2. Atomic Mass
Atoms are so small, it is difficult to discuss how much they weigh in grams.
Use atomic mass units.
an atomic mass unit (amu) is one twelth the mass of a carbon-12 atom.
This gives us a basis for comparison.
The decimal numbers on the table are atomic masses in amu.

3. They are not whole numbers
Because they are based on averages of atoms and of isotopes.
can figure out the average atomic mass from the mass of the isotopes and their relative abundance.
add up the percent as decimals times the masses of the isotopes.

4. Examples
There are two isotopes of carbon 12C with a mass of amu(98.892%), and 13C with a mass of amu (1.108%).
There are two isotopes of nitrogen , one with an atomic mass of amu and one with a mass of amu. What is the percent abundance of each?

5. The Mole The mole is a number.
A very large number, but still, just a number.
6.022 x 1023 of anything is a mole
A large dozen.
The number of atoms in exactly 12 grams of carbon-12.

6. The Mole
Makes the numbers on the table the mass of the average atom.

7. Representative particles
The smallest pieces of a substance.
For a molecular compound it is a molecule.
For an ionic compound it is a formula unit.
For an element it is an atom.

8. More Stoichiometry

9. Molar mass Mass of 1 mole of a substance.
Often called molecular weight.
To determine the molar mass of an element, look on the table.
To determine the molar mass of a compound, add up the molar masses of the elements that make it up.

10. Find the molar mass of
CH4
Mg3P2
Ca(NO3)3
Al2(Cr2O7)3
CaSO4 · 2H2O

11. Examples How much would 2.34 moles of carbon weigh?
How many moles of magnesium in g of Mg?
How many atoms of lithium in 1.00 g of Li?
How much would 3.45 x 1022 atoms of U weigh?

12. Percent Composition
One can find the percentage of the mass of a compound that comes from each of the elements in the compound by using this equation:
% element =
(number of atoms)(atomic weight)
(FW of the compound)
x 100
© 2009, Prentice-Hall, Inc.

13. Percent Composition So the percentage of carbon in ethane is… %C =
(2)(12.0 amu)
(30.0 amu)
24.0 amu
30.0 amu =
x 100
= 80.0%
© 2009, Prentice-Hall, Inc.

14. Mass percent of an element:
For iron in iron(III) oxide, (Fe2O3):
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15. Percent Composition Percent of each element a compound is composed of.
Find the mass of each element, divide by the total mass, multiply by a 100.
Easiest if you use a mole of the compound.
Find the percent composition of CH4
Al2(Cr2O7)3
CaSO4 · 2H2O

16. Finding Empirical Formulas
© 2009, Prentice-Hall, Inc.

17. Calculating Empirical Formulas
One can calculate the empirical formula from the percent composition.
© 2009, Prentice-Hall, Inc.

18. Calculating Empirical Formulas
The compound para-aminobenzoic acid (you may have seen it listed as PABA on your bottle of sunscreen) is composed of carbon (61.31%), hydrogen (5.14%), nitrogen (10.21%), and oxygen (23.33%). Find the empirical formula of PABA.
© 2009, Prentice-Hall, Inc.

19. Calculating Empirical Formulas
Assuming g of para-aminobenzoic acid,
C: g x = mol C
H: g x = 5.09 mol H
N: g x = mol N
O: g x = mol O
1 mol
12.01 g
14.01 g
1.01 g
16.00 g
© 2009, Prentice-Hall, Inc.

20. Calculating Empirical Formulas
Calculate the mole ratio by dividing by the smallest number of moles:
C: =  7
H: =  7
N: = 1.000
O: =  2
5.105 mol
mol
5.09 mol
1.458 mol
© 2009, Prentice-Hall, Inc.

21. Calculating Empirical Formulas
These are the subscripts for the empirical formula:
C7H7NO2
© 2009, Prentice-Hall, Inc.

22. Working backwards
From percent composition, you can determine the empirical formula.
Empirical Formula the lowest ratio of atoms in a molecule.
Based on mole ratios.
A sample is 59.53% C, 5.38%H, 10.68%N, and 24.40%O what is its empirical formula.

23. Pure O2 in
CO2 is absorbed
Sample is burned completely to form CO2 and H2O
H2O is absorbed

24. A gram sample of a compound (vitamin C) composed of only C, H, and O is burned completely with excess O g of CO2 and g of H2O are produced. What is the empirical formula?

25. Empirical To Molecular Formulas
Empirical is lowest ratio.
Molecular is actual molecule.
Need Molar mass.
Ratio of empirical to molar mass will tell you the molecular formula.
Must be a whole number because...

26. Example
A compound is made of only sulfur and oxygen. It is 69.6% S by mass. Its molar mass is 184 g/mol. What is its formula?

27. Exercise
The composition of adipic acid is 49.3% C, 6.9% H, and 43.8% O (by mass). The molar mass of the compound is about 146 g/mol.
What is the empirical formula?
C3H5O2
What is the molecular formula?
C6H10O4
Assume 100.0g g of carbon is mol C (49.3/12.01) g of hydrogen is mol H (6.9/1.008) g of oxygen is mol O (43.8/16.00). The ratio of C:H:O is 1.5:2.5:1 (C: 4.105/ = 1.5; H: 6.845/ = 2.5). Multiplying each by 2 to get a whole number becomes a ratio of 3:5:2. The empirical formula is therefore C3H5O2. The molar mass of the empirical formula is g/mol, which goes into the molar mass of the molecular formula 2 times (146/73.07). The molecular formula is therefore C6H10O4.
Copyright © Cengage Learning. All rights reserved

28. Chemical Equations Are sentences.
Describe what happens in a chemical reaction.
Reactants ® Products
Equations should be balanced.
Have the same number of each kind of atoms on both sides because ...

29. Meaning
A balanced equation can be used to describe a reaction in molecules and atoms.
Not grams.
Chemical reactions happen molecules at a time
or dozens of molecules at a time
or moles of molecules.

30. Stoichiometry
Given an amount of either starting material or product, determining the other quantities.
use conversion factors from
molar mass (g - mole)
balanced equation (mole - mole)
keep track.

31. Examples How many moles is 4.56 g of CO2 ?
How many grams is 9.87 moles of H2O?
How many molecules in 6.8 g of CH4?
49 molecules of C6H12O6 weighs how much?

32. Examples
One way of producing O2(g) involves the decomposition of potassium chlorate into potassium chloride and oxygen gas. A 25.5 g sample of Potassium chlorate is decomposed. How many moles of O2(g) are produced?
How many grams of potassium chloride?
How many grams of oxygen?

33. Examples
A piece of aluminum foil 5.11 in x 3.23 in x in is dissolved in excess HCl(aq). How many grams of H2(g) are produced?
How many grams of each reactant are needed to produce 15 grams of iron form the following reaction? Fe2O3(s) + Al(s) ® Fe(s) + Al2O3(s)

34. Examples K2PtCl4(aq) + NH3(aq) ® Pt(NH3)2Cl2 (s)+ KCl(aq)
what mass of Pt(NH3)2Cl2 can be produced from 65 g of K2PtCl4 ?
How much KCl will be produced?
How much from 65 grams of NH3?

35. Gases and the Mole

36. Gases Many of the chemicals we deal with are gases.
They are difficult to weigh.
Need to know how many moles of gas we have.
Two things effect the volume of a gas
Temperature and pressure
Compare at the same temp. and pressure.

37. Standard Temperature and Pressure
0ºC and 1 atm pressure
abbreviated STP
At STP 1 mole of gas occupies 22.4 L
Called the molar volume
Avagadro’s Hypothesis - at the same temperature and pressure equal volumes of gas have the same number of particles.

38. Examples What is the volume of 4.59 mole of CO2 gas at STP?
How many moles is L of O2 at STP?
What is the volume of 8.8g of CH4 gas at STP?

39. Density of a gas D = m /V for a gas the units will be g / L
We can determine the density of any gas at STP if we know its formula.
To find the density we need the mass and the volume.
If you assume you have 1 mole than the mass is the molar mass (PT)
At STP the volume is 22.4 L.

40. Examples Find the density of CO2 at STP.
Find the density of CH4 at STP.

41. The other way
Given the density, we can find the molar mass of the gas.
Again, pretend you have a mole at STP, so V = 22.4 L.
m = D x V
m is the mass of 1 mole, since you have 22.4 L of the stuff.
What is the molar mass of a gas with a density of g/L?
2.86 g/L?

42. Stoichiometry Greek for “measuring elements”
The calculations of quantities in chemical reactions based on a balanced equation.
We can interpret balanced chemical equations several ways.

43. Look at it differently 2H2 + O2 ® 2H2O
2 dozen molecules of hydrogen and 1 dozen molecules of oxygen form 2 dozen molecules of water.
2 x (6.02 x 1023) molecules of hydrogen and 1 x (6.02 x 1023) molecules of oxygen form 2 x (6.02 x 1023) molecules of water.
2 moles of hydrogen and 1 mole of oxygen form 2 moles of water.

44. Mole to mole conversions
2 Al2O3 ® 4Al + 3O2
every time we use 2 moles of Al2O3 we make 3 moles of O2
2 moles Al2O3
3 mole O2 or
3 mole O2
2 moles Al2O3

45. Mole to Mole conversions
How many moles of O2 are produced when 3.34 moles of Al2O3 decompose?
2 Al2O3 ® 4Al + 3O2
3.34 moles Al2O3
3 mole O2 =
5.01 moles O2
2 moles Al2O3

46. Your Turn
2C2H2 + 5 O2 ® 4CO2 + 2 H2O
If 3.84 moles of C2H2 are burned, how many moles of O2 are needed?
How many moles of C2H2 are needed to produce 8.95 mole of H2O?
If 2.47 moles of C2H2 are burned, how many moles of CO2 are formed?

47. Periodic Table
Balanced Equation
Periodic Table
Mass g A
MolesA
MolesB
Mass g B
Decide where to start based on the units you are given and stop based on what unit you are asked for

48. For example...
If 10.1 g of Fe are added to a solution of Copper (II) Sulfate, how much solid copper would form?
Fe + CuSO4 ® Fe2(SO4)3 + Cu
2Fe + 3CuSO4 ® Fe2(SO4)3 + 3Cu
1 mol Fe
63.55 g Cu
10.1 g Fe
3 mol Cu
55.85 g Fe
2 mol Fe
1 mol Cu
17.3 g Cu =

49. More Examples
To make silicon for computer chips they use this reaction
SiCl4 + 2Mg ® 2MgCl2 + Si
How many moles of Mg are needed to make 9.3 g of Si?
3.74 mol of Mg would make how many moles of Si?
How many grams of MgCl2 are produced along with 9.3 g of silicon?

50. For Example The U. S. Space Shuttle boosters use this reaction
3 Al(s) + 3 NH4ClO4 ® Al2O3 + AlCl3 + 3 NO + 6H2O
How much Al must be used to react with 652 g of NH4ClO4 ?
How much water is produced?
How much AlCl3?

51. Gases and Reactions

52. We can also change Liters of a gas to moles At STP
0ºC and 1 atmosphere pressure
At STP 22.4 L of a gas = 1 mole
If 6.45 moles of water are decomposed, how many liters of oxygen will be produced at STP?

53. For Example
If 6.45 grams of water are decomposed, how many liters of oxygen will be produced at STP?
H2O ® H2 + O2
2H2O ® 2H2 + O2
1 mol H2O
1 mol O2
22.4 L O2
6.45 g H2O
18.02 g H2O
2 mol H2O
1 mol O2

54. Your Turn
How many liters of CO2 at STP will be produced from the complete combustion of 23.2 g C4H10 ?
What volume of oxygen will be required?

55. How much you get from an chemical reaction
Yield
How much you get from an chemical reaction

56. Limiting Reagent
If you are given one dozen loaves of bread, a gallon of mustard and three pieces of salami, how many salami sandwiches can you make?
The limiting reagent is the reactant you run out of first.
The excess reagent is the one you have left over.
The limiting reagent determines how much product you can make

57. Limiting Reagent
Reactant that determines the amount of product formed.
The one you run out of first.
Makes the least product.
Book shows you a ratio method.
It works.
So does mine

58. Limiting reagent
To determine the limiting reagent requires that you do two stoichiometry problems.
Figure out how much product each reactant makes.
The one that makes the least is the limiting reagent.

59. How do you find out? Do two stoichiometry problems.
The one that makes the least product is the limiting reagent.
For example
Copper reacts with sulfur to form copper ( I ) sulfide. If 10.6 g of copper reacts with 3.83 g S how much product will be formed?

60. If 10. 6 g of copper reacts with 3. 83 g S
If 10.6 g of copper reacts with 3.83 g S. How many grams of product will be formed?
2Cu + S ® Cu2S
Cu is Limiting Reagent
1 mol Cu
1 mol Cu2S
g Cu2S
10.6 g Cu
63.55g Cu
2 mol Cu
1 mol Cu2S
= 13.3 g Cu2S
= 13.3 g Cu2S
1 mol S
1 mol Cu2S
g Cu2S
3.83 g S
32.06g S
1 mol S
1 mol Cu2S
= 19.0 g Cu2S

61. Example
Ammonia is produced by the following reaction N2 + H2 ® NH What mass of ammonia can be produced from a mixture of 100. g N2 and 500. g H2 ?
How much unreacted material remains?

62. How much excess reagent?
Use the limiting reagent to find out how much excess reagent you used
Subtract that from the amount of excess you started with

63. Excess Reagent The reactant you don’t run out of.
The amount of stuff you make is the yield.
The theoretical yield is the amount you would make if everything went perfect.
The actual yield is what you make in the lab.

64. Your turn Mg(s) +2 HCl(g) ® MgCl2(s) +H2(g)
If 10.1 mol of magnesium and 4.87 mol of HCl gas are reacted, how many moles of gas will be produced?
How much excess reagent remains?

65. Your Turn II
If 10.3 g of aluminum are reacted with 51.7 g of CuSO4 how much copper will be produced?
How much excess reagent will remain?

67. Percent Yield % yield = Actual x 100% Theoretical
% yield = what you got x 100% what you could have got

68. Yield The amount of product made in a chemical reaction.
There are three types
Actual yield- what you get in the lab when the chemicals are mixed
Theoretical yield- what the balanced equation tells you you should make.
Percent yield = Actual x 100 % Theoretical

69. Example
6.78 g of copper is produced when 3.92 g of Al are reacted with excess copper (II) sulfate.
2Al + 3 CuSO4 ® Al2(SO4)3 + 3Cu
What is the actual yield?
What is the theoretical yield?
What is the percent yield?
If you had started with 9.73 g of Al, how much copper would you expect?

70. Examples
Aluminum burns in bromine producing aluminum bromide. In a laboratory 6.0 g of aluminum reacts with excess bromine g of aluminum bromide are produced. What are the three types of yield.

71. Examples
Years of experience have proven that the percent yield for the following reaction is 74.3% Hg + Br2 ® HgBr If 10.0 g of Hg and 9.00 g of Br2 are reacted, how much HgBr2 will be produced?
If the reaction did go to completion, how much excess reagent would be left?

72. Examples
Commercial brass is an alloy of Cu and Zn. It reacts with HCl by the following reaction Zn(s) + 2HCl(aq) ® ZnCl2 (aq) + H2(g) Cu does not react. When g of brass is reacted with excess HCl, g of ZnCl2 are eventually isolated. What is the composition of the brass?