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Student name: _______________________________________________ Class: _______

SSttaattee GGooaall//EExxppeeccttaattiioonnss:: Goal: The student will demonstrate the ability to use scientific skills and processes to explain composition and interactions of matter in the world in which we live. Expectation(s): The student will explain that matter undergoes transformations, resulting in products that are different from the reactants. Indicator(s): The student will demonstrate that adjusting quantities of reactants in a chemical reaction may affect the amounts of products formed during a chemical reaction. Performance objectives:

Explain the mole concept, and convert between grams, moles, liters and number of atoms and/or molecules.

Determine mass percent composition of a sample from experimental data.

Determine mass percent composition of a compound from its formula.

Determine empirical formula of a compound from its mass percent composition.

Topic 1: The Mole & its relationships (Ch. 10-1 & 10-2)

Day 1: Measuring familiar quantities (10-1)

1) What are three methods for measuring the amount of something?

Book, Page 289; Problems 1 & 2

Study Guide, page 3; 1-12 Topic 2 – By count: Mole to particle relationship (10-1)

2) How is Avogadro’s number related to a mole of any substance?

Book, page 291; Problems 3 & 4

Book, page 292; Problems 5 & 6

Study Guide, page 4; 1-10

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Topic 3 – By mass: Molar mass (10-1)

3) How is the atomic mass of an element related to the molar mass of a substance?

Book, page 296; 7 & 8

Study Guide, page 5; 1-24 Topic 4 – By mass & count: Mass/Mole/Particle relationship (10-2)

4) How do you convert the mass of a substance to the number of moles of the substance?

Book, page 298; 16 & 17

Book, page 299; 18 & 19

Study Guide, page 6; 1-16 Topic 5 – By volume: Mass/Mol/Volume relationship (10-2)

5) What is the volume of a gas at STP?

Book, page 301; 20-21

Book, page 302; 22-23 Study Guide, Page 7; 23-30

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Practice: Familiar Quantities

Using diagram 1 and the information associated with it, answer the following questions.

1. How many oranges are in 8 bags of oranges?

2. How many bags do I need for 37 oranges?

3. What is the mass of the 8 bags of oranges?

4. If I had 24 kilograms of oranges, how many bags would I have?

5. How many bags of oranges would take up 23 dm3 of space?

6. About how many oranges is this?

Using diagram 2 and the information associated with it, answer the following questions.

7. How many grains of sugar are in 2 boxes of sugar?

8. How many boxes do I need to hold 1x1012 grains of sugar?

9. What is the mass of 4.3 boxes of sugar?

10. If I had 2.4 kilograms of sugar, how many boxes would I have?

11. How many boxes of sugar would take up 27 dm3 of space?

12. About how many grains of sugar is this?

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Practice: By Count (Mole ↔ Particle conversions) Use Avogadro’s number to make the following conversions:

1) How many molecules are there in 2 moles of FeF3?

2) How many molecules are there in 4 moles of Na2SO4?

3) How many moles are there in 2.3 x 1024 atoms of silver?

4) How many moles are there in 7.4 x 1023 molecules of AgNO3?

5) How many moles are there in 7.5 x 1023 molecules of H2SO4?

6) How many molecules are there in 1.2 moles of Cu(NO3)2?

7) How many moles are there in 9.4 x 1025 molecules of H2?

8) How many molecules are there in 2.30 moles of CoCl2?

9) How many molecules are there in 2.3 moles of NH4SO2?

10) How many moles are there in 3.3 x 1029 molecules of N2I6?

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Practice: By Mass (Gram formula/Molar mass) Use your periodic table to determine the Gram formula or Molar mass for each compound: The gram formula or molar mass of a compound is simply the sum of the masses of all the atoms in the formula.

1) NaBr (sodium bromide)

2) HNO3 (nitric acid)

3) H2O (Hydrogen hydroxide)

4) NaOH (sodium hydroxide)

5) CaCO3 (calcium carbonate)

6) LiI (lithium iodide)

7) AgCl (silver chloride)

8) C6H12O6 (glucose)

9) Ba(NO3)2(barium nitrate)

10) Fe2(SO4)3 (iron(III) sulfate)

11) Cl2

12) KOH

13) BeCl2

14) FeCl3

15) BF3

16) CCl2F2

17) Mg(OH)2

18) UF6

19) SO2

20) H3PO4

21) (NH4)2SO4

22) CH3COOH

23) Pb(NO3)2

24) Ga2(SO3)3

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Practice: By Mass (Mole ↔ Mass conversions) Use your periodic table to make the following conversions: 1) Find the mass of 0.89 mol of CaCI2. 2) A bottle of PbSO4 contains 158.1 g of the

compound. How many moles of PbSO4 are in the bottle?

3) Find the mass of 1.112 mol of HF. 4) Determine the number of moles of C5H12 that are

in 362.8 g of the compound. 5) Find the mass of 0.159 mol of Si02. 6) You are given 12.35 g of C4H802. How many moles

of the compound do you have? 7) Find the mass of 3.66 mol of N2. 8) A bottle of KMnO4 contains 66.38 g of the. How

many moles of KMnO4 does it contain?

9) Determine the number of atoms that are in 0.58 mol of Se.

10) How many moles of barium nitrate (BaNO3)

contain 6.80 x 1024 formula units? 11) Determine the number of atoms that are in 1.25

mol of 02. 12) How many moles of magnesium bromide (MgBr2)

contain 5.38 x 1024 formula units? 13) Determine the number of formula units that are

in 0.688 mol of AgNO3. 14) How many moles of ethane (C2H6) contain 8.46 x

1024 formula units? 15) Determine the number of formula units that are

in 1.48 mol of NaF. 16) How many formula units are in 3.5 g of

compound NaOH?

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Practice: By Volume (Mole ↔ Volume conversions) Use Avogadro’s number to make the following conversions:

1) What is STP? 2) A container with a volume of 893 L contains how

many moles of air at STP? 3) A chemical reaction produces 0.37 mol of N2 gas.

What volume will that gas occupy at STP? 4) A canister with a volume of 694 L contains how

many moles of oxygen at STP? 5) A chemical reaction produces 13.8 mol of CO gas.

What volume will that gas occupy at STP?

6) A tube with a volume of 3.68 L contains how

many moles of neon gas at STP? 7) A chemical reaction produces 0.884 mol of H2S

gas. What volume will that gas occupy at STP? 8) A container with a volume of 101 L contains how

many moles of argon gas at STP? 9) A chemical reaction produces 138 mol of HBr gas.

What volume will that gas occupy at STP?

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Topic 2: Chemical formulas (Ch. 10-3)

Day 6: Percent composition (10-3)

6) How do you calculate the percent by mass of an element in a compound?

Book, Page 306; 32 & 33

Book, Page 307; 34 & 35

Study Guide, page 9; 1-14 Day 7: Empirical formulas (10-3)

7) What does the empirical formula of a compound show?

Book, page 310; 36 & 37

Book, page 292; 5 & 6

Study Guide, page 10; 1-10 Day 8: Molecular formulas (10-3)

8) How does the molecular formula compare with the empirical formula?

Book, page 312; 38 & 39

Study Guide, page 11; 1-6

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Practice: Percent composition by mass Use your periodic table to make the following calculations: 1) H2O 2) H2SO4 3) KMnO4 4) C12H22O11 5) MgO 6) Find the percent composition of a compound that

contains 1.94 grams of carbon, 0.48 grams of hydrogen, and 2.58 grams of sulfur.

7) A sample of an unknown compound with a mass

of 0.847 grams has the following composition: 50.51% fluorine and 49.49 % iron. When this compound is decomposed into its elements, what mass of each element should be recovered?

8) Find the percent composition of a compound that

contains 2.63 grams of carbon, 0.370 grams of hydrogen, and 0.580 grams of oxygen.

9) A sample of an unknown compound with a mass of 2.876 grams has the following composition: 66.07% carbon, 6.71 % hydrogen, 4.06 % nitrogen and 23.16 % oxygen. What is the mass of each element in the compound?


contains 2.7369 grams of chlorine, 0.4116 grams of oxygen, and 3.9460 grams of phosphorus.


contains 1.51 grams of chromium, 1.13 grams of potassium, and 1.62 grams of oxygen.

12) A sample of a compound with a mass of 0.432

grams is analyzed. The sample is found to be made up of oxygen and fluorine only. Given that the sample contains 0.128 grams of oxygen, calculate the percent composition of the compound.

13) What is the percentage composition of a carbon-

oxygen compound, given that a sample contains 40.8 grams of carbon and 54.4 grams of oxygen?

14) What is the percentage composition of a sulfur-

chlorine compound, given that a sample contains 9.63 grams of sulfur and 21.3 grams of chlorine?

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Practice: Empirical Formula Use your periodic table to make the following calculations: 1) Determine the empirical formula of a

compound containing 2.644 g of gold and 0.476 g of chlorine.

2) Determine the empirical formula of a

compound containing 0.928 g of gallium and 0.412 g of phosphorus.


compound containing 1.723 g of carbon, 0.289 g of hydrogen, and 0.459 g of oxygen.

4) Find the empirical formula of a compound,

given that the compound is found to be 47.9 percent zinc and 52.1 percent chlorine by mass.

5) Find the empirical formula of a compound,

given that a 4&5-g sample of the compound contains 1.75 g of carbon and 46.75 g of bromine.

6) Determine the empirical formula of a compound containing 20.23 percent aluminum and 79.77 percent chlorine.


compound containing 24.74 percent potassium, 34.76 percent manganese, and 40.50 percent oxygen.


compound containing 4.288 g of carbon and 5.712 g of oxygen.


compound containing 2.16 g of aluminum, 3.85 g of sulfur and 7.68 g of oxygen.


compound containing 3.611 g of calcium and 6.389 g of chlorine.

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Practice: Molecular Formulas Use your periodic table to make the following calculations:

1) Find the molecular formula of a compound

that contains 42.56 g of palladium and O.SO g

of hydrogen. The molar mass of the

compound is 216.8 g/mol.

2) Octane, a compound of hydrogen and carbon,

has a molar mass of 114.26 g/mol. If one

mole of the compound contains 18.17 g of

hydrogen, what is its molecular formula?


that contains 30.45 percent nitrogen and

69.55 percent oxygen. The molar mass of the


4) Find the molecular formula of a compound,

given that a 212.1-g sample of the compound

contains 42.4 g of hydrogen and 169.7 g of

carbon and the molar mass is 30.0 g/mo!.

5) A compound is known to have a molar mass

of 391.5 g/mol. Find the molecular formula of

the compound, given the results of an

analysis of a 310.8-g sample that revealed

that the sample contains only boron and

iodine. The mass of the iodine in the sample

is found to be 302.2 g.


that contains 56.36 g of oxygen and 43.64 g

of phosphorus. The molar mass of the


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Review

Chemists relate units of counting (# atoms), of mass (grams), and of volume (liters) to a single quantity called

the _________________________. The number of representative particles in a mole of a substance is

_________________________. To find the mass of a mole of a compound, scientists add together the

_________________________ of the atoms making up the compound. When you substitute the unit grams for AMU,

you obtain the _________________________ of the compound. There are _________________________

representative particles in a mole of any substance.

At STP (0°C and 1 atmosphere pressure), one mole of any gas occupies a volume of

_________________________ L. This quantity is known as the _________________________ of the gas. To

determine the volume in liters of 2.00 mol of SO2 gas at STP, you would use _________________________ as a

conversion factor. _________________________ expressed in the units g/L, is used as a conversion factor when

converting from _________________________ to molar mass. When converting between numbers of representative

particles, _________________________, and volumes, you must always convert to _________________________ as

an intermediate step.

The _________________________ of a compound is the percent by mass of each element in a compound. The

percent by mass of an element in a compound is the number of grams of the element per

_________________________ g of the compound, multiplied by 100%. To calculate the percent by mass of an

element in a known compound, divide the mass of the element in one mole by the _________________________ and

multiply by 100%. A(n) _________________________ formula represents the lowest _________________________

ratio of the elements in a compound. It can be calculated from a compound’s percent composition. The

_________________________ formula of a compound is either the same as its empirical formula, or it is some whole-

number multiple of it.

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