4- 3

phosphate

2. KNO2 NO

2- __________ _________

3. NaHCO3 _________ ____ __________ _________

4. (NH4)

2CO

3 _________ ____ __________ _________

5. MgSO4 _________ ____ __________ _________

6. Ca(OH)2

_________ ____ __________ _________

7. FeCO3 _________ ____ __________ _________

8. HNO3 _________ ____ __________ _________

9. CaSO3

_________ ____ __________ _________

ANSWERS TO SELF-TESTS

Self Test #11. matter 5. atoms2. atoms/elements 6. atoms3. about 112 7. atoms4. hydrogen, oxygen

Self Test #21. iodine 10. Na2. iron 11. chlorine3. Fe 12. K4. Ca 13. hydrogen, oxygen5. carbon 14. S6. calcium 15. Ag7. sodium8. C 9. lead

Self Test #31. neutrons, protons, electrons 6. nucleus, neutral2. neutrons 7. outside, negative3. protons 8. p+, six protons4. electrons 9. 10e-

5. positive 10. no


Self Test #41. one 6. 192. one 7. 113. three 8. 114. six 9. 115. six

Self Test #5 For isotopes of Mg For isotopes of ironatomic number 12 12 12 26 26 26 26# of protons 12 12 12 26 26 26 26# of neutrons 12 13 14 28 30 31 32mass number 24 25 26 54 56 57 58atomic weight 24.30 55.84

Self Test #61. matter 11. b 19. electron 29. H2. atoms/elements 12. d 20. negative 30. Ag3. atom 13. e 21. protons 31. K4. about 112 14. c 22. protons 32. P5. a 15. f 23. 3 33. protons, neutrons6. c 16. neutrons 24. 11 34. Na7. d protons 25. 6 35. 128. e electrons 26. 8 36. 119. b 17. protons 27. 20 37. isotopes10. a 18. neutral 28. F 38. isotopes

Self Test #71. 2 7. 1 11. 2, donate, 2, +22. 8 8. ion 12. cation3. outer 9. an ionic bond 13. anion4. 1, donate 10. 15, 15, 5 5. +1 accept6. 7, accept 3, -3

Self Text #81. Answer already provided.

2. Hydrogen (H), 1; Iodine (I), 1 MW = 128

3. Sodium (Na), 2 ; Oxygen (0), 4 ; Sulfur (S), 1 MW = 142


4. Carbon (C), 2 Hydrogen (H), 6 MW = 30

5. Potassium (K), 1 ; Hydrogen (H), 1; Oxygen (0), 1 MW = 56

6. Magnesium (Mg), 1; Chlorine (Cl), 2 MW = 94

7. Hydrogen (H), 1; Oxygen (0), 3; Nitrogen (N), 1 MW = 63

8. Nitrogen (N), 2 ; Carbon (C), 1; Hydrogen (H), 8 ; Oxygen (0), 3 MW = 96

Self Test #9

1. Answer already provided.2. - - - nitrite 3. HCO

3- bicarbonate

4. NH4

+& CO3

- 2 ammonium & carbonate 5. SO

4- 2 sulfate

6. OH- hydroxide7. CO

3- 2 carbonate

8. NO3

- nitrate9. SO

3- 2 sulfi te

final module 2.2007.sep - weber state university … · ... has mass and occupies space. ... of any...

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