Matter

Is it uniform throughout

HeterogeneousMixture

HomogeneousSubstance

Pure SubstanceHomogeneous

Mixture

Element Compound

Can it be separatedby physical means?

Can it be broken intoSimpler substances By chemical means?

Classification of Matter

Consists of two or more substancesthat are easily distinguishable.

Ex. soil, milk, & granite

YES

YES

YES

NO

NO

NO

Physical Separationsdo not change thecomposition of the

Matter.Homogeneous mixturesare also called solutions.Two substances that are

mixed in equal proportionsthroughout.

Ex. Air, Kool-Aid

During a chemical changethe composition of the substance is altered. It

does not possess the samecharacteristics as before.Elements are substances that

are composed of one type ofatom. These are the simplest

units of matter that are able toretain all the properties of

that matter.Ex. Gold, carbon, silver

Compounds are substances Containing atoms that are

Chemically bonded to each Other. Separating the atomsRequires a chemical change.

Ex. Salt (sodium chloride), water,Oxygen (as found in the nature)

PHASE DIAGRAMS

This is a phase-change diagram for water.

1. Along LEG ‘A’ water exists as a solid (ice), and the temperature increases as HEAT energy is

absorbed. 2. At 0 °C a phase change begins:

a) Moving from left to right along LEG ‘B’, ice is melting to form liquid water b) Moving from right to left along LEG ‘B’, liquid water is freezing to form ice c) The distance of LEG ‘B’ along the Heat axis (x-axis) is known as the Heat of Fusion d) Note that temperature remains constant during a phase change!

3. Once ice has completely melted, the temperature begins to increase again (LEG ‘C’), as the heat absorbed by water is no longer going toward changing the phase of the substance.

4. At 100 °C, a second phase change begins:

a) Moving from left to right along LEG ‘D’, water is boiling to form water vapor b) Moving from right to left along LEG ‘D’, water vapor is undergoing condensation to form

liquid water c) The distance of LEG ‘D’ along the Heat axis (x-axis) is known as the Heat of

Vaporization d) Note that temperature remains constant during a phase change!

5. Once all of the liquid water has vaporized, the temperature begins to increase again (LEG ‘E’), as the heat absorbed by water is no longer going toward changing the phase of the substance.

This a phase diagram for water, showing all three phases. The lines represent conditions of equilibrium between phases.

Phase Changes From…. To…. …is called …and energy is…

Solid Liquid Melting Absorbed Liquid Solid Freezing Released Liquid Vapor Boiling or vaporization Absorbed Vapor Liquid Condensation Released Solid Vapor Sublimation Absorbed Vapor Solid Deposition Released

1. The temperature at which a substance melts/freezes at standard pressure (1 atm) is known as

the Normal melting point or Normal freezing point. For water, this is 0 °C. 2. The temperature at which a substance boils/condenses at standard pressure is known as the

Normal boiling point or Normal condensation point. For water, this is 100 °C. 3. The Triple point is the condition of temperature and pressure in at which all three phases exist

together at equilibrium. For water, this is 0.0099 °C and 0.006 atmospheres. 4. The Critical Temperature, Tc is the temperature beyond which the solid and liquid phases of

the substance cannot exist. Put another way, above the critical temperature, the substance can only be found as a gas. For water, this temperature is 373.99 °C.

5. The Critical Pressure, Pc is the pressure above which the substance cannot exist as a gas. For water, this pressure is 217.75 atmospheres.

6. The Critical Point is the point defined by the critical temperature and the critical pressure. 7. The slope of the line between the solid and liquid phase provides important information about

the substance: a) If the slope is negative (as it is for water), then the substance is more dense as a liquid

than it is as a solid. b) If the slope is positive (as it is for most substances), then the substance is more dense

as a solid than it is as a liquid.

Notes on Mass, Volume and Density Mass measures the amount of _____________ in an object. Solids, liquids and gases all have mass because they are made of matter. The SI unit for mass is the _____________. The mass of an object is measured using a triple beam _____________. Don't confuse mass with weight! Mass is a property of an object and does not depend on the object's __________. Weight is a measure of the force of _______________ acting on an object and therefore can vary depending on the object's location. Volume is the amount of space that an object occupies. Solids, liquids and gases all have volume. There are 3 methods to determine volume: 1) To determine the volume of a liquid use a _____________ ________________. 2. To determine the volume of an object with an irregular shape subtract ________ volume – _________ volume. 3. To determine the volume of a rectangular solid, multiply length x width x height. The unit will be ________________ Remember: 1 ml of water is equal to 1 cm3 meniscus-curved surface of water in a graduated cylinder Read the volume at the ___________________ of the meniscus Density is the measurement of the mass of an object divided by its volume Density unit is generally g/cm3 or g/mL (remember that 1 mL = 1 cm3) Two units are needed because mass is measured in grams and volume in cm3 or mL Two objects can have the same ______________________ but different _____________ ex: empty box vs full box

Notes on Mass, Volume and Density Two objects can have the ____________________________but differ in __________________ ex

Density is also a physical ____________________. Density can be determined using the formula Density = Mass/Volume (D=M/V) Lets practice: Find the density of a book that has a mass of 35 g and a volume of 7 cm3 Find the density of a bottle that has a mass of 125 g and a volume of 25 mL A boat has a mass of _____________ and a volume of _____________________. Will this boat float? The density of water is 1 g/cm3 Hint: an object that is more dense than water will sink. Mystery Density Find the density of your book. What will we need? Hint: D=M/V SHOW ALL WORK HERE:

History of the Atom

John Dalton (1766 – 1844):

John Dalton was an English chemist. His ideas form the atomic theory of matter. Here are his ideas.

• All elements are composed (made up) of atoms. It is impossible to divide ordestroy an atom.

• All atoms of the same elements are alike. (One atom of oxygen is like anotheratom of oxygen.)

• Atoms of different elements are different. (An atom of oxygen is different from anatom of hydrogen.)

• Atoms of different elements combine to form a compound. These atoms have to bein definite whole number ratios. For example, water is a compound made up of 2atoms of hydrogen and 1 atom of oxygen (a ratio of 2:1). Three atoms of hydrogenand 2 atoms of oxygen cannot combine to make water.

1. What is the name of John Dalton’s theory? _____________________________________

2. What are elements made of? ________________________________________________

3. An atom of hydrogen and an atom of carbon are _________________________________.

4. What are compounds made of? _______________________________________________

5. The ratio of atoms in HCl is: a) 1:3 b) 2:1 c) 1:1

J. J. Thompson (Late 1800s):

J. J. Thompson was an English scientist. He discovered the electron when he was experimenting with gas discharge tubes. He noticed a movement in a tube. He called the movement cathode rays. The rays moved from the negative end of the tube to the positive end. He realized that the rays were made of negatively charged particles – electrons.

1. What did J.J. Thompson discover? _____________________________________________

2. What is the charge of an electron? ____________________________________________

3. What are cathode rays made of? ______________________________________________

4. Why do electrons move from the negative end of the tube to the positive end?__________________________________________________________________________

5. What was Thompson working with when he discovered the cathode rays?__________________________________________________________________________

Lord Ernest Rutherford (1871 – 1937):

Ernest Rutherford conducted a famous experiment called the gold foil experiment. He used a thin sheet of gold foil. He also used special equipment to shoot alpha particles (positively charged particles) at the gold foil. Most particles passed straight through the foil like the foil was not there. Some particles went straight back or were deflected (went in another direction) as if they had hit something. The experiment shows:

• Atoms are made of a small positive nucleus; positive nucleus repels (pushes away) positive alpha particles

• Atoms are mostly empty space

1. What is the charge of an alpha particle? _______________________________________

2. Why is Rutherford’s experiment called the gold foil experiment? _____________________

__________________________________________________________________________

3. How did he know that an atom was mostly empty space? __________________________

__________________________________________________________________________

4. What happened to the alpha particles as they hit the gold foil? _____________________

__________________________________________________________________________

5. How did he know that the nucleus was positively charged? _________________________

__________________________________________________________________________

Niels Bohr (Early 1900s):

Niels Bohr was a Danish physicist. He proposed a model of the atom that is similar to the model of the solar system. The electrons go around the nucleus like planets orbit around the sun. All electrons have their energy levels – a certain distance from the nucleus. Each energy level can hold a certain number of electrons. Level 1 can hold 2 electrons, Level 2 - 8 electrons, Level 3 - 18 electrons, and level 4 – 32 electrons. The energy of electrons goes up from level 1 to other levels. When electrons release (lose) energy they go down a level. When electrons absorb (gain) energy, they go to a higher level.

1. Why could Bohr’s model be called a planetary model of the atom? __________________

__________________________________________________________________________

2. How do electrons in the same atom differ? _____________________________________

__________________________________________________________________________

3. How many electrons can the fourth energy level hold? ____________________________

4. Would an electron have to absorb or release energy to jump from the second energy level to the third energy level? _____________________________________________________

5. For an electron to fall from the third energy level to the second energy level, it must ___________________________________ energy.

Properties of Ionic, Covalent, and Metallic Compounds

Ionic Compounds Covalent Compounds Metallic Compounds -Formed from a combination of metals and nonmetals. -Electron transfer from the cation to the anion. -Opposite charged ions attract each other.

-Formed from a combination of nonmetals. -Electron sharing between atoms.

-Formed from a combination of metals -“sea of electrons”; electrons can move among atoms

Solids at room temperature Can be solid, liquid, or gas at room temperature.

Solids at room temperature

High melting points Low melting points Various melting points

Dissolve well in water Do not dissolve in water Do not dissolve in water.

Conduct electricity only when dissolved in water; electrolytes

Do not conduct electricity; nonelectrolytes

Conduct electricity in solid form.

Brittle, hard Soft Metallic compounds range in hardness. Group 1 and 2 metals are soft; transition metals are hard. Metals are malleable, ductile, and have luster.

*MOST compounds are a mixture between ionic and covalent

Chapter 6.3 & 6.4 Atomic/Ionic Radius Worksheet Name: _________________________Chemistry2 points Date: _____________ Hour: _______

Atomic Radius (for help, see pp.187 to 189 in the textbook)

1. What is meant by the term atomic radius? ______________________________________

________________________________________________________________________

2. What is the trend across a row for atomic radius? ________________________________

________________________________________________________________________

3. What is the trend down a column for atomic radius? ______________________________

_________________________________________________________________________

4. Predict which atom has a smaller atomic radius.

(a) magnesium or chlorine (circle one)

_____ same column or _____ same row (check one)

(b) strontium or barium (circle one)

_____ same column or _____ same row (check one)

(c) vanadium or niobium (circle one)

_____ same column or _____ same row (check one)

5. Predict which atom has a larger atomic radius.

(a) potassium or calcium (circle one)

_____ same column or _____ same row (check one)

(b) neon or argon (circle one)

_____ same column or _____ same row (check one)

(c) cesium or polonium (circle one)

_____ same column or _____ same row (check one)

6. Put the following elements in order from largest atom to smallest atom: carbon, fluorine,

beryllium, and lithium __________________________________________________

7. Put the following elements in order from smallest to largest: aluminum, indium, boron

__________________________________________________________________________

Ionic Radius (for help, see pp. 189-191 in the textbook)

8. What is an ion? ___________________________________________________________

_________________________________________________________________________

9. A positive ion has ______________ an electron. It is also called a __________________.

10. A negative ion has ______________ an electron. It is also called an _________________.

11. When an atom forms a positive ion, it gets ________________________.

12. When an atom forms a negative ion, it gets ________________________.

13. Metals tend to _________________ electrons and form ___________________ ions.

14. Nonmetals tend to _________________ electrons and form ___________________ ions.

15. Predict which particle is larger:

(a) Rb or Rb (circle one)+

p _____ e _____ p _____ e _____+ - + -

(b) I or I (circle one)-

p _____ e _____ p _____ e _____+ - + -

16. Predict which particle is smaller:

(a) Al or Al (circle one)3+

p _____ e _____ p _____ e _____+ - + -

(b) S or S (circle one)2-

p _____ e _____ p _____ e _____+ - + -

Review

17. Rows on the periodic table are also called ____________________ or

____________________ or ____________________.

18. Columns on the periodic table are also called ____________________ or

____________________.

19. There are _______________ rows and _______________ columns on the periodic table.

20. (a) The most stable electron configurations are ______________________________.

(b) The second most stable electron configurations are ________________________.

(c) The third most stable electron configurations are __________________________.

1. Draw the Bohr's Model for Elements 1 through 20 in the correct box on the periodic table. 2. Include the number of protons, neutrons, and electrons. Atomic Number = number of protons and electrons, Atomic Mass - Atomic Number = number of neutrons

1. Draw the Lewis Dot Diagram for Elements 1 through 20 in the correct box on the periodic table. 2. For each element write the valance electron number and the oxidation number.

1. Draw the Orbital Diagram for Elements 1 through 20 in the correct box on the periodic table. 2. In each box write it's electron configuration.

Ch. 19:4 Name: _____________________

Period:_____________________

Copyright © 2003, C. Stephen Murray www.aisd.net/smurray

Naming Compounds

What’s it Made of?

Metal and non-metal

2 non-metals

3 or more elements

polyatomic compound

covalent compound

ionic compound

USE “- IDE” ENDING (NO PREFIXES!)

Name the metal and non-metal and change the ending to “ide”.

USE GREEK PREFIXES

Put prefixes in front of element names to tell how many atoms are there.

Don’t use “mono” for first name,

but always for second name.

CHECK THE CHART BELOW (NO PREFIXES!)

Use the names on the chart.

If the polyatomic ion is the cation end the second name with “-ide”.

Li2S

Metal and non-metal— ionic

Lithium Sulfide

(not dilithium sulfide— no prefixes for ionic compounds)

N2O4

2 non-metals—covalent

(di =2 and tetra =4)

“Dinitrogen tetroxide”

NaNO3

3 elements — polyatomic

Check chart (see below) Na - sodium

NO3 - nitrate (on chart)

Sodium nitrate

Polyatomic Ions

Oxidation # Name Formula

1+ ammonium NH4+

1- acetate C2H3O2-

2- carbonate CO32-

2- chromate CrO42-

1- hydrogen carbonate

HCO31-

1+ hydronium H3O+

1- hydroxide OH1-

1- nitrate NO31-

2- peroxide O22-

3- phosphate PO43-

2- sulfate SO42-

2- sulfite SO32-

Greek Prefixes

Mono - 1

Di – 2 Tri – 3

Tetra – 4 Penta – 5

Hexa – 6 Hepta – 7 Octa – 8 Nona – 9 Deca – 10

Exception— O2 is “peroxide” and can make polyatomic com-pounds with only 2 ele-

ments! O2 with a non-metal is dioxide. O2 with a metal OR Hydrogen (acting as a

metal) is peroxide.

Why are ionic compounds so easy to name? Because most ionic com-

pounds can only form one way, using the oxidation numbers. In covalent compounds, though, non-metals can sometimes combine in multiple ways (carbon monoxide; carbon dioxide).

So, covalent compounds use prefixes.

How to remember prefixes:

Monorail – one rail train Monocle – glasses for one eye;

single lens (Colonel Klink). Dilemma – struggle between 2 choices. Tricycle – 3 wheels

Pentagon – 5 five sided military building in Washington, D.C.

Octopus – 8 legs Decade – 10 years

Transition Metals Can Have More Than One Oxidation Number

Iron (II) has an oxidation number of 2+

Iron (III) has an oxidation number of 3+. When naming them you must specify

WHICH ONE.

FeO—Iron (II) oxide Fe2O3— Iron (III) oxide

How to use this chart— Determine what the

compound is made of and follow the arrows. The

chart will tell you how to name the compound.

Ch. 19:4 Name: _____________________

Period:_____________________

Copyright © 2003, C. Stephen Murray www.aisd.net/smurray

Use the Polyatomic Ion Chart on the front of the worksheet to name these Polyatomic Ions:

HCO31-

SO42-

O22-

SO32-

NO31-

NH4+

CrO42-

OH1-

PO43-

CO32-

Hydrogen carbonate

___________________

___________________

___________________

___________________

___________________

___________________

___________________

___________________

___________________

Metal or Non-metal?

Ionic or Covalent?

Iron Oxide Ionic Barium Chloride ____________ Carbon Dioxide ____________ Magnesium Oxide ____________ Aluminum Fluoride ____________ Nitrogen Tribromide ____________ Chromium Fluoride ____________ Potassium Oxide ____________

M N

Name These Ionic Compounds

MgF2 Magnesium Fluor-ide

Li2O Lithium Ox- __________________

NaCl Sodium Chlor- ________________

K2O Potassium Ox-_________________

CaS _______________ Sulf- _________

BeI2 _______________ Iod-__________

AlBr3 _______________ Brom-________

CaF2 ____________________________

MgO ____________________________

LiCl ____________________________

Classify and Name These Compounds

1. BaCl2 Ionic __ Barium chloride _

2. CO _____________ _______________________________

3. Ag2O _____________ _______________________________

4. K2SO4 _____________ _______________________________

5. MgBr2 _____________ _______________________________

6. SO3 _____________ _______________________________

7. P2O4 _____________ _______________________________

8. Be(CrO4) _____________ _______________________________

9. LiF _____________ _______________________________

11. CO2 _____________ _______________________________

12. OF2 _____________ _______________________________

Ionic, Covalent, or Polyatomic Name

Define these Greek Prefixes

Penta = ______ Nona = ______ Mono = ______ Octa = ______ Tri = ______

Tetra = ______ Hexa = ______ Hepta = ______ Deca = ______

Di = ______

1. CO2

2. C2O4

3. C3O5

4. CO

5. C2O

6. CO8

A. Carbon monoxide

B. Carbon dioxide

C. Dicarbon monoxide

D. Tricarbon pentoxide

E. Dicarbon tetroxide

F. Carbon octoxide

Si2O3 Disilicon _____oxide

N3Cl4 _____nitrogen tetrachloride

SO2 Sulfur _____oxide

PO5 Phosphorous ______ox____

S2F4 ____sulfur _____fluor____

Name These Covalent Compounds

Name these Polyatomic Compounds (Remember — no prefixes!)

CaSO4 Calcium _________________

K2CO3 ________________ carbonate

CuNO3 Copper (I) ________________

NH4Cl _________________ chloride

Mg(NO3)2 Magnesium _______________

K3PO4 Potassium _________________

Li2(CrO4) Lithium _____________________

Mg(OH)2 M___________ H_____________

Al(PO4) A______________ P___________

K(NO3) _____________ ______________

Ca2SO3 _____________ ______________

Name: _________________ Reaction Types Worksheet I Instructions: i. Balance the equations ii. Classify each of the following reactions (i.e. synthesis, decomposition, single displacement, double displacement, or combustion)

1. ___ KBr + ___ Cl2 à 2 KCl + ___Br2

2. ___ Ca(OH)2 + ___ H2SO4 à ___ CaSO4 + ___ H2O

3. ___ Zn + ___ H2SO4 à ___ ZnSO4 + ___ H2 4. ___ AgNO3 + ___ NaCl à ___ AgCl + ___ NaNO3 Reaction Types Worksheet II Instructions: i. Write the equation ii. Balance the chemical equations iii. Classify each of the following reactions 1. aluminum oxide à aluminum + oxygen 2. carbon tetrafluoride + bromine à carbon tetrabromide + fluorine 3. mercury iodide + oxygen à mercury oxide + iodide 4. sodium + oxygen à sodium oxide 5. barium chloride + magnesium sulfate à barium sulfate + magnesium chloride

Finding  Empirical  Formula  given  grams    A  sample  of  a  brown  gas,  a  major  air  pollutant,  is  found  to  contain  2.34  g  N  and  5.34g  O.    Determine  a  formula  for  this  substance.    

1. Determine  the  mass  in  grams  of  each  element  present.  Molar  mass  of  N                                        Molar  mass  of  O  =14.0067g/mol  of  N                            =15.999g/mol  of  O    

2. Convert  grams  of  each  element  to  moles  2.34  g  of  N            1  mole  of  N            =  0.167  moles  of  N                              14.0067  g  of  N  

 5.34  g  of  O    1  mole  of  O                  =  0.334  moles  of  O                                          15.999  g  of  O        

3. Divide  each  by  the  smallest  number  of  moles  to  obtain  the  simplest  whole  number  ratio.  

   

 4. If  whole  numbers  are  not  obtained*  in  step  3),  multiply  through  by  the  

smallest  number  that  will  give  all  whole  numbers    • BE  CAREFUL!  DO  NOT  ROUND  OFF  NUMBERS  PREMATURELY!  

 Finding  Empirical  Formula  given  percent  composition  

1.     Drop  the  “%”  sign  and  add  a  “g”  2. Convert  grams  of  each  element  to  moles  4. Divide  each  by  the  smallest  number  of  moles  to  obtain  the  simplest  

whole  number  ratio.  3. If  whole  numbers  are  not  obtained*  in  step  3),  multiply  through  by  the  

smallest  number  that  will  give  all  whole  numbers    • BE  CAREFUL!  DO  NOT  ROUND  OFF  NUMBERS  PREMATURELY!  

 

N0.167 O0.334

0.167 0.334 20.167 0.167

N O NO=

Empirical Formula Worksheet Empirical Formula – a formula showing the smallest whole-number mole ratio.

1. Convert the grams of each element to moles. 2. To find the simplest ratio, divide each element’s mole value by the smallest mole value. This

will ensure a something to 1 ratio. 3. If the mole ratios are NOT all whole numbers, you must multiply a whole number in order to

obtain a whole number ratio. 4. It is helpful to recognize decimal forms of common fractions and multiply by the

denominator.

1. A 15.0g sample of a compound is found to contain 8.83g sodium and 6.17g sulfur. Calculate the empirical formula of this compound.

8.83g Na 1 mol Na = 0.384 mol Na 0.384 = 2 Na 23.0g Na 0.192

6.17g S 1 mol S = 0.192 mol S 0.192 = 1 S 32.1g S 0.192 Na2S

2. A compound is found to contain 36.48% Na, 25.41% S, and 38.11% O. Find its empirical formula.

36.48g Na 1 mol Na = 1.586 mol Na 1.586 = 2 Na 23.0g Na 0.7916

25.41g S 1 mol S = 0.7916 mol S 0.7916 = 1 S 32.1g S 0.7916

38.11g O 1 mol O = 2.382 mol O 2.382 = 3 O 16.0g O 0.7916 Na2SO3

3. Find the empirical formula of a compound that contains 53.70% iron and 46.30% sulfur. 6. Qualitative analysis shows that a compound contains 32.38% sodium, 22.65% sulfur, and 44.99%

oxygen. Find the empirical formula of this compound. Write the molecular formulas of the following compounds: A compound with an empirical formula of C2OH4 and a molar mass of 88 grams per mole What is the percent composition?

Atomic Weight Carbon Atomic weight is average of isotope example Carbon exists in the form of a carbon 12 atom, carbon 13 atom, and carbon 14 atom -98.89% of all carbon atoms are carbon 12 -1.11% of all carbon atoms are carbon 13 -The percentage of carbon 14 atoms present in all carbon atoms is so small that it plays no part in this calculation. -.9889 (12 amu) + .0111 (13 amu) = 12.01 -This is how the atomic weight is calculated. The Mole The number of atoms in 12 grams of carbon 12 atoms corresponds to Avogadro's Number, which is 6.02 * 1023. One mole of Mr Mack weighs 250. 2moles of Mr Mack weighs? .75 moles of Mr Mack weighs? Examples 1 mole = the total of the elements atomic masses expressed in grams ( 1 mole/amu(g) ) Mole x amu(g)/1 mole = mass(g)

.25 mol Ag = .25 (107.9 g) = 26.98 g If given g of atom, you can find how many moles of that substance Given(g) x 1 mole/amu (g)= moles 27.93 g Fe = 27.93 g/55.85 g = .5 mol Atom to moles. # atoms x 1 mole/6.02x1023 atoms = moles 3.01*1023 atoms K = (3.01*1023/6.02*1023) = .5 mol Moles to atoms. # moles x 6.02x1023 atoms/ 1 mole .25 mol Ag = .25 (6.02*1023) = 1.51*1023 atoms

Section 1: Directions: Determine the mol. of each sample. 1. 5.75 g Na 2. 17.6 g K

Section 2: Directions: Determine how much each sample weighs. 1. 1.98 mol Zn 2. .57 mol Rb

Section 3: Directions: Determine how many atoms of each sample there are. 1. 1.6 mol Na 2. .016 mol K 3. . Section 4: Directions: Determine how many moles of each substance there are. 1. 1.19*1023 atoms V 2. 3.43*1023 atoms Fe

Molecular Mass (Molecular Weight) How much one mole of Al2S3 (Aluminum Sulfide) weighs? So using this assumption, there are 2 moles of Al and 3 moles of S, so take 26.98 grams (which is the weight of 1 mole of Al) and multiply it by 2 (because there are two moles of Al in Al2S3). Then take 32.06 grams (which is the weight of 1 mole of S) and multiply it by 3 (because there are three moles of S in Al2S3). Then add the two amounts up. Al+3 = 2 mol Al+3 = 2 (26.98) = 53.96 g S-2 = 3 mol S-2 = 3 (32.06) = 96.18 g Total - 150.14 g Molecular Mass Examples

1. 1 mol HBr = 1 mol H and 1 mol Br 2. 3.5 mol Al3P = 10.5 mol Al and 3.5 mol P 3. .75 mol CaO = .75 mol Ca and .75 mol O

.75 (40.1 g Ca) = 30.08 g Ca

.75 (16 g O) = 12 g O Section 1: Directions: Determine the mass of each element in the substance.

1. .345 mol LiCl 2. 1.98 mol KBr

Section 2: Directions: Determine the number of moles of the compound from the mass of each element.

1. 42 g N, 48 g O2, in NO

Percent Composition Percent composition is defined as the mass of an element in a compound divided by the total weight of the compound. So if Bob wanted to find the percent composition of Aluminum and Sulfur in Aluminum Sulfide, he would: % Al = (g of Al/g of Al2S3) % S = (g of S/g of Al2S3) % Al = (53.96 g / 150.14 g) * 100 = 35.94% Al+3 % S = (96.18 g / 150.14 g) * 100 = 64.06% S-2 So, Aluminum is 35.94 percent by mass of Aluminum and Sulfur is 64.06 percent by mass. Examples

1. HBr - molar mass = 80.9 g H+ = 1.0 g; Br- = 79.9 g % H+ = (1 g H+/80.9 g HBr ) * 100 = 1.2 % % Br- = 79.9 g Br-/80.9 g HBr) * 100 = 98.8 %

Section 1: Directions: Determine the percent composition of each element in each substance.

1. MgS 2. LiCl

Solubility Trends !  The solubility of MOST solids increases

with temperature. !  The rate at which solids dissolve increases

with increasing surface area of the solid. !  The solubility of gases decreases with

increases in temperature. !  The solubility of gases increases with the

pressure above the solution.

Therefore… Solids tend to dissolve best when:

o Heatedo Stirredo Ground into small particles

Gases tend to dissolve best when: o The solution is coldo Pressure is high

Molarity  

1. What is the molarity of a solution that contains 1.724 moles of H2SO4 in 2.50 L of solution?

2. What is the molarity of a solution prepared by dissolving 25.0 g of HCl (g) in enough water to make 150.0 mL of solution?

3. How many grams of iron (IV) oxide are needed to make 5.6 liters of a 2.1 M solution?

                                                 

Molarity  

1. What is the molarity of a solution that contains 1.724 moles of H2SO4 in 2.50 L of solution?

2. What is the molarity of a solution prepared by dissolving 25.0 g of HCl (g) in enough water to make 150.0 mL of solution?

3. How many grams of manganese (IV) oxide are needed to make 5.6 liters of a 2.1 M solution?

 

 

Periodic  Table  

1  mol  

Periodic  Table  

Periodic  Table  

1  mol  

 

grams mol mol grams

liters

particles

l iters

particles  

Given   Find  

Use coefficients

Worksheet: Mixed Problems—Mole/Mole Name______________ and Mole/Mass

CHEMISTRY: A Study of Matter © 2004, GPB

8.13

Answer each of the following questions using the equation provided. BE SURE TO BALANCE EACH EQUATION BEFORE SOLVING ANY PROBLEMS. SHOW ALL WORK. 1. ___Cu + ___O2 ___CuO

a. If 101 grams of copper is used, how many moles of copper (II) oxide will be formed?

b. If 5.25 moles of copper are used, how many moles of oxygen must also be used?

c. If 78.2 grams of oxygen react with copper, how many moles of copper (II) oxide will be produced?

2. ___C4H10 + ___O2 ___CO2 + ___H2O

a. How many moles of butane, C4H10, are needed to react with 5.5 moles of oxygen?

b. How many grams of carbon dioxide will be produced if 3.5 moles of O2 react?

CHEMISTRY: A Study of Matter © 2004, GPB

8.14

3. ___Mg + ___HCl ___MgCl2 + ___H2

a. What mass of HCl is consumed by the reaction of 2.50 moles of magnesium?

b. What mass of MgCl2 is produced if 3.67 moles of HCl react?

c. How many moles of hydrogen gas are produced when 3.0 moles of magnesium react? 4. ___NH3 + ___O2 ___N2 + ___H2O

a. How many moles of oxygen react with 0.23 moles of NH3?

b. How many grams of water will be produced if 0.55 moles of oxygen react?

c. How many moles of nitrogen gas will be produced if 12.6 grams of ammonia react?