Stoichiometry Objectives

1
Stoichiometry Objectives

1. 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.

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 its 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 6.022 136 7 x 1023 representative
  particles.
• -representative particle any particle such
  as atoms, molecules, formula units, electrons, or
  ions.
• -called Avogadros 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 Avogadros number to three
  significant figures 6.02 x 1023.
• - If you write out Avogadros number, it
  looks like this
• - 602 000 000 000 000 000 000 000

4
One Mole Quantities
5
Other Mole Stoichiometry Vocabulary

• 1. What is stoichiometry?

Study of quantitative relationships among amounts
of reactants and products.

1. 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 Avogadros
  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 Avogadros number
  as a conversion factor.

8
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

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. 5.75 x 1024 atoms Al
• b. 3.75 x 1024 molecules CO2
• c. 3.58 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 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.

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. 3.57 mol Al
• b. 42.6 mol Si
• Determine the number of moles of
• a. 25.5 g Ag
• b. 300.0 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 Avogadros
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 1 mol He
  0.0914 mol He
• 6.02 x1023 atoms He
• 0.0914 mol He x 4.00 g He 0.366 g He
• 1 mol He

16
Mole Practice 2

• How many atoms are in the following samples?
• 1. 1.24 g cobalt
• 2. 0.575 g cesium
• How many grams are in the following samples?
• 3. 4.16 x1023 atoms of radium
• 4. 1.50 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(14.007 g) 3(1.008 g)
17.031 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. 55.2 g Li b. 0.230 g Pb
• What is the mass of
• a. 6.02x1023 atoms Bi
• b. 1.00x1024 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 2.50 mol Al
  ions 1 mole Al2O3

21

• -Mole relationships from a formula (p 321)
• P321 20-21
• 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) 72.06 g/mol
• 10(1.01 g/mol) 10.10 g/mol
• 1(32.07 g/mol) 32.07 g/mol
• 114.23 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 114.23g (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 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
25

• 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

26
Practice

• p 323 27-28, p 324 30ab
• 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. 22.6 g AgNO3
• b. 6.50 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 compositiona 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 Ca2 ions and
  2mol Cl- ions.
• 1mol Ca2 ions x 40.08g Ca2 ions
  40.08g Ca2 ions
• 1mol Ca2
  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.
• - 40.08g Ca2 ions 70.90g Cl- ions
  110.98 g CaCl2
• 1 mole CaCl2
  1 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 40.08 g Ca2 x 100 36.11 Ca2
• 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 Ca2 63.98 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 compounds 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 compounds empirical
  formula.
• 1. Calculate the number of moles of each element
  in the 100 g of compound.

36
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 compounds 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

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 compounds empirical
  formula.
• 3. Obtain the simplest whole-number ratio of
  moles
• Mn
  C O
• 1 2
  4
• 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
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 CHO is 111, 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 110.0 g/mol.
  Determine the molecular formula.
• 53. A compound contains 46.68 N and 53.32 O.
  It has a molar mass of 60.01 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 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

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.
• 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

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 0.0400 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 0.0400 mol of K into water,
  how many moles of hydrogen gas will be produced?
• R 2K(s) 2H2O(l) ? 2KOH(aq)
  H2(g)
• (0.400 mol)
  (_____ mol)
• I 0.400 0
• C -0.400 1/2(0.400)
• E 0 0.200
• subtract from reactants, add to products
• products start at 0 mol
• what about water? it is in excess.
• why multiply by ½? b/c KH2 21 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 0.0400 mol of K into water,
  how many grams of hydrogen gas will be produced?
• R 2K(s) 2H2O(l) ? 2KOH(aq)
  H2(g)
• (0.0400 mol)
  (_____ g)
• I 0.0400 0
• C -0.0400 1/2(0.0400)
• E 0 0.0200
• 0.0404 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) 2H2O(l) ? 2KOH(aq)
  H2(g)
• 15.0g
  (_____ g)
• I 0.384 0
• C -0.384
  1/2(0.384)
• E 0 0.192
• 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 379-380 61, 64, 70

59

• QUIZ!!!!

60
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
• C -1.00 -1/2(1.00)
  1.00
• E 0 0.112
  1.00
• 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
• C -1.00 -1/2(1.00)
  1.00
• E 0 0.112
  1.00
• change mol product to mass
• show work
• to determine amount in excess, take moles excess
  and change to mass
• show work

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 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.