Chemistry HP Unit 5 Stoichiometry Learning Targets …5+Packet+Honors+… · 1 Chemistry HP Unit 5 – Stoichiometry Learning Targets (Your exam at the end of Unit 5 will assess the

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Chemistry HP Unit 5 – Stoichiometry

Learning Targets (Your exam at the end of Unit 5 will assess the following:)

5. Stoichiometry

5-1. Define the quantity one mole by stating Avogadro’s number.

5-2. Determine the number of atoms of each element in a given compound from the chemical formula.

5-3. Define molar mass.

5-4. Determine the molar mass for a given element/compound with the appropriate units.

5-5. Determine the percent composition for a compound from the chemical formula.

5-6. Determine the percent composition of water in a hydrate.

5-7. Determine the empirical formula and molecular formula for a compound from percent composition or mass

information.

5-8. Perform conversions between moles, mass, and atoms/molecules and solve problems involving these

quantities, giving answers with the appropriate units.

5-9. Name and write formulas for hydrates.

5-10. Define stoichiometry.

5-11. Determine mole ratios for a reaction from the balanced chemical equation in order to convert between the

moles different substances.

5-12. Perform stoichiometric calculations involving mass of a reactant or product, giving answers with the

appropriate units.

5-13. Perform stoichiometric calculations involving limiting and excess reactants.

5-14. Determine the mass of the excess reactant remaining upon completion of the reaction.

5-15. Determine percent yield of a reaction given the actual yield and actual yield given percent yield.

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5-1. Define the quantity one mole by stating Avogadro’s number.

5-2. Determine the number of atoms of each element in a given compound from the chemical formula.

5-3. Define molar mass.

5-4. Determine the molar mass for a given element/compound with the appropriate units.

What is a “mole?”

The “mole” is an SI (System International) unit of measurement based on carbon-12. It is also known as Avogadro’s number. A mole is the number of atoms in 12 g of carbon-12 which is 6.02 × 1023 atoms. How much does a mole weigh?

How much a mole weighs depends on what it’s made of. A mole of carbon, for instance, or 6.02 x 1023 atoms of carbon,

weighs 12 grams. Moles of different atoms weigh differently, because different atoms weigh differently. We refer to

the mass of 1 mole of an atom or molecule as its molar mass. The mass of 1 mole of any atom on the Periodic Table is

equal to its atomic mass in grams.

Molar Mass = mass of 1 mole

Units: grams per mole, or g/mol

Sample Problem 1: What is the molar mass of water?

Now you try.

Practice Problem 2: Calculate the molar mass of nitric acid, HNO3.

63.0 g/mol

Solving Mole Problems: Converting Between Grams and Moles

To solve mole problems, we need to find out three things:

1. What the problem wants us to find.

2. What's given to us by the problem.

3. The conversion factors, which we'll talk more about as we go along.

The setup will take the form:

? Wanted = Given x ------------------

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5-8. Perform conversions between moles, mass, and atoms/molecules and solve problems involving these quantities,

giving answers with the appropriate units.

Sample Problem 3: How many moles of water, H2O, are in 6.83 g of water, H2O?

Sample Problem 4: How many grams are equivalent to 21.3 moles of BaCO3?

Now you try.

Practice Problem 5: Determine the number of moles of CO2 in 454 g.

10.3 mol

Practice Problem 6: Find the number of grams in 0.760 mole of H2SO4.

74.6 grams H2SO4

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Converting Between Molecules and Moles

Sample Problem 7: How many moles of H2O is 4.55 x 1023 molecules of H2O?

Sample Problem 8: How many formula units are in 0.400 mol KCl?

Sample Problem 9: How many ions are in 0.400 mol KCl?

Now you try.

Practice Problem 10: How many atoms are in 3.00 mol Sn?

1.81 x 1024 atoms Sn

Practice Problem 11: How many moles is 2.80 x 1024 atoms of silicon?

4.65 mol

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WS #1 (Learning Targets 5.4 and 5.8)

Show all calculations in the space below. Do not round answers.

(1) Molar mass is defined as the _______ of one _______ of an element or compound.

(2) The unit for molar mass is _______ per _______, written as ___/_____.

(3) Determine the molar mass of the following elements:

(a) iron (b) W

(c) carbon (d) Zn

(e) lead (f) Sb

(4) Calculate the molar mass of the following compounds:

(a) oxygen (b) hydrogen

(c) iodine (d) chlorine

(5) Calculate the molar mass of the following compounds. Name each compound.

Name of Compound Molar Mass

(a) AgCl

(b) Ba3(PO4)2

(c) Al(OH)3

(d) Cr(C2H3O2)2

(e) N2O4

(f) H2SO4

(6) Write the chemical formula for each compound. Calculate the molar mass.

Chemical Formula Molar Mass

(a) magnesium fluoride

(b) calcium chlorate

(c) tin (IV) bromide

(d) iron (III) nitrate

(e) phosphorus trichloride

(f) nitrous acid

Show all calculations required to solve each problem. Give answers with the appropriate units and significant figures.

(7) 1.50 mol Be _____________ Be atoms?

(8) 1.60 mol Cr = _____________ g Cr?

(9) 135 g F2O____________ mol F2O?

(10) 3.01x1023 molecules NaCl = _____________ mol NaCl?

(11) What is the mass of 0.800 mol of CuCl?

(12) How many moles of magnesium are in 1.806 x 1023 atoms?

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Answers:

(1) Molar mass is defined as the mass of one mole of an element or compound (2) The unit for molar mass is grams per mole,

written as g/mol.

(3) (a) 55.85 g/mol (b) 183.9 g/mol (c) 12.01 g/mol (d) 65.38 g/mol (e) 207.2 g/mol (f) 121.8 g/mol

(4) (a) 32.00 g/mol (b) 2.016 g/mol (c) 253.8 g/mol (d) 70.90 g/mol

(5) (a) 143.35 g/mol, silver chloride (b) 601.84 g/mol, barium phosphate (c) 78.004g/mol, aluminum hydroxide (d) 170.088 g/mol,

chromium (II) acetate (e) 92.02 g/mol dinitrogen tetroxide (f) 98.076 g/mol, sulfuric acid

(6) (a) MgF2, 62.31 g/mol (b) Ca(ClO3)2, 206.98 g/mol (c) SnBr4, 438.34 g/mol (d) Fe(NO3)3, 241.88 g/mol (e) PCl3, 137.22 g/mol (f)

HNO2, 47.018 g/mol

(7) 9.03 x 1023 atoms Be (8) 83.2 g Cr (9) 2.50 mol F2O (10) 0.500 mol NaCl (11) 79.2 g CuCl (12) 0.3000 mol Mg

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Multi-Step Molar Conversions

Sample Problem 12: How many atoms are in 36 g of bromine?

Sample Problem 13: What is the mass in grams of 1.20 x 1025 atoms of sulfur?

Sample Problem 14: How many formula units are in 22.4 g SnO2?

Now you try.

Practice Problem 15: What is the mass in grams of 25 x 1025 molecules of CO2?

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WS #2 (Learning Target 5.8: Perform conversions between moles, mass, and atoms/molecules and solve problems

involving these quantities, giving answers with the appropriate units.)

Show all calculations required to solve each problem. Give answers with the appropriate units and significant figures.

(1) What is the mass of 0.350 mol of KIO3?

(2) How many moles of radon are there in 1.11 g?

(3) What is the mass of 2.408 x 1024 atoms of iron?

(4) How many atoms are there in 157.6 g of gold?

(5) Diantimony trioxide is commonly used as a flame retardant. Give the formula for this compound. How many

molecules are contained in 72.90 g of diantimony trioxide? How many atoms are there of each element?

(6) How many molecules are in 0.4000 mol of oxygen? How many atoms of oxygen are there?

(7) Sodium nitrate is used in fertilizers, explosives, and glass manufacturing. Give the formula for this compound. How

many moles are in 1.7 g of sodium nitrate?

(8) What is the mass of 6.00 mol of carbon dioxide?

(9) Iron (III) chloride is used in the production of circuit boards. Give the formula for this compound. How many moles

are in 73.0 g of iron (III) chloride?

(10) Potassium sulfate is often used in fertilizers. Give the formula for this compound. How many molecules are

contained in 35 g of potassium sulfate? How many atoms are there of each element?

(11) Phosphoric acid is commonly added to beverages to give a sour taste. Give the formula for this compound. Calculate

the mass of 3.3 x 1023 molecules of phosphoric acid.

(12) Calcium thiosulfate can be used in water purification. How many molecules of calcium thiosulfate are in 0.761 kg?

How many atoms of each element are present? How many atoms are there in total?

(13) What is the mass of 5.84 x 1023 molecules of sucrose (C12H22O11), the chemical commonly known as table sugar?

(14) How many molecules are contained in 500.0 mg of cholesterol (C27H46O)?

Answers.

(1) 74.9 g KIO3 (2) 0.00500 mol radon (3) 223.4 g iron (4) 4.816 x 1023 atoms Au (5) 1.505 x 1023 molecules Sb2O3, 3.010 x 1023 atoms

Sb, 4.515 x 1023 atoms O, 7.525 x 1023 atoms total (6) 2.408 x 1023 molecules O2, 4.816 x 1023 atoms O (7) 0.020 mol NaNO3 (8) 264 g

CO2 (9) 0.450 mol FeCl3 (10) 1.2 x 1023 molecules K2SO4, 2.4 x 1023 atoms K, 1.2 x 1023 atoms S, 4.8 x 1023 atoms O (11) 54 g H3PO4

(12) 3.01 x 1024 molecules CaS2O3, 3.01 x 1024 atoms Ca, 6.02 x 1024 atoms S, 9.03 x 1024 atoms O, and 1.806 x 1025 atoms total (13)

332 g C12H22O11 (14) 7.785 x 1020 molecules C27H46O

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5-5. Determine the percent composition for a compound from the chemical formula.

Percent Composition

Percent Composition is defined as the percent by mass of each different element present in a compound. These are the steps to follow in order to calculate it.

1. Calculate the total mass of each element present in the formula of the compound. 2. Calculate the molecular weight of the compound. 3. Divide the total mass of each element by the molecular mass, and multiply by 100.

Sample Problem 16: Determine the percent composition of silver and chloride in the compound silver chloride, AgCl.

Sample Problem 17: Determine the percent composition of each element in ammonium nitrate, NH4NO3.

Now you try.

Practice Problem 18: Determine the percent composition of each element in glucose, C6H12O6.

39.99% carbon, 6.73% hydrogen, and 53.28% oxygen

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Percent Water in a Hydrate

Hydrates are compounds that contain water. Their chemical formulas have suffixes that take the format: dot nH2O, where n is an integer, corresponding to the number of water molecules attached to the compound. For Example:

Ba(OH)2·8H2O Barium hydroxide is bound to 8 water molecules, to form barium hydroxide octahydrate. MgCO3·5H2O Magnesium carbonate is bound to 5 water molecules, to form magnesium carbonate pentahydrate. Naming Hydrates

In order to name hydrates, follow these steps:

1. Name the ionic compound 2. Follow with the number of hydrates, preceded by the appropriate prefix:

Number of water molecules

Prefix Number of water molecules

Prefix

1 Mono- 7 Hepta-

2 Di- 8 Octa-

3 Tri- 9 Nona-

4 Tetra- 10 Deca-

5 Penta- 11 Undeca-

6 Hexa- 12 Dodeca-

Ba(OH)2•8H2O Barium hydroxide octahydrate. LiClO4•3H2O Lithium perchlorate trihydrate Now you try:

Practice Problem 19: Name the following compounds: MgCO3•5H2O Hg2(NO3)2•2H2O FeCl3•6H2O

Magnesium carbonate pentahydrate

Mercury(I) nitrate dihydrate

Iron(III) chloride hexahydrate

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Molar Mass of Hydrates

To calculate the molar mass of hydrates, add the molar mass of each element multiplied by the number of times that element appears in the formula (including the hydrate).

Sample Problem 20: Calculate the molar mass of barium hydroxide octahydrate, Ba(OH)2•8H2O

Sample Problem 21: What is the percent of water in Ba(OH)2•8H2O?

Sample Problem 22: When a 1.000 g sample of CuSO4•5H2O(s) was heated so that the waters of hydration

were driven off, the mass of the anhydrous salt remaining was found to be 0.6390 g. What is the experimental

value of the percent water of hydration?

Now You Try:

Practice Problem 23: A hydrate of Na2CO3 has a mass of 4.31 g before heating. After heating, its mass is found to be 3.22 g. Determine the percent by mass of water in the hydrate. 25.3%

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WS #3 (Learning Targets 5.6 and 5.7)

I. Hydrates

(1) Name the hydrate.

(a) Cs2CO2·2H2O ______________________ (c) MgSO4·7H2O ______________________

(b) FeCl3·6H2O ______________________ (d) Ca(NO3)2·4H2O ______________________

(2) Write the formula for each hydrate.

(a) calcium chloride dihydrate _______________ (c) magnesium hydrogen phosphate trihydrate _________________

(b) cobalt (II) chloride hexahydate ____________ (d) aluminum hypochlorite octahydrate

II. Percent Composition

Give answers to one decimal place

(1) Determine the percent composition for each element in the following compounds.

(a) SnCl2 (d) copper (II) sulfate

(b) CaCO3 (e) propane (C3H8)

(c) Mg3(PO4)2 (f) iron (III) nitrate hexahydrate

(2) Penicillin was one of the first antibiotics used to treat bacterial diseases. It was discovered by accident by Alexander

Fleming in 1928. The formula for penicillin is C14H20N2SO4. Determine the percent composition for each element.

(3) Gypsum (CaSO4·2H2O) is used in combination with limestone, sand, shale, and clay to make cement.

(a) What is the chemical name for gypsum?

(b) Determine the percent composition for each element in gypsum.

(c) Determine the percent composition of water in gypsum.

Answers:

I. Hydrates (1)(a) cesium carbonite dihydrate (b) iron (III) chloride hexahydrate (c) magnesium sulfate heptahydrate (d) calcium

nitrate tetrahydrate (2) (a) CaCl2·2H2O (b) CoCl2·6H2O (c) MgHPO4·3H2O (d) Al(ClO)3·8H2O

II. Percent Composition (1) (a) 62.6% Sn, 37.4% Cl (b) 40.0% Ca, 12.0% C, 48.0% O (c) 27.7% Mg, 23.6 % O, 48.7% O (d) 39.8% Cu,

20.1% S, 40.1% O (e) 81.7%C, 18.3% H (f) 16.0% Fe, 12.0% N, 68.6% O, 3.5% H (2) 53.8% C, 6.5% H, 9.0% N, 10.3% S, 20.5% O (3) (a)

calcium sulfate dihydrate (b) 23.3% Ca, 18.6% S, 55.8% O, 2.3% H (c) 20.9% H2O

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information.

Empirical Formulas

An empirical formula is a formula that gives the simplest whole-number ratio of atoms in a compound. For example, C6H12O6 is the molecular formula of glucose. But, if we divided each of these subscripts by the greatest common factor, which in this case is 6, we could reduce this formula to CH2O. CH2O would be the empirical formula: the simplest whole-number ratio of atoms in the compound.

Sample Problem 24: Convert H2O2 to its empirical formula.

Now you try.

Practice Problem 25: Convert C6H9 to its empirical formula.

C2H3

How to Calculate Empirical Formulas

Use the following steps:

1. Percent to Mass 2. Mass to Moles 3. Divide by Small 4. Times to Whole How to Determine Molecular Formulas

The molecular formula is either the same or a whole number multiple of the empirical formula. For example, the empirical formula of glucose is CH2O. Its molecular formula is C6H12O6. C6H12O6 is a whole number multiple of CH2O. 1. Determine the empirical formula.

2. Calculate the molar mass of the empirical formula. 3. Divide the molar mass of the molecular formula by the molar mass of the empirical formula to get the multiple. 4. Multiply all subscripts in the empirical formula by that multiple.

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information.

Sample Problem 26: A compound is composed of 85.7 % carbon and 14.3 % hydrogen. Find the empirical

formula for this compound.

Sample Problem 27: NutraSweet is 57.14 % C, 6.16 % H, 9.52 % N, and 27.18 % O. Calculate the empirical

formula of NutraSweet. Calculate its molecular formula, given that its molar mass is 294.30 g/mol.

Now you try.

Practice Problem 28: A compound is composed of 53.30% carbon 11.19% hydrogen and 35.51% oxygen by

mass. Calculate the empirical formula of the compound. If its molar mass is 90.12 g/mol, what is the molecular

formula for the compound?

C2H5O, C4H10O2

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WS #4 (Learning Target 5.7: Determine the empirical formula and molecular formula for a compound from percent

composition or mass information.)

I. Empirical Formulas

(1) Determine the empirical formula for the following compounds. Name each compound.

(a) 77.7% Fe, 22.3% O

(b) 70.0% Fe, 30.0% O

(c) 64.0% K, 9.8% C, 26.2% O

(d) 56.6% K, 8.7% C, 34.7% O

(2) Methylamine has the following composition: 38.7% carbon, 16.2 % hydrogen, and the remainder nitrogen.

Determine the empirical formula.

(3) A compound contains only calcium, carbon, and oxygen. A 34.5 g sample is analyzed to show that there is 13.8 g of

calcium, 4.2 g of carbon, and 16.5 g of oxygen. Determine the empirical formula. Name the compound.

(4) The compound called acetone contains only carbon, hydrogen, and oxygen. A 6.50 g sample of the compound is

analyzed to show that there is 3.54 g of carbon and 0.59 g of hydrogen. The remainder is oxygen. Determine the

empirical formula.

II. Molecular Formulas

(1) What is the difference between an empirical formula and a molecular formula?

(2) The molecular formulas of some substances are as follows. Write the empirical formulas.

(a) acetylene: C2H2 (used in welding torches)

(b) glucose: C6H12O6 (the simplest form of sugar)

(c) octane: C8H18 (a component of gasoline)

(3) Styrene is a chemical used in the manufacture of hard plastics. Styrene contains 92.3% carbon and 7.7% hydrogen.

The molecular weight of styrene is 104.144 g/mol. Determine the empirical and molecular formula.

(4) Potassium persulfate is used in the process of film developing. A 0.8162 g sample was found to contain 0.2361 g of

potassium, 0.1936 g of sulfur, and 0.3865 of oxygen. The molecular weight of potassium persulfate is 675.8 g/mol.

Determine the empirical and molecular formula.

(5) Isobutylene is a raw material for making synthetic rubber. A sample with a mass of 0.6481 g was found to contain

0.5555 g of carbon and the rest was hydrogen. The molecular weight of isobutylene is 56.104 g/mol. Determine the

empirical and molecular formula.

(6) A 44.860 g sample of white powder was found to contain 19.58 g of phosphorus and 25.28 g of oxygen. The

molecular weight of the compound is 283.88 g/mol. Determine the empirical and molecular formula. Name the

molecular compound.

Answers: I. Empirical Formulas (1)(a) FeO iron (II) oxide (b) Fe2O3 iron (III) oxide (c) K2CO2 potassium carbonite (d) K2CO3 potassium carbonate (2) CH5N (3) CaCO3 calcium carbonate (4) C2H4O II. Molecular Formulas (1) empirical formula- the simplest formula molecular formula- the formula of the naturally occurring compound (2) (a) CH (b) CH2O (c) C4H9 (3) CH, C8H8 (4) KSO4, K5S5O20 (5) CH2, C4H8 (6) P2O5, P4O10

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5-11. Determine mole ratios for a reaction from the balanced chemical equation in order to convert between the moles

different substances.

Mole Ratios

The single most important thing in Stoichiometry you need to be able to understand is what the coefficients in a balanced chemical equation really represent.

N2(g) + 3H2(g) 2NH3(g)

N2(g) 3H2(g) 2NH3(g)

1 molecule N2 3 molecules H2 2 molecules NH3

2 molecules N2 6 molecules H2 4 molecules NH3

10 molecules N2 30 molecules H2 20 molecules NH3

6.02 x 1023 molecules

N2

3 x (6.02 x 1023) molecules N2 2 x (6.02 x 1023) molecules NH3

1 mole N2 3 moles H2 2 moles NH3

The coefficients in a balanced equation can be used to write mole ratios.

Mole ratios are conversion factors that can be used to relate the moles of any reactants to the moles of any other reactants and products.

We write mole ratios as follows:

1 mole N2 = 3 moles H2

1 mole N2 = 2 moles NH3

3 moles of H2 = 2 moles of NH3

Sample Problem 29: How many moles of H2O can be produced from 4 moles of NH3? 4 NH3(g) + 5 O2(g) 4 NO(g) + 6 H2O(g)

Now you try.

Practice Problem 30: How many moles of O2 do you need to make 9 moles of H2O? 2 H2(g) + O2(g) 2 H2O(g)

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5-10. Define stoichiometry.

5-12. Perform stoichiometric calculations involving mass of a reactant or product, giving answers with the appropriate

units.

Stoichiometry

“Stoich” = ______________________

“metry” = ______________________

Stoichiometry = __________________________________________________________________________

Sample Problem 31: How many grams of ammonia are produced from 2.00 moles of nitrogen? N2 + 3H2 2NH3

Sample Problem 32: How many grams of ammonia are produced from 5.00 grams of hydrogen? N2 + 3H2 2NH3

Now you try.

Practice Problem 33: How many grams of water will be produced from 50.0 g of NH3? 4 NH3 + 3 O2 2 N2 + 6H2O

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WS #5 (Learning Targets 5-10, 5-11, 5-12)

I. Mole Ratios

(1) AlCl3 + K3PO4 → AlPO4 + 3KCl

(a) If 10.2 mol of AlCl3 reacts, determine the moles of KCl produced

(b) If 0.40 mol of AlCl3 reacts, determine the moles of K3PO4 reacting

(c) If 4.3 mol of AlPO4 is produced, determine the moles of AlCl3 reacting

(d) If 2.7 mol of KCl is produced, determine the moles of AlPO4 produced

(2) 2Fe2O3 + 3C → 4Fe + 3CO2

(a) If 2.4 mol of C reacts, determine the moles of Fe2O3 reacting

(b) If 0.04 mol of Fe2O3 reacts, determine the moles of Fe produced

(c) If 0.300 mol of Fe is produced, determine the moles of CO2 produced

(d) If 5.7 mol of CO2 is produced, determine the moles of C reacting

(3) 2C4H10 + 13O2 → 8 CO2 + 10H2O

(a) If 2.6 mol of O2 reacts, determine the moles of C4H10 reacting

(b) If 0.030 mol of C4H10 reacts, determine the moles of CO2 produced

(c) If 8.00 mol of H2O is produced, determine the moles of O2 reacting

(d) If 12.4 mol of CO2 is produced, determine the moles of H2O produced

II. Stoichiometry

(4) Sodium is reacted with oxygen to form sodium oxide according to the following balanced equation:

4Na + O2 → 2Na2O

What mass of oxygen is required to react with 92.0 g of sodium? What mass of sodium oxide is produced?

(5) Silver nitrate is reacted with zinc to produce silver and zinc nitrate according to the following balanced equation:

2AgNO3 + Zn → 2Ag + Zn(NO3)2.

If 340 g of silver nitrate are present, what mass of zinc would be required? What would be the mass of each of the

products?

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(6) Magnesium chloride can be decomposed into magnesium and chlorine. Write a balanced chemical equation for this

reaction. What mass of magnesium chloride would be required to produce 36.5 g of magnesium? What mass of chlorine

would be produced?

(7) Chlorine reacts with sodium bromide to produce sodium chloride and bromine. Write a balanced chemical equation

for this reaction. What mass of sodium bromide is required to react with 14.18 g of chlorine? What would be the mass of

each of the products?

(8) Ammonium sulfate is reacted with sodium nitrate. Write a balanced chemical equation for this reaction.

(a) If 2.64 g of ammonium sulfate are present, what mass of sodium nitrate would be required in the reaction?

What would be the mass of each of the products?

(b) If 170 g of sodium nitrate are present, what mass of ammonium sulfate would be required in the reaction?


(9) Aluminum chloride is reacted with lead (II) nitrate. Write a balanced chemical equation for this reaction.

(a) If 26.7 g of aluminum chloride are present, what mass of lead (II) nitrate would be required in the reaction?


(b) If 8.28 g of lead (II) nitrate are present, what mass of aluminum chloride would be required in the reaction?


Answers: (1) (a) 30.6 mol KCl (b) 0.40 mol K3PO4 (c) 4.3 mol AlCl3 (d) 0.90 mol AlPO4 (2) (a) 1.6 mol Fe2O3 (b) 0.08 mol Fe

(c) 0.225 mole CO2 (d) 5.7 mol C (3) (a) 0.40 mol C4H10 (b) 0.12 mol CO2 (c) 10.4 mol O2 (d) 15.5 mol H2O

(4) 32.0 g O2, 124 g Na2O (5) 65.4 g Zn, 216 g Ag, 189 g Zn(NO3)2 (6) MgCl2 → Mg + Cl2, 143 g MgCl2, 106 g Cl2

(7) Cl2 + 2NaBr → 2NaCl + Br2, 41.16 g NaBr, 23.38 g NaCl, 31.96 g Br2 (8) (NH4)2SO4 + 2NaNO3 → 2NH4NO3 + Na2SO4

(a) 3.40 g NaNO3, 3.20 g NH4NO3, 2.84 g Na2SO4 (b) 132 g (NH4)2SO4, 160 g NH4NO3, 142 g Na2SO4

(9) 2AlCl3 + 3Pb(NO3)2 → 2Al(NO3)3 + 3PbCl2 (a) 99.5 g Pb(NO3)2, 42.7 g Al(NO3)3, 83.5 g PbCl2 (b) 2.22 g AlCl3, 3.55 g

Al(NO3)3, 6.95 g PbCl2

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Percent Yield

In real life, reactions do not produce as much product as one would predict.

Therefore, some stoichiometry problems will ask you to calculate the percent yield of an actual reaction.

Percent yield is a percentage indicating the amount of product actually produced as compared to the amount of

product predicted.

IMPORTANT!

In Percent Yield problems, they always give you the amount of reactant AND the amount of product.

The WANTED will always be the amount of product.

The GIVEN will always be the amount of reactant.

To calculate percent yield, use the following steps:

1. Use stoichiometry to determine the theoretical yield of product.

a. The GIVEN will always be the amount of reactant.

b. The WANTED will always be the amount of product.

2. Divide the actual yield by this theoretical yield and multiply by 100.

a. The ACTUAL YIELD is the amount of product GIVEN TO US in the problem.

Sample Problem 37: What is the percent yield of H2O if 138 g H2O is produced from 16.0 g H2 and excess O2?

2 H2 + O2 --> 2 H2O

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Sample Problem 38: A student adds 200.0g of C7H6O3 to an excess of C4H6O3, this produces C9H8O4 and C2H4O2. Calculate the percent yield if 231 g of aspirin (C9H8O4) is produced.

C7H6O3 + C4H6O3 C9H8O4 + C2H4O2

Now you try.

Practice Problem 39: When 34.7 g of PCl3 in the laboratory is reacted with excess C2H5OH, 23.7 g of C2H5Cl is produced, according to the following equation. What is the percent yield of C2H5Cl?

3 C2H5OH + PCl3 3 C2H5Cl + H3PO3

48.6%

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WS #6 (Learning Target 5-15: Determine percent yield of a reaction given the actual yield and actual yield given

percent yield.)

Percent Yield

(1) Strontium nitride can be decomposed into strontium and nitrogen. Write a balanced chemical equation for this

reaction. If 58.2 g of strontium nitride yields 45.0 g of strontium, determine the percent yield for the reaction. What

mass of nitrogen will actually be obtained?

(2) Mercury (II) oxide can be decomposed into mercury and oxygen. Write a balanced chemical equation for this

reaction. If 54.15 g of mercury (II) oxide yields 3.250 g of oxygen, determine the percent yield for the reaction. What

mass of mercury will actually be obtained?

(3) Zinc and fluorine can be combined to form zinc fluoride. Write a balanced chemical equation for this reaction. (a) If

16.35 g of zinc yields 16.50 g of zinc fluoride, determine the percent yield of the reaction. (b) If 0.380 g of fluorine reacts

with a percent yield of 60.0%, what mass of zinc fluoride will actually be obtained?

(4) Magnesium chloride reacts with silver nitrate. Write a balanced chemical equation for this reaction. (a) If 21.2 g of

silver nitrate react with a percent yield of 80.0%, what mass would actually be obtained for each of the products. (b) If

476 g of magnesium chloride yields 400 g of magnesium nitrate, determine the percent yield of the reaction. What mass

of silver chloride will actually be obtained?

(5) Aluminum is reacted with copper (II) sulfate. Write a balanced chemical equation for this reaction. If 5.60 g of

aluminum react, determine the mass of copper (II) sulfate required. If 18.0 g of copper are actually produced, determine

the percent yield of the reaction. What mass of aluminum sulfate will actually be obtained?

Answers:

(1) Sr3N2 → 3Sr + N2, 85.6%, actual mass N2: 4.80g

(2) 2HgO → 2Hg + O2, 81.25%, actual mass Hg: 40.75 g

(3) Zn + F2 → ZnF2 (a) 63.82% (b) actual mass ZnF2: 0.620 g

(4) MgCl2 + 2AgNO3 → Mg(NO3)2 + 2AgCl (a) actual mass Mg(NO3)2: 7.40 g, actual mass AgCl: 14.3 g

(b) 53.9%, actual mass AgCl: 771 g

(5) 2Al + 3CuSO4 → Al2(SO4)3 + 3Cu 49.7 g CuSO4, 90.9 %, actual mass Al2(SO4)3: 32.3 g

23



Limiting Reagents

The Limiting Reagent is the reactant that is totally consumed in a chemical reaction. It’s what we don’t have enough of. Once all of the limiting reagent is used up, the reaction is finished. How do you know you have a limiting reagent problem? You’re given the mass of more than one reactant. How to solve any limiting reagent problem:

1. Realize you have a limiting reagent (LR) problem.

When you are given the quantities of more than one species in a reaction, chances are you have a limiting reagent (LR) problem. 2. Convert the given quantities to moles. 3. Divide each mole quantity by its coefficient. 4. The smallest result is the limiting reagent. 5. Use this limiting reagent to solve the rest of the problem.

Sample Problem 34: A 50.6 g sample of Mg(OH)2 is reacted with 45.0 g of HCl according to the following

reaction. How many grams of MgCl2 are produced? How much of the excess reactant is left over?

Mg(OH)2 + 2 HCl MgCl2 + 2 H2O

24



Sample Problem 35: A 20.1 g sample of Pb(NO3)2 is reacted with 25.0 g of KI according to the following reaction. How many grams of lead iodide is precipitated? How many grams of excess reagent remain?

Pb(NO3)2 + 2KI PbI2 + 2KNO3

Now you try.

Practice Problem 36: A 100. g sample of CaCO3 is reacted with 50.0 g of FePO4 according to the following reaction. Determine the amount of products produced. Determine the mass of excess reagent remaining.

3 CaCO3 + 2 FePO4 Ca3(PO4)2 + Fe2(CO3)3

40.9 g Ca3(PO4)2, 48.4 g Fe2(CO3)3, 50.2 g CaCO3

25

WS #7 (Learning Targets 5-13 and 5-14)

Limiting and Excess Reactions

(1) Copper (II) sulfate is reacted with zinc to produce copper and zinc sulfate. Write a balanced chemical equation for

this reaction. If 2.00 mol of copper (II) sulfate are combined with 3.00 mol of zinc, which reactant is limiting and which is

excess? What is the mass of each of the products?

(2) Calcium hydroxide reacts with phosphoric acid to produce water and calcium phosphate. Write a balanced chemical

equation for this reaction. If 4.20 mol of calcium hydroxide are reacted with 2.20 mol of phosphoric acid, which reactant

is limiting and which is excess. What is the mass of each of the products?

(3) Lithium is reacted with iodine. Write a balanced chemical equation for this reaction. If 15.10 g of lithium are

combined with 200.4 g of iodine, which reactant is limiting and which is excess? What is the mass of the product? What

mass of the excess reactant is used in the reaction, and what mass of the excess reactant remains after the reaction?

(4) Barium carbonate is reacted with ammonium fluoride. Write a balanced chemical equation for this reaction. If 240 g

of barium carbonate and 95.0 g ammonium fluoride are combined, which reactant is limiting and which is excess? What

is the mass of each of the products? If the reaction yields 100 g of ammonium carbonate, determine the percent yield.

What mass of barium fluoride would actually be obtained? What mass of the excess reactant is used in the reaction, and

what mass of the excess reactant remains after the reaction?

(5) Aluminum bromide is reacted with chlorine. Write a balanced chemical equation for this reaction. If 135 g of

aluminum bromide and 50.0 g chlorine are combined, which reactant is limiting and which is excess? What is the mass of

each of the products? If the reaction yields 104 g of bromine, determine the percent yield. What mass of aluminum

chloride would actually be obtained? What mass of the excess reactant is used in the reaction, and what mass of the

excess reactant remains after the reaction?

Answers: (1) CuSO4 + Zn → Cu + ZnSO4 lim: CuSO4 ex: Zn, 127 g Cu, 323 g ZnSO4 (2) 3Ca(OH)2 + 2H3PO4 → 6H2O + Ca3(PO4)2 lim: H3PO4

ex: Ca(OH)2, 119 g H2O, 341 g Ca3(PO4)2 (3) 2Li + I2 → 2LiI lim: I2 ex: Li, 211.4 g LiI used: 10.96 g Li, remaining: 4.14 g Li (4) BaCO3 +

2NH4F→ BaF2 + (NH4)2CO3 lim: BaCO3 ex: NH4F, 213 g BaF2, 117 g (NH4)2CO3 85.5%, actual mass BaF2: 182 g used: 90.1 g NH4F,

remaining: 4.9 g NH4F (5) 2AlBr3 + 3Cl2 → 2AlCl3 + 3Br2 lim: Cl2 ex: AlBr3, 62.7 g AlCl3, 113 g Br2 92.0%, actual mass AlCl3: 57.7g used:

125 g AlBr3, remaining: 10g AlBr3

26

WS #8 (Unit 5 Review)

I. Reactions

Classify the reaction, predict the products, and balance the chemical equation.

____ (a) aluminum combines with nitrogen

____ (b) chloric acid reacts with barium hydroxide

____ (c) bromine reacts with potassium iodide

____ (d) aluminum chloride reacts with sodium phosphate

II. Mole Ratios

(1) C5H12 + 8O2 → 6H2O + 5CO2

(a) If 0.20 mol of C5H12 reacts, determine the moles of O2 reacting

(b) If 4.8 mol of O2 reacts, determine the moles of H2O produced

(c) If 30 mol of H2O is produced, determine the moles of CO2 produced

(d) If 0.015 mol of CO2 is produced, determine the moles of C5H12 reacting

(2) Calcium iodide is reacted with sodium. Write a balanced chemical equation for this reaction.

(a) If 2.2 mol of calcium iodide reacts, determine the moles of sodium reacting and the moles of each product.

(b) If 0.016 mol of sodium reacts, determine the moles of calcium iodide reacting and the moles of each product.

III. Mass-Mass Stoichiometry

(3) Antimony and iodine can be combined to form antimony triiodide. Write a balanced chemical equation for this

reaction. If 60.9 g of antimony are present, what mass of iodine will be required in the reaction? What mass of antimony

triiodide will be produce?

(4) Strontium chloride reacts with sodium oxalate to produce strontium oxalate and sodium chloride. Write a balanced

chemical equation for this reaction. If 3.20 g of strontium chloride are present, determine the mass of sodium oxalate

required in the reaction. What mass of strontium oxalate and sodium chloride will be produced?

27

WS #8, p. 2 (Unit 5 Review)

IV. Percent Yield

(5) Silver nitrate is reacted with sodium sulfate. Write a balanced chemical equation for this reaction.

(a) If 98.6 g of silver nitrate react with a percent yield of 75.5%, what mass would actually be obtained for each

of the products.

(b) If 30.0 g of sodium sulfate yields 25.0 g of sodium nitrate, determine the percent yield of the reaction. What

mass of silver sulfate will actually be obtained?

V. Limiting and Excess Reactants

(6) Iron (III) iodide reacts with bromine. Write a balanced chemical equation for this reaction.

If 218 g of iron (III) iodide reacts with 90.0 g of bromine, which reactant is limiting and which is excess?

What is the mass of each of the products?

What mass of the excess reactant is used in the reaction, and what mass of the excess reactant remains after the

reaction?

If the reaction yields 115 g of iodine, determine the percent yield. What mass of iron (III) bromide would actually be

obtained?

Answers:

I. Reactions

(a) S: 2Al + N2 → 2AlN (b) N: 2HClO3 + Ba(OH)2 → Ba(ClO3)2 + 2H2O (c) SR: Br2 + 2KI → 2KBr + I2

(d) DR: AlCl3 + Na3PO4 → AlPO4 + 3NaCl

II. Mole Ratios

(1) (a) 1.6 mol O2 reacting (b) 3.6 mol H2O produced (c) 25 mol CO2 produced (d) 0.0030 mol C5H12 reacting

(2) CaI2 + 2Na → Ca + 2NaI (a) 4.4 mol Na, 4.4 mol NaI, 2.2 mol Ca (b) 0.0080 mol CaI2, 0.0080 mol Ca, 0.016 mol NaI

III. Mass-Mass Stoichiometry

(3) 2Sb + 3I2 → 2SbI3 190 g I2, 251 g SbI3 (4) SrCl2 + Na2C2O4 → SrC2O4 + 2 NaCl 2.71 g Na2C2O4, 3.55 g SrC2O4, 2.36 g NaCl

IV. Percent Yield

(5) 2AgNO3 + Na2SO4 → Ag2SO4 + 2NaNO3 (a) 68.3 g Ag2SO4 , 37.2 g NaNO3 (b) 69.6%, 45.9 g Ag2SO4

V. Limiting and Excess Reactions

(6) 2FeI3 + 3Br2 → 2FeBr3 + 3I2 lim: Br2 ex: FeI3, 111 g FeBr3, 143 g I2 80.5%, actual mass: FeBr3 89.3 g used: 164 g FeI3, remaining 54 g

FeI3

Chemistry HP Unit 5 Stoichiometry Learning Targets …5+Packet+Honors+… · 1 Chemistry HP Unit 5 – Stoichiometry Learning Targets (Your exam at the end of Unit 5 will assess the

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