1. Stoichiometry Objectives
Identify the quantitative relationships in a balanced chemical chemical equation.
Determine the mole ratios from from a balanced chemical equation.
Explain the sequence of steps used in solving stoichiometric problems.
Use the steps to solve stoichiometric problems.
Identify the limiting reactant in a chemical equation.
Identify the excess reactant and calculate the amount remaining after the reaction is complete.
Calculate the mass of a product when the amounts of more than one reactant are given.
Calculate the theoretical yield of a chemical reaction from data.
Determine the percent yield for a chemical reaction.

2. 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

3. The Mole
Through experimentation, it has been established 1 mole = 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:

4. One Mole Quantities

5. Other Mole & Stoichiometry Vocabulary
1. What is stoichiometry?
Study of quantitative relationships among amounts of reactants and products.
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

6. 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.

7. 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.

8. Mole Practice 1
Identify and calculate the number of representative particles in each of the following quantities.
moles of gold
mole of nitrogen oxide
moles of potassium bromide
Calculate the number of moles of the substance that contains the following number of representative particles.
x 1023 atoms of barium
x 1023 molecules of carbon monoxide
x 1023 formula units of potassium iodide

9. Practice-Homework p 311 # 1-3; p 312 # 4a-c
Determine the number of atoms in 2.50 mol Zn
Determine the number of formula units in 3.25 mol AgNO3
Determine the number of molecules in 11.5 mol H2O
Determine the number of moles in
a x 1024 atoms Al
b x 1024 molecules CO2
c x 1023 formula units ZnCl2

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

11. Mass to Moles (11.2)
How many moles of copper are in a roll of copper that has a mass of 848g?

12. Practice-Homework p 316 # 11ab-12ab Determine the mass in grams of
a mol Al
b mol Si
Determine the number of moles of
a g Ag
b g S

13. 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.
Determine moles

14. 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.

15. Atoms to Mass Example problem 11-5, p 318
A party balloon has 5.50x1022 atoms of helium. What is the mass in grams of the helium?
5.50x1022 atoms He x mol He = mol He
6.02 x1023 atoms He
mol He x g He = g He
1 mol He

16. Mole Practice 2 How many atoms are in the following samples?
g cobalt
g cesium
How many grams are in the following samples?
x1023 atoms of radium
x 1020 atoms of cadmium

17. Practice-Homwork
p 316 # 11ab-12ab, p 318 # 13ab-14ab

18. 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( g) + 3(1.008 g) = g/mol

19. Practice
p 322 # 25
P 318 # 13ab-14ab
p322
25. Determine the molar mass of each of the following:
NaOH, CaCl2, Sr(NO3)2
p318
How many atoms are in
a g Li b g Pb
What is the mass of
a x1023 atoms Bi
b x1024 atoms Mn

20. -Mole relationships from a formula (p 321)
Determine the number of moles of aluminum ions in 1.25 moles of aluminum oxide (Al2O3).
First we need the ratio of Al ions to Al2O3.
2 mol Al ions
1 mole Al2O3
1.25 mol Al2O3 x 2 mol Al ions = mol Al ions mole Al2O3

21. -Mole relationships from a formula (p 321)
Determine the number of moles of chloride ions in 2.50 mol ZnCl2.
Calculate the number of moles of each element in 1.25 mole glucose (C6H12O6).

22. -Mole to Mass for compounds ( p 323)
What is the mass of 2.50 moles of allyl sulfide, (C3H5)2S?
Calculate the mass of allyl sulfide.
6(12.01 g/mol) = g/mol
10(1.01 g/mol) = g/mol
1(32.07 g/mol) = g/mol
g/mol

23. -Mole to Mass for compounds ( p 323)
What is the mass of 2.50 moles of allyl sulfide, (C3H5)2S?
2. Convert the moles to mass.
2.50mol (C3H5)2S x g (C3H5)2S
1 mol (C3H5)2S
= 286g (C3H5)2S

24. Mass of Compound to Moles
Calculate the number of moles of water that are in 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

25. 2. Now you can use the molar mass of water as a conversion factor to determine moles of water.
-Notice kg is converted to x 103 g

26. Practice
p 323 # 27-28, p 324 # 30a&b
27. What is the mass of 3.25 moles H2SO4?
What is the mass of 4.35x10-2 moles of ZnCl2?
Determine the number of moles present in each of the following:
a g AgNO3
b g ZnSO4

27. Mole Practice 3 Calculate the molar mass of the following: C2H5OH HCN
What is the mass of the following:
2.25 moles of KMnO4
1.56 moles of H2O
Determine the number of moles in the following:
35.0 g HCl
254 g PbCl4
What is the mass in grams of one molecule of the following:
7. H2SO4

28. 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.

29. 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

30. Percent Composition
Example: Determine the percent composition of calcium chloride (CaCl2).
2. Calculate molar mass of CaCl2.
g Ca+2 ions g Cl- ions = g CaCl2
1 mole CaCl mole CaCl2
3. Determine percent by mass of each element.

31. Percent Composition
Example: Determine the percent composition of calcium chloride (CaCl2).
3. Determine percent by mass of each element.
% Ca = g Ca+2 x 100 = % Ca+2
g CaCl2
% Cl = g Cl x 100 = % Cl-
4. Make sure your percent compositions equal 100%.
36.11% Ca % Cl- = 100%

32. Practice
p 331 # 43, 45
Calculate the percent composition of sodium sulfate (Na2SO4).
45. What is the percent composition of phosphoric acid (H3PO4).

33. 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.

34. 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.

35. 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.

36. Empirical Formula 0.6995 mol 0.6995 mol 0.6995 mol
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.
mol Mn : mol C : mol O
mol mol mol
: :

37. 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
: :
4. Determine the empirical formula.
MnC2O4

38. Practice
p 333 # 46-47
A blue solid is found to contain 36.84% N and 63.16% O. What is the empirical formula?
Determine the empirical formula for a compound that contains 35.98% Al and 64.02% S.

39. 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.

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

41. 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,

42. 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

43. Practice
p 335 # 51, 53
51. A substance has a chemical composition of 65.45% C, 5.45% H and 29.09% O. The molar mass of the molecular formula is g/mol. Determine the molecular formula.
53. A compound contains % N and % O. It has a molar mass of g/mol. What is the molecular formula?

44. 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 = mol Fe
55.85 g Fe
4.64g Ti x 1 mol Ti = mol Ti
47.88g Ti
4.65g O x 1 mol O = mol O
16.00g O

45. 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.
mol Fe : mol Ti : mol O
mol mol mol
: :
3. Calculate the empirical formula.
FeTiO3

46. Practice
p 337 # 54-55

47. TEST!!!!

48. Stoichiometry Practice – Conservation of Mass
For the following balanced chemical equations, determine all possible mole ratios.
HCl(aq) + KOH(aq)  KCl(aq) + H2O(l)
2. 2Mg(s) + O2(g)  2MgO(s)

49. Stoichiometric Calculations
Many 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
3. molar mass

50. Stoichiometric Calculations: mole-mole
Example: If you put mol of K into water, how much hydrogen gas will be produced?
2K(s) + 2H2O(l) → 2KOH(aq) +H2(g)
Use ‘modified’ RICE table:
R : Reaction (must be balanced)
I : Initial (amount in moles)
C : Change (also in moles)
E : “End” (really means equilibrium, but….)

51. Stoichiometric Calculations: mole-mole
Example: If you put mol of K into water, how many moles of hydrogen gas will be produced?
R : 2K(s) H2O(l) → 2KOH(aq) H2(g)
(0.400 mol) (_____ mol)
I :
C : /2(0.400)
E :
~subtract from reactants, add to products
~products start at ‘0’ mol
~what about water? it is in ‘excess’.
~why multiply by ½? b/c K:H2 = 2:1 mole ratio (1/2 as much H2 as K)

52. Stoichiometric Calculations Practice
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?

53. Practice p 359 #10 10. CH4(g) + S8(s)  CS2(l) + H2S(g)
a. Balance the equation.
b. Calculate moles of CS2 produced when 1.50 mol S8 is
used.
c. How many mol H2S is produced?

54. Stoichiometric Calculations: mole-mass
Example: If you put mol of K into water, how many grams of hydrogen gas will be produced?
R : 2K(s) H2O(l) → 2KOH(aq) H2(g)
( mol) (_____ g)
I :
C : /2(0.0400)
E : g
~must change moles into grams if asked for mass!!

55. Practice
p 360 #11-12
If you begin with 1.25 mol TiO2, what mass of Cl2 is
needed? TiO2 + C + Cl2  TiCl4 + CO2
12. Sodium chloride is decomposed into the elements sodium and chlorine by means of electrical energy. How many grams of chlorine gas are produced from 2.50 mol sodium chloride?

56. Stoichiometric Calculations: mass-mass
Example: If you put 15.0g of K into water, how many grams of hydrogen gas will be produced?
R : 2K(s) H2O(l) → 2KOH(aq) H2(g)
15.0g (_____ g)
I :
C : /2(0.384)
E :
0.388 g
~now we must change initial mass to moles!!

57. Practice
p 361 #13-14
Determine the mass of N2 produced if 100.0g NaN3 is
decomposed. NaN3(s)  Na(s) + N2(g)
14. If 2.50 g sulfur dioxide reacts with excess oxygen and water, how many grams of sulfuric acid are produced?

58. Extra Practice
p #61, 64, 70

59. QUIZ!!!!

60. - The left-over reactants are called excess reactants
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

61. 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.

62. Limiting Reactants R : 40.0g 60.0 g (_____ g) I : 1.00 0.612 0
~change both initial masses to moles
~compare ratio, I need a NaOH : H2SO4 of 2 : 1, so ask:
“Do I have 2x as much NaOH as H2SO4?”
No, b/c 2(0.612) = 1.22 and I only have 1.00, therefore
NaOH is my limiting reactant and H2SO4 is excess.

63. Limiting Reactants R : 40.0g 60.0 g (_____ g) I : 1.00 0.612 0
~change mol product to mass
show work:
~to determine amount in excess, take moles excess and change to mass

64. Practice
p 368 # 20-21

65. 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.

66. 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)

67. Percent Yield
When potassium chromate is added to a solution containing g of silver nitrate, solid silver chromate is formed.
a. Determine the theoretical yield of silver chromate
b. If g of silver chromate is actually obtained,
calculate the percent yield.