1. 2.C – Conserving Matter

2. Objectives State and apply the Law of Conservation of Matter.
Learn how to balance chemical equations.
Explain the mole concept.
Calculate molar mass of a compound.
Calculate percent composition of elements.
Distinguish between renewable and nonrenewable resources.

3. C.1: Keeping Track of Atoms
Complete the reading guide handout

4. C.1: Keeping Track of Atoms
In some chemical reactions, matter seems to be created.
When a nail rusts.
In other reactions, matter seems to disappear.
When paper burns and apparently vanishes
However, neither creation or destruction of matter occurs.
Matter may undergo physical or chemical changes.

5. C.1: Keeping Track of Atoms
In a car engine gasoline is burned. What happens to the molecules of gasoline?
Gasoline is made up of carbon and hydrogen atoms (C and H atoms)
When gasoline burns these atoms react with oxygen atoms in air to form carbon dioxide (CO2), carbon monoxide (CO) and water (H2O).
The original atoms of gasoline are not destroyed but become rearranged.

6. Basic chemical equation: CxHx + O2 CO2 + CO + H2O
reactants
products
Molecules can be converted or decompose by chemical reactions; but the atoms remain.

7. C.1: Keeping Track of Atoms
Law of Conservation of Matter: Matter is neither created nor destroyed.
Since chemical reactions cannot create or destroy atoms, chemical equations representing the reactions must always be BALANCED.
This means that the number of atoms of each element is the same on the reactant and product sides of a chemical equation.

8. Example:
The burning of coal is a reaction where carbon reacts with oxygen to produce carbon dioxide. The number of carbon and oxygen atoms is the same on both sides of the equation.
C O2
CO2

9. Atomic Perspective: C O2  CO2
1 Carbon atom oxygen molecule  carbon dioxide molecule
What are the reactants in this chemical equation?
What are the products in this chemical equation?
Are there the same number of atoms on both sides of the equation?
Were any atoms destroyed or created?
Was the Law of Conservation of Matter maintained?

10. Atomic Perspective: Cu (s) + O2 (g)  CuO (s)
From the burning of copper lab:
Is this reaction balanced?
Does it obey the Law of Conservation of Matter?

11. 2 Cu (s) + O2 (g)  2 CuO (s)
COEFFICIENTS - indicates the number of units of each substance involved.
Does the oxygen molecule have a coefficient?
What do the subscripts represent?
Can subscripts be removed from chemical equations?
The above reaction reads: 2 copper atoms react with 1 oxygen
to produce 2 molecules of copper oxide.

12. C.2: Accounting for Atoms
To obey the Law of Conservation of Matter atoms can not be created or destroyed.
All atoms must be accounted for on both sides of the reaction arrow.

13. Accounting for Atoms
Homework:
Complete C1 Supplement worksheet
Due:

14. C.2: Accounting for Atoms
Classwork:
Page
Questions 1-5, parts A, C and D only

15. C.3: Balancing Chemical Equations

16. Balancing Equations

17. Balancing Equations Law of Conservation of Atoms:
The number of atoms of each type of element must be the same on each side of the equation.

18. Balancing Equations Balancing hints: Balance the metals first.
Balance the non metals next.
Save the oxygen and hydrogen atoms until the end.

19. How do we Balance Equations?
Below is the chemical equation for the reaction of hydrogen gas with oxygen gas to produce water.
H2 + O2  H2O
Is this reaction balanced?

20. Hydrogen + oxygen water
Balancing Equations
Hydrogen + oxygen water
H2 + O H2O
Hydrogen and oxygen are diatomic elements.
Their subscripts cannot be changed.
The subscripts on water cannot be changed.

21. Balancing Equation H2 + O2 H2O Count the atoms on each side.
Reactant side: 2 atoms H and 2 atoms O
Product side: 2 atoms H and 1 atom O

22. Balancing Equations H2 + O2 H2O
If the subscripts cannot be altered, how can the atoms be made equal?
Adjust the number of molecules by changing the coefficients.

23. Balancing Equations H2 + O2 2H2O
Balance one atom at a time: Let’s start with balancing oxygen!
Reactants: 2 atoms of H and 2 atoms of O
Products: 4 atoms of H and 2 atoms of O
H is no longer balanced!

24. Balancing Equations 2H2 + O2 2H2O
Now try to balance the hydrogens by placing a 2 in front of the reactant H2
Reactant side: 4 atoms of H and 2 atoms of O
Product side: 4 atoms of H and 2 atoms of O
It’s Balanced!

25. How do we Balance Equations?
Number of compounds in the reaction
Coefficients
2 H2 + O2  2 H2O
Subscripts
# of atoms in a compound
Subscripts can not be changed.
Coefficients balance atoms in an equation

26. How to Balance By Inspection:1
Make a table of elements
_____ C H4 +
_____ O2 
_____ H2 O +
_____ C O2
Reactants
Products C H O

27. How to Balance By Inspection:2
Count the number of each element or ion on the reactants and products side.
Don’t forget to add all the atoms of the same element together—even if it appears in more than one compound!
_____ C H4 +
_____ O2 
_____ H2 O +
_____ C O2
Reactants
Products C 1 1 H 4 2 O 2 3

28. How to Balance By Inspection:3
Add coefficients to balance the numbers
Each time you add a coefficient, update your table with the new quantities of each atom.
_____ C H4 +
_____ 2 O2 
_____ 2 H2 O +
_____ C O2
Reactants
Products C 1 1 H 4 2 4 O 2 4 3 4

29. How to Balance By Inspection:4
Place a “1” in any empty coefficient location
Filling each coefficient location lets you and the grader know that you finished the problem rather than you left some blank because you weren’t done!
_____ 1 C H4 +
_____ 2 O2 
_____ 2 H2 O +
_____ 1 C O2
Reactants
Products C 1 1 H 4 2 4 O 2 4 3 4

30. What do Coefficients Really Mean?
CH4 + 2 O2  CO2 + 2 H2O
Total:
1 C
4 H
4 O
Total:
1 C
4 H
4 O
The equation is balanced.

31. Nitrogen + hydrogen ammonia
Balancing Equations
N2 + H NH3
Nitrogen + hydrogen ammonia
Count atoms.
Reactants: 2 atoms N and 2 atoms H
Products: 1 atom N and 3 atoms of H

32. Balancing Equations N2 + H2 2NH3 Nothing is balanced.
Balance the nitrogen first by placing a coefficient of 2 in front of the NH3.
N2 + H NH3

33. Balancing Equations N2 + 3H2 2NH3 Hydrogen is not balanced.
Place a 3 in front of H2.
Reactant side: 2 atoms N, 6 atoms H
Product side: 2 atoms N, 6 atoms H
It’s balanced!
N H NH3

34. Balancing Equations Cu + H2SO4 CuSO4 + H2O + SO2 Count Atoms:
Reactants: Cu – 1, H – 2, S – 1, O – 4
Products: Cu – 1, H – 2, S - 2, O - 7

35. Balancing Equations Sulfur is not balanced.
Place a two in front of sulfuric acid.
Count Atoms:
Reactants: Cu – 1, H – 4, S – 1, O – 8
Products: Cu – 1, H – 2, S - 2, O - 7
Cu + 2H2SO CuSO H2O + SO2

36. Balancing Equations
Hydrogen needs to be balanced so place a 2 in front of the H2O.
Count the number of atoms.
Cu + 2H2SO CuSO H2O + SO2

37. Balancing Equations Cu + 2H2SO4 CuSO4 + 2H2O + SO2
Reactants: Cu – 1, H – 4, S – 2, O – 8
Products: Cu – 1, H – 4, S – 2, O – 8
It’s balanced!

38. Balancing Equations Ca3(PO4)2 + H2SO4 CaSO4 + H3PO4 Count atoms.
Reactants: Ca – 3 atoms, P – 2 atoms, O – 12 atoms, H – 2 atoms, S – 1 atom

39. Balancing Equations Side note on Ca3(PO4)2
The subscript after the phosphate indicates two phosphate groups.
This means two P and eight O atoms.

40. Balancing Equations Count atoms in the product.
Ca3(PO4)2 + H2SO CaSO4 + H3PO4
Count atoms in the product.
Ca atoms – 1, S atom – 1, O atoms – 8; H atoms – 3, P atom – 1,

41. Balancing Equations
Balance the metal first by placing a coefficient of 3 in front of CaSO4.
Ca3(PO4)2 + H2SO CaSO4 + H3PO4

42. Balancing Equations Now balance the S atoms followed by the P atoms.
Ca3(PO4) H2SO CaSO4 + H3PO4

43. Balancing Equations
A coefficient of 2 placed in front of H3PO4 which balances both hydrogen and phosphate.
Ca3(PO4) H2SO CaSO H3PO4

44. Balancing Equations
This method of balancing equations is trial and error.
Practice.

45. Balance the following equation
Let’s Practice #1
Example:
Balance the following equation
__ HCl + __ Ca(OH)2  __ CaCl2 + __ H2O

46. Balance the following equation
Let’s Practice #2
Example:
Balance the following equation
__ H2 + __ O2  __ H2O

47. Balance the following equation
Let’s Practice #3
Example:
Balance the following equation
__ Fe + __ O2  ___ Fe2O3

48. C.4: Balancing Equations
Classwork:
Complete questions 1-6 on page
Due:

49. Balancing Equations
Homework:
Complete C3 Supplement worksheet
Due:

50. Balancing Equations
Additional Worksheets

51. C.5 The Mole Concept Definition:
Mole – Unit for counting for chemists.
Counting atoms is impractical because atoms are
so small and are not visible to the naked eye.
You can not weigh individual atoms on
laboratory balance.

52. Relating Mass to Numbers of Atoms
Introduction of three very important concepts:
The mole
Avogadro’s number
Molar mass

53. What is a counting unit?
You’re already familiar with one counting unit…a “dozen”
A dozen = 12
“Dozen” 12
A dozen doughnuts
12 doughnuts
A dozen books
12 books
A dozen cars
12 cars
A dozen people
12 people

54. A Mole of Particles Contains 6.02 x 1023 particles Avogadro’s Number
1 mole C = x 1023 C atoms
1 mole H2O = x 1023 H2O molecules
1 mole NaCl= x 1023 NaCl “molecules”

55. How big is a mole?
Enough soft drink cans to cover the surface of the earth to a depth of over 200 miles.
If you had Avogadro's number of unpopped popcorn kernels, and spread them across the United States of America, the country would be covered in popcorn to a depth of over 9 miles.
If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole.
One mole of paper clips would wrap around the earth 400 trillion times!

56. What does a “mole” count in?
A mole = 6.02  1023 (called Avogadro’s number)
6.02  1023 = 602,000,000,000,000,000,000,000
“mole”
6.02  1023
1 mole of doughnuts
6.02  1023 doughnuts
1 mole of atoms
6.02  1023 atoms
1 mole of molecules
6.02  1023 molecules

57. Mole of a substance = grams of substance/MW of substance
The mole enables chemists to move from the microscopic world of atoms and molecules to the real world of grams and kilograms.

58. Molar Mass
What is molar mass?

59. Atomic Weight (used for atoms)
Definition
Molar Mass – The mass for one mole of an atom or molecule.
Other terms commonly used for the same meaning:
Molecular Weight
Atomic Weight (used for atoms)

60. Mass for 1 mole of atoms The average atomic mass = grams for 1 mole
Average atomic mass is found on the periodic table
Element
Mass
1 mole of carbon atoms
12.01 g
1 mole of oxygen atoms
16.00 g
1 mole of hydrogen atoms
1.01 g
Unit for molar mass: g/mole or g/mol

61. One mole of carbon (12 grams) and one mole of copper (63.5 grams)

62. Molar Mass Examples: Molar mass of lithium (Li) = 6.94 g/mol
Molar mass of helium (He) = 4.00 g/mol
Molar mass of mercury (Hg) = g/mol

63. Molar Mass A molar mass of an element contains one mole of atoms.
4.00 g helium = 1mole = 6.02 x 1023 atoms.
6.94 g lithium = 1mole = 6.02 x 1023 atoms.
200.6 g mercury = 1mole = 6.02 x 1023 atoms.

64. Molar mass for molecules
The molar mass for a molecule = the sum of the molar masses of all the atoms

65. Calculating a Molecule’s Mass
To find the molar mass of a molecule: 1
Count the number of each type of atom 2
Find the molar mass of each atom on the periodic table 3
Multiple the # of atoms  molar mass for each atom 4
Find the sum of all the masses

66. Find the molar mass for CaBr2
Example: Molar Mass
Example:
Find the molar mass for CaBr2

67. Find the molar mass for CaBr2
Example: Molar Mass 1
Count the number of each type of atom
Example:
Find the molar mass for CaBr2 Ca 1 Br 2

68. Find the molar mass for CaBr2
Example: Molar Mass 2
Find the molar mass of each atom on the periodic table
Example:
Find the molar mass for CaBr2 Ca 1
40.08 g/mole Br 2
79.91 g/mole

69. Find the molar mass for CaBr2
Example: Molar Mass 3
Multiple the # of atoms  molar mass for each atom
Example:
Find the molar mass for CaBr2 Ca 1 
40.08 g/mole =
40.08 g/mole Br 2 
79.91 g/mole =
g/mole

70. Find the molar mass for CaBr2
Example: Molar Mass 4
Find the sum of all the masses
Example:
Find the molar mass for CaBr2 Ca 1 
40.08 g/mole =
40.08 g/mole Br 2 
79.91 g/mole = +
g/mole
g/mole
1 mole of CaBr2 molecules would have a mass of g

71. Molar Mass
A molar mass of a compound is the sum of the molar masses of the elements.
Example: Water, H2O:
2 H = 2 x 1 g/mole = 2 g/mole
1 O = 1 x 16 g/mole = 16 g/mole
molar mass of H20 =18 g/mol

72. Molar Mass
A molar mass of a compound is the sum of the molar masses of the elements.
Example: methane, CH4:
4 H = 4 x 1 g = 4 g/mole
1 C = 1 x 12g = 12 g/mole
molar mass of CH4 =16 g/mol

73. Example: Molar Mass & Parenthesis
Be sure to distribute the subscript outside the parenthesis to each element inside the parenthesis.
Example:
Find the molar mass for Sr(NO3)2

74. Find the molar mass for CH2Cl2
Let’s Practice #1
Example:
Find the molar mass for CH2Cl2

75. Find the molar mass for Al(OH)3
Let’s Practice #2
Example:
Find the molar mass for Al(OH)3

76. Classwork: C.6: Molar Masses
Page 163; questions1-9

77. Homework: C.5 Supplement
Worksheet
Due:

78. C.8: Molar Relationships
What is molar mass?

79. Molar mass for molecules
The molar mass for a molecule = the sum of the molar masses of all the atoms in the molecule.
Example:
1 mole of H2O = 18 grams
1 mole of H2O weighs 18 grams.

80. What happens when you do not have 18 grams
of water?
How do you calculate the number of moles of water?

81. Conversion Factors
Conversion factor – a way that can be used to convert from one unit to the other.
Example: the conversion between quarters and dollars:
4 quarters dollar
1 dollar or quarters

82. Conversion Factors Example:
Determine the number of quarters in 12 dollars?
(conversion factor)
? Quarters = 12 dollars x = 48 quarters
4 quarters
1 dollar

83. Lets return to the term dozen.
Conversion Factors
Lets return to the term dozen.
12 roses = 1 dozen roses (conversion factor)
6 roses = ½ dozen
Calculations:
6 roses____ = dozen
12 roses/dozen

84. Molar masses can be used as a conversion factor in chemical calculations.
Molar masses allow the chemist to convert from grams to moles of a compound and from moles to grams of a compound.
Grams Moles
molar mass

85. Gram/Mole Conversions
When converting between grams and moles, the molar mass is needed
Example: The molar mass of helium is 4.00 g/mol. To calculate how many grams of helium are in 2 moles of helium:
amount of He in moles amount of He in grams
2.00 mol He x = 8.00 g He
4.00 g He mol He

86. Example: Moles to Grams
How many grams are in 1.25 moles of water?

87. Example: Moles to Grams
When converting between grams and moles, the molar mass is needed
Example:
How many grams are in 1.25 moles of water? H O 2 1
1.01 g/mole
16.00 g/mole  =
2.02 g/mole +
18.02 g/mole
1 mole H2O molecules = g
1.25 mol H2O
18.02
g H2O
= _______ g H2O
22.53 1
mol H2O

88. How many grams are in 2.25 moles of iron, Fe?
Problem
How many grams are in 2.25 moles of iron, Fe?

89. How many grams are in 2.25 moles of iron, Fe?
Problem
How many grams are in 2.25 moles of iron, Fe?
Answer: 126 grams Fe

90. How many grams are in 0.375 moles of potassium, K?
Problem
How many grams are in moles of potassium, K?

91. How many grams are in 0.375 moles of potassium, K?
Problem
How many grams are in moles of potassium, K?
Answer: grams

92. How many grams are in 20 moles of methane, CH4?
Problem
How many grams are in 20 moles of methane, CH4?

93. How many grams are in 20 moles of methane, CH4?
Problem
How many grams are in 20 moles of methane, CH4?
Answer: 320 grams

94. Grams to Moles
If you are given a quantity in grams you can determine how moles you have by using the molar mass.
molar mass
Grams Moles

95. How many moles are in 25.5 g NaCl?
Let’s Practice
Example:
How many moles are in 25.5 g NaCl?

96. How many moles are in 25.5 g NaCl?
Let’s Practice
Example:
How many moles are in 25.5 g NaCl? Na Cl 1
22.99 g/mole
35.45 g/mole  = +
58.44 g/mole
1 mole NaCl molecules = g
25.5 g NaCl 1
mole NaCl
= _______ mole NaCl
0.44
58.44
g NaCl

97. Example: Grams to Moles
A chemist produced 11.9 g of aluminum, Al. How many moles of aluminum were produced?
mass of Al in grams amount of Al in moles
1 mol Al
27 g Al
11.9 g Al x = 0.44 mol Al

98. Problem
How many moles of calcium, Ca, are in 5.00 grams of calcium?

99. Problem
How many moles of calcium, Ca, are in 5.00 grams of calcium?
Answer: moles

100. Problem
How many moles of sodium chloride, NaCl, are in 75 grams of NaCl?

101. Problem
How many moles of sodium chloride, NaCl, are in 75 grams of NaCl?
Answer: moles

102. C.8: Molar Relationships
Classwork: C.8 Problems
Page 166 questions 1-3
(hold on 4)

103. Homework: C.7-1 Supplement
Worksheet
Due:

104. C.7:Equations and Molar Relationships
How are chemical equations and molar masses related?
2 CuO C Cu CO2
2 moles 1 mole moles 1mole

105. C.7:Equations and Molar Relationships
2 CuO C Cu + CO2
2 moles mole moles 1mole
159.1 g g g g
171.1 g total g total
Same total mass on the reactant and product side

106. C.7:Equations and Molar Relationships
Stoichiometry
PowerPoint

107. Example: Grams to Molecules
How many molecules are in
25.5 g NaCl?
Skip

108. Example: Grams to Molecules
How many molecules are in
25.5 g NaCl? Na Cl 1
22.99 g/mole
35.45 g/mole  = +
58.44 g/mole
1 moles NaCl molecules = g
1 mol = 6.021023 molecules
25.5 g NaCl 1
mol NaCl
6.021023
molecules NaCl
58.44
g NaCl 1
mol NaCl
= _________ molecules NaCl
2.63  1023

109. How many grams is a sample of 2.75 × 1024 molecules of SrCl2?
Let’s Practice #4
Example:
How many grams is a sample of 2.75 × 1024 molecules of SrCl2?
Skip

110. Homework: C.7-2 Supplement
Worksheet
Due:

111. C.9: Percent Composition
Percent Composition - The percent mass of each element found in a mixture

112. Mixtures
Sample problem 1: A post-1982 penny with a mass of 2.50 g is composed of 2.44 g of zinc and 0.06 g copper. What is the percent composition of each element?

113. The percent composition of the penny can be found by dividing the mass of each element by the total mass of the penny and multiplying by 100%.
Zinc: g zinc x 100% = 97.5%
2.50 g total
Copper: g Cu x 100% = 2.5%
Total percent must equal 100%

114. Relating molar mass and percent composition
Sample problem 2: The formula for the copper containing mineral chalcocite is Cu2S.
1) What percentage of copper (Cu) is in this mineral?
2) What percentage of sulfur (S) is in this mineral?

115. The formula for chalcocite indicates one mole of Cu2S contains two molecules of Cu (127.1 g) and one mole of S (32 g). The molar mass of Cu2S = g.
% Cu = mass of Cu x 100%
mass of Cu2S
% Cu: g Cu x 100% = 79.9%
159.2 g Cu2S

116. The formula for chalcocite indicates one mole of Cu2S contains two molecules of Cu (127.1 g) and one mole of S (32 g). The molar mass of Cu2S = g.
% S = mass of S x 100%
mass of Cu2S
% Cu: g Cu x 100% = 20.1%
159.2 g Cu2S

117. C.10: Percent Composition
Classwork: page 168,
problems 2 and 4

118. C.10: Percent Composition
Homework: C.10: Worksheet
percent composition
Due:

119. C.11:
Lab: Retrieving Copper
Read over the lab procedure

120. C.12: Conservation
Renewable resources: resources that can be replenished.
Examples include fresh water, air, fertile soil, plants and animals.
As long as natural cycles are not disturbed too much, supplies of renewable resources can be maintained indefinitely.

121. Nonrenewable resources: resources that can not be readily replenished.
Examples include metals, natural gas, coal and petroleum.
The length of time to replenish these resources are very large.

122. End of Unit

123. Replacing: replace a resource by finding a substitute, preferably from renewable resources.
Reusing: refurbish or repair an item to use again. Examples include car parts and printer cartridges.
Recycle: reprocess to use again. Examples include aluminum cans, newspaper and glass bottles.

124. C.12: Conservation C12: Complete worksheet.
These terms will appear again on a test or final exam!