Stoichiometry Objectives1. Identify the quantitative relationships in a balanced chemical

chemical equation.2. Determine the mole ratios from from a balanced chemical

equation.3. Explain the sequence of steps used in solving stoichiometric

problems.4. Use the steps to solve stoichiometric problems.5. Identify the limiting reactant in a chemical equation.6. Identify the excess reactant and calculate the amount remaining

after the reaction is complete.7. Calculate the mass of a product when the amounts of more than

one reactant are given.8. Calculate the theoretical yield of a chemical reaction from data.9. Determine the percent yield for a chemical reaction.

The Mole (Ch 11)Chemists need a convenient method for counting accurately

the number of atoms, molecules, or formula units in a sample of a substance.

-atoms and molecules are extremely small

-even in the smallest sample it’s impossible to actually count each individual atom.

To fix this problem chemists created their own counting unit called the mole.

-mole: commonly abbreviated mol, is the SI base unit used to measure the amount of a substance

The Mole

Through experimentation, it has been established 1 mole = 6.022 136 7 x 1023 representative particles.

-representative particle: any particle such as atoms, molecules, formula units, electrons, or ions.

-called Avogadro’s number in honor of the Italian physicist and lawyer Amedeo Avogadro who, in 1811, determined the volume of one mole of a gas.

-we round Avogadro’s number to three significant figures— 6.02 x 1023.

- If you write out Avogadro’s number, it looks like this:

- 602 000 000 000 000 000 000 000

One Mole Quantities

Other Mole & Stoichiometry Vocabulary

1. What is stoichiometry?Study of quantitative relationships among amounts of reactants an products.

2. What is a mole ratio?Ratio between the moles of any two substances in a balanced chemical equation

3. What is molar mass? Mass, in grams, of one mole of pure substance.

-numerically equal to its atomic mass

-has the units g/mol

Converting Moles to Particles (11.1)

Determine how many particles of sucrose are in 3.50mol of sucrose.

- write a conversion factor using Avogadro’s number that relates representative particles to moles of a substance.

Converting Particles to Moles (11.1)

Now, suppose you want to find out how many moles are represented by a certain number of representative particles, such as 4.50 x 1024 atoms of zinc.

-You can use the inverse of Avogadro’s number as a conversion factor.

Mole Practice 1

Identify and calculate the number of representative particles in each of the following quantities.

1. 2.15 moles of gold

2. 0.151 mole of nitrogen oxide

3. 11.5 moles of potassium bromide

Calculate the number of moles of the substance that contains the following number of representative particles.

4. 8.92 x 1023 atoms of barium

5. 5.50 x 1023 molecules of carbon monoxide

6. 2.66 x 1023 formula units of potassium iodide

Practice-Homeworkp 311 # 1-3; p 312 # 4

Moles to Mass (11.2)

To convert between moles and mass, you need to use the atomic mass found on the periodic table.

Calculate the mass of 0.625 moles of calcium.

-According to the periodic table, the atomic mass of calcium is 40.078 amu, so the molar mass of calcium is 40.078 g/mol.

Mass to Moles (11.2)

How many moles of copper are in a roll of copper that has a mass of 848g?

Practice-Homeworkp 316 # 11ab-12ab

Mass to Atoms (11.2)

To find the number of atoms in a sample, you must first determine the number of moles.

Calculate the number of atoms in 4.77 g lead.

1. Determine moles

Mass to Atoms (cont.)

2. Determine atoms

You can also convert from number of particles to mass by reversing the procedure above and dividing the

number of particles by Avogadro’s number to determine the number of moles present.

Atoms to Mass

Example problem 11-5, p 318

Mole Practice 2

Determine the mass in grams of the following quantities. 1. 1.24 moles of beryllium 2. 3.35 moles of calciumCalculate the number of moles in each of the following.

3. 6.35 g lithium4. 346 g zinc

How many atoms are in the following samples?5. 1.24 g cobalt6. 0.575 g cesium

How many grams are in the following samples?7. 4.16 x1023 atoms of radium8. 1.50 x 1020 atoms of cadmium

Practice-Homworkp 316 # 11ab-12ab

p 318 # 13ab-14ab

Moles of Compounds (11.3)

A mole of a compound contains as many moles of each element as are indicated by the subscripts in the formula.

-For example, a mole of ammonia (NH3) consists of one mole of nitrogen atoms and three moles of hydrogen atoms.

-the molar mass of the compound is found by adding the molar masses of all of the atoms in the representative particle.

Molar mass of NH3 = 1(molar mass of N) + 3(molar mass of H)

Molar mass of NH3 = 1(14.007 g) + 3(1.008 g) = 17.031 g/mol

Practicep 322 # 25

Book Practice

P 312 # 4

-Mole relationships from a formula (p 321)

-Mole to Mass for compounds ( p 323)

Mass of Compound to Moles

Calculate the number of moles of water that are in 1.000 kg of water?

1. Before you can calculate moles, you must determine the molar mass of water (H2O).

molar mass H2O = 2(molar mass H) + molar mass O

2. Now you can use the molar mass of water as a conversion factor to determine moles of water.

-Notice 1.000 kg is converted to 1.000 x 103 g

Practicep 323 # 27-28

p 324 # 30a&b

Mole Practice 3

Calculate the molar mass of the following:

1. C2H5OH

2. HCN

What is the mass of the following:

3. 2.25 moles of KMnO4

4. 1.56 moles of H2O

Determine the number of moles in the following:

5. 35.0 g HCl

6. 254 g PbCl4

What is the mass in grams of one molecule of the following:

7. H2SO4

Percent Composition, Molecular & Empirical Formulas

Recall that every chemical compound has a definite composition—a composition that is always the same wherever that compound is found.

The composition of a compound is usually stated as the percent by mass of each element in the compound, using the following process.

Percent Composition

Example: Determine the percent composition of calcium chloride (CaCl2).

1. Determine mass of each ion in CaCl2.

-1mol CaCl2 consists of 1mol Ca+2 ions and 2mol Cl- ions.

1mol Ca+2 ions x 40.08g Ca+2 ions = 40.08g Ca+2 ions

1mol Ca+2 ions

2mol Cl- ions x 35.45g Cl- ions = 70.90g Cl- ions

1mol Cl- ions

Percent Composition

Example: Determine the percent composition of calcium chloride (CaCl2).

2. Calculate molar mass of CaCl2.

- 40.08g Ca+2 ions + 70.90g Cl- ions = 110.98 g CaCl2

1 mole CaCl2 1 mole CaCl2

3. Determine percent by mass of each element.

Percent Composition

Example: Determine the percent composition of calcium chloride (CaCl2).

3. Determine percent by mass of each element.

% Ca = 40.08 g Ca+2 x 100 = 36.11 % Ca+2

110.98 g CaCl2

% Cl = 70.90 g Cl- x 100 = 63.89 % Cl-

110.98 g CaCl2

4. Make sure your percent compositions equal 100%.

36.11% Ca+2 + 63.98% Cl- = 100%

Practice

p 331 # 42-43

Empirical Formula

You can use percent composition data to help identify an unknown compound by determining its empirical formula.

-empirical formula-simplest whole-number ratio of atoms of elements in the compound.

~In many cases, the empirical formula is the actual

formula for the compound.

the empirical formula of sodium chloride is Na1Cl1,

or NaCl, which is the true formula

~sometimes, the empirical formula is not the actual

formula of the compound.

the empirical formula for N2O4 (the actual) is NO2.

Empirical Formula

Example: The percent composition of an unknown compound is found to be 38.43% Mn, 16.80% C, and 44.77% O. Determine the compound’s empirical formula.

- Because percent means “parts per hundred parts,” assume that you have 100 g of the compound.

1. Calculate the number of moles of each element in the 100 g of compound.

Empirical Formula

Example: The percent composition of an unknown compound is found to be 38.43% Mn, 16.80% C, and 44.77% O. Determine the compound’s empirical formula.

1. Calculate the number of moles of each element in the 100 g of compound.

Empirical Formula

Example: The percent composition of an unknown compound is found to be 38.43% Mn, 16.80% C, and 44.77% O. Determine the compound’s empirical formula.

2. The results show the following relationship

3. Obtain the simplest whole-number ratio of moles:

-divide each number of moles by the smallest number of moles.

0.6995 mol Mn : 1.339 mol C : 2.798 mol O

0.6995 mol 0.6995 mol 0.6995 mol

1 : 2 : 4

Empirical Formula

Example: The percent composition of an unknown compound is found to be 38.43% Mn, 16.80% C, and 44.77% O. Determine the compound’s empirical formula.

3. Obtain the simplest whole-number ratio of moles:

Mn : C : O

1 : 2 : 4

4. Determine the empirical formula.

MnC2O4

Practice

p 333 # 46-47

Molecular Formula

For many compounds, the empirical formula is not the true formula.

-Chemists have learned, though, that acetic acid is a molecule with the formula C2H4O2, which is the molecular formula for acetic acid.

-molecular formula tells the exact number of atoms of each element in a molecule or formula unit of a compound.

Notice the molecular formula for acetic acid (C2H4O2) has exactly twice as many atoms of each element as the empirical formula (CH2O).

-The molecular formula is always a whole-number multiple of the empirical formula.

Molecular Formula

Example: Determine the molecular formula for maleic acid, which has a molar mass of 116.1g/mol.

1. empirical formula of the compound

-composition of maleic acid is 41.39% C, 3.47% H, and

55.14% O (change the % to g)

Molecular Formula

Example: Determine the molecular formula for maleic acid.

1. empirical formula of the compound

-the ratio of C:H:O is 1:1:1, making the empirical formula

CHO

2. calculate the molar mass of CHO (empirical formula).

-29.01g/mol

3. Determine the molecular formula for maleic acid,

Molecular Formula

Example: Determine the molecular formula for maleic acid.

3. Determine the molecular formula for maleic acid,

-shows the molar mass of maleic acid is 4x that of CHO. 4. Multiply CHO by 4 to get C4H4O4

Practice

p 335 # 51-52

Empirical Formula from Mass

You can also calculate the empirical formula of a compound from mass of individual elements.

Example: Determine the empirical formula for ilmenite, which contains 5.41g Fe, 4.64g Ti and 4.65g O.

1. Multiply the mass by molar mass to get moles

5.41g Fe x 1 mol Fe = 0.0969 mol Fe

55.85 g Fe

4.64g Ti x 1 mol Ti = 0.0969 mol Ti

47.88g Ti

4.65g O x 1 mol O = 0.291 mol O

16.00g O

Empirical Formula from Mass

Example: Determine the empirical formula for ilmenite, which contains 5.41g Fe, 4.64g Ti and 4.65g O.

2. Multiply by the smallest number to get the mole ratio.

0.0969 mol Fe : 0.0969 mol Ti : 0.291 mol O

0.0969 mol 0.0969 mol 0.0969 mol

1 : 1 : 3

3. Calculate the empirical formula.

FeTiO3

Practice

p 337 # 54-55

TEST!!!!

Stoichiometry Practice – Conservation of Mass

For the following balanced chemical equations, determine all possible mole ratios.

1. HCl(aq) + KOH(aq) KCl(aq) + H2O(l)

2. 2Mg(s) + O2(g) 2MgO(s)

3. 2HgO(s) 2Hg(l) + O2(g)

Stoichiometric CalculationsMany times we need to determine a certain amount of product

from a reaction or want to know how much product will form from a given amount of reactant.

To do this you need:

1. Balanced chemical equations

2. Mole ratios

Moles of known x moles of unknown =

moles of known

Stoichiometric Calculations: mole-mole

Example: If you put 0.0400 mol of K into water, how much hydrogen gas will be produced?

2K(s) + 2H2O(l) → 2KOH(aq) +H2(g)

Mole ratio between K and KOH is 2 mol K or 1 mol H2

1 mol H2 2 mol K

moles of known x moles of unknown

moles of known

0.0400 mol K x 1 mol H2 =

2 mol K

0.0200 mol H2

Stoichiometric Calculations Practice

1. How many moles of carbon dioxide are produced when 10.0 moles of propane (C3H8) are burned in excess oxygen in a gas grill. Water is also a product.

2. Sulfuric acid is formed when sulfur dioxide reacts with oxygen and water. Write the balanced chemical equation for the reaction. If 12.5 mol SO2 reacts, how many moles H2SO4 can be produced? How many mole O2 is needed?

Practice

p 357 #2a, 3a

p 359 #10

Stoichiometric Calculations: mole-massExample: If you put 0.0400 mol of K into water, how

many grams of hydrogen gas will be produced?

2K(s) + 2H2O(l) → 2KOH(aq) +H2(g)

Mole ratio between K and KOH is 2 mol K or 1 mol H2

1 mol H2 2 mol Kmoles of known x moles of unknown

moles of known

0.0400 mol K x 1 mol H2 =

2 mol K

0.0200 mol H2

0.0200 mol H2 x 2.02 g H2 =

1 mol H2

0.0404 g H2

Practice

p 360 #11-12

Stoichiometric Calculations: mass-mass

Example: If you put 15g of K into water, how many grams of hydrogen gas will be produced?

2K(s) + 2H2O(l) → 2KOH(aq) +H2(g)

Mole ratio between K and KOH is 2 mol K or 1 mol H2

1 mol H2 2 mol K

0.384 mol K x 1 mol H2 =

2 mol K

0.192 mol H2

15 g K x 1 mol K = 0.384 mol K

39.098 g K

0.192 mol H2 x 2.02 g H2 =

1 mol H2

0.388 g H2

Practice

p 361 #13-14

Stoichiometric Review

1. Why is a balanced chemical equation needed in solving stoichiometric calculations?

2. When solving stoichiometric problems, how is the correct mole ratio expressed?

3. List the steps in solving stoichiometric equations.

4. How many grams of zinc (II) chloride can be obtained in 50.0 g of zinc reacts completely in the following reaction:

Zn + 2HCl → ZnCl2 + H2

5. How many grams of KCl is produced when 0.500 mol of KClO3 decomposes in the following reaction:

2KClO3 + heat → 2KCl + 3O2

Stoichiometric Calculations Practice 1

1. What mass of carbon dioxide are produced when 10.0 moles of propane (C3H8) are burned in excess oxygen in a gas grill. Water is also a product.

2. Sulfuric acid is formed when sulfur dioxide reacts with oxygen and water. Write the balanced chemical equation for the reaction. If 12.5 mol SO2 reacts, how many grams of H2SO4 can be produced?

Stoichiometric Calculations Practice 2

1. What mass of carbon dioxide are produced when 20.0 grams of propane (C3H8) are burned in excess oxygen in a gas grill. Water is also a product.

2. Sulfuric acid is formed when sulfur dioxide reacts with oxygen and water. Write the balanced chemical equation for the reaction. If 10.0 grams SO2 reacts, how many grams of H2SO4 can be produced?

QUIZ!!!!

Limiting Reactants

Rarely are the reactants in a chemical reaction present in the exact mole ratios specified in the balanced equation.

-usually, one or more of the reactants are present in excess, and the reaction proceeds until all of one reactant is used up.

-the reactant that is used up is called the limiting reactant

The limiting reactant limits the reaction and, thus, determines how much of the product forms.

- The left-over reactants are called excess reactants

Limiting Reactants

In the reaction below, 40.0 g of sodium hydroxide (NaOH) reacts with 60.0 g of sulfuric acid (H2SO4).

a. Calculate the limiting reactant

b. Determine the reactant in excess

c. Calculate the mass of water produced

d. Calculate the mass of reactant in excess.

Limiting Reactants

In the reaction below, 40.0 g of sodium hydroxide (NaOH) reacts with 60.0 g of sulfuric acid (H2SO4).

a. Calculate the limiting reactant

1. calculate the actual number of moles of reactants

Limiting Reactants

In the reaction below, 40.0 g of sodium hydroxide (NaOH) reacts with 60.0 g of sulfuric acid (H2SO4).

a. determine the limiting reactant

2. look at the actual ratio to the available ratio.

~actual ratio = 1.00 mol NaOH

0.612 mol H2SO4

~available ratio = 2.00 mol NaOH = 1.00 mol NaOH

1.00 mol H2SO4 0.500 mol H2SO4

You can see that when 0.500 mol H2SO4 has reacted, all of the 1.00 mol of NaOH would be used up.

-NaOH is the limiting reactant

Limiting Reactants

In the reaction below, 40.0 g of sodium hydroxide (NaOH) reacts with 60.0 g of sulfuric acid (H2SO4).

b. determine the reactant in excess

-since NaOH is the limiting reactant, H2SO4 is the reactant in excess

c. calculate the mass of sodium sulfate produced

-use the limiting reactant

1.00mol NaOH x 1.00mol Na2SO4 x 142g Na2SO4 = 71.0g Na2SO4

2.00mol NaOH 1.00mol Na2SO4

Limiting Reactants

In the reaction below, 40.0 g of sodium hydroxide (NaOH) reacts with 60.0 g of sulfuric acid (H2SO4).

d. calculate the mass of reactant in excess

-use NaOH, the limiting reactant to determine moles and mass of H2SO4 used

~1.00mol NaOH x 1.00mol H2SO4 x 98.09g H2SO = 49.05g H2SO4

2.00mol NaOH 1.00mol H2SO4

-subtract the mass needed from the mass available.

60.0g H2SO4 – 49.05g H2SO4 = 11.0g H2SO4 in excess

Practice

p 368 # 20-21

Percent Yield

Most reactions never succeed in producing the predicted amount of product.

-not every reaction goes cleanly or completely

●liquids may stick to surfaces of containers

●liquids may vaporize/evaporate

●solids may be left behind on filter paper

●solids may be lost in the purification process

●sometimes unintended products form

The amount you have been calculating so far has been the theoretical yield, the maximum amount of product that can be produced from a given amount of reactant.

Percent Yield

A chemical reaction rarely produces the theoretical yield.

-actual yield is the amount of product produced in the chemical reaction

We can measure the efficiency of the reaction by calculating the percent yield.

-percent yield (% yield) of the product is the ratio of actual yield to the theoretical yield, expressed as a percent.

% yield = actual yield (from the experiment) x 100

theoretical yield (from calculations)

Percent Yield

When potassium chromate is added to a solution containing 0.500 g of silver nitrate, solid silver chromate is formed.

a. Determine the theoretical yield of silver chromate

b. If 0.455 g of silver chromate is actually obtained,

calculate the percent yield.