1. Moles and Stoichiometry
Ms. Randall/Mr. Palermo

2. The Mole
A mole is like a dozen. You can have a dozen of anything and it means you have 12.
A mole means you have 6.02x1023 particles which could be ATOMS or MOLECULES. (like a really big dozen)

3. Example:
1 mol of C = 6.02 x 1023 atoms = g of C

4. What do you call a mole of avacado’s?... Guacamole!!
Therefore….
1 mole is 6.02 x of anything!!
6.02 x is called Avogadro's number.
A mole can be 6.02 x 1023 atoms or molecules
What do you call a mole of avacado’s?... Guacamole!!

5. Atomic Mass Units (amu)
Used to measure the mass of an atom
Example: The AMU of 1 atom of oxygen is

6. MONOATOMIC ELEMENTS = one atom of an element -not bonded to anything
DIATOMIC ELEMENTS or DIATOMS = elements whose atoms always travel in pairs (N2, O2, F2, Cl2, Br2, I2, At2, H2)—bonded to another atom of the same element
Remember 7 up rule….

7. So, what would the mass be of one molecule of oxygen (O2)?
This means that the mass of O2 =
2 x amu = amu

8. What is the mass of F2 ?

9. Calculating Gram Formula Mass (GFM)
Gram Formula Mass- the mass of one MOLE of an atom, molecule or compound in GRAMS (g)
If there is more than 1 atom multiply the GFM by the # of atoms
Round to nearest tenth

10. Example: Calculate the GFM of H2(g) ?
Element # of atoms Mass Product
H x amu = 2.0g
(this is the mass of 1 mol H2(g))

11. Example: What is the GFM of AgNO3?
Element # of atoms Mass Product
Ag = 1 × amu = g
N = 1 × 14amu = g
O = 3 × amu = g
170.0g/mol

12. Practice:
What is the GFM of Calcium(s) ?

13. Practice:
What is the GFM of Br2(g)?

14. Practice:
What is the GFM of CO2(g)?

16. Application of the Mole: The Math of Chemistry
We will need to convert from grams to moles and vice versa for this class.
The diagram below summarizes these processes:

17. Converting Grams to Moles:
From Table T, you would use the MOLE FORMULA:
Locate the formula from Table T and fill in the box below:

18. Example: How many moles are in 4.75 g of sodium hydroxide? (NaOH)
Step 1: Calculate the GFM for the compound.
Na = 1 x = =
O = 1 x = =
H = 1 x = =
40.0g/mol

19. Step 2:
Plug the given value and the GFM into the “mole calculations” formula and solve for the number of moles.
4.75g
40.0 g/mol
= mol

20. Practice:
How many moles are in 39.0 grams of LiF?

21. Practice:
What is the number of moles of potassium chloride present in 148 g?

22. Practice:
3) How many moles are in 168 g of KOH?

23. Converting Moles to Grams:
From Table T, rearrange Mole Calculations Formula, to solve for MASS in GRAMS now: Try and rearrange the formula now…

24. Problem: You have a 2. 50 mole sample of sulfuric acid (H2SO4)
Problem: You have a 2.50 mole sample of sulfuric acid (H2SO4). What is the mass of your sample in grams?
Step 1: Calculate the GFM for the compound. H = 2 x = S = 1 x = O = 4 x =
= g/mol

25. Step 2:
Plug the given value and the GFM into the “mole calculations” formula and solve for the mass of the sample.
mass of sample (g) = # of moles (mol) x GFM (g/mol)
=(2.50 mol) x (98.1 g/mol)
=245 g H2SO4

26. Practice:
1) What is the mass of 4.5 moles of KOH?

27. Practice:
2) What is the mass of 0.50 mol of CuSO4?

28. Practice:
3) What is the mass of 1 mole of nitrogen gas?

29. Calculating Mole/Volume of a Gas
It is DIFFICULT to measure MASS of a GAS so we measure VOLUME but we have to determine how many moles of gas we have.

30. TEMPERATURE and PRESSURE affect the VOLUME of a gas….
Avogadro’s Hypothesis- at the same temperature and pressure EQUAL VOLUMES of gas have the same number of particles.
TEMPERATURE and PRESSURE affect the VOLUME of a gas….
At STP 1 mole of ANY gas occupies 22.4 L
Called the MOLAR VOLUME

32. Example What is the volume of 4.59 mole of CO2 gas at STP? 1mol=22.4L
Volume= moles x 22.4L
4.59 Moles CO2 x 22.4L = 103L of CO2

33. Example: How many moles are present in 101.0 L of CO2 gas at STP?
Moles= volume
22.4L
Moles= 101.0L = mol CO2

34. Practice
What is the volume in liters of 0.60 mol SO2 gas at STP?

35. Practice
How many moles of CO2 (g) are present in a 63.5 L sample at STP?

36. Mole Road Map

37. Percent Composition
Nutrition Facts on foods can tell you just how much of a substance you are consuming and how that relates to how much you should eat in a day.

38. Percent Composition
How much of a element or compound is in a mixture.

39. Calculating Percent Composition
Formula located on Table T (Write the formula below)

40. Ex: What is the percent composition of Calcium in CaCl2
Step 1: Calculate the GFM for the compound.
= g/mol
40.1 g
70.9 g

41. Step 2: Use the formula to find the percent composition of each element or “part” in our compound (to the nearest tenth of a %).
%Ca in CaCl2 = 40.1g = %
111.1g

42. Practice
1) What is the percentage by mass of carbon in CO2?

43. Practice
2) What is the percent by mass of nitrogen in NH4NO3?

44. Practice
3) What is the percent by mass of oxygen in magnesium oxide?

45. % Composition of Hydrates
Hydrate- Ionic solids with water trapped in the crystal lattice.
It is written like this:
Ionic Compound’s Formula • n H2O
(n) is a whole number
Notice how WATER molecules are BUILT INTO the chemical formula

46. Anhydride- A crystal with NO WATER trapped inside its lattice

47. How does water get trapped in the crystal?
Crystal lattices have open spaces between the atoms so water molecules get trapped in the spaces (like food between your teeth)
Water molecules have charged ends which are attracted to the charged ions in the crystal.

48. Everyday example of a Hydrated Crystal
The little white packet in your new shoe box that says “SILICA GEL DO NOT EAT”
The crystals absorb moisture (water) from the air so your shoes do not get mildew on them. Also used in electronics packaging.
Anhydride

49. Calculate % Composition of Hydrate
Step 1: Calculate GFM of the HYDRATE
Step 2: Plug into % composition formula from Table T
% composition by mass = mass of part X 100
mass of whole

50. Example: Step 1- Calculate GFM of Hydrate
What is the percentage by mass of water in sodium carbonate crystals (Na2CO310H2O)?
Step 1- Calculate GFM of Hydrate
GFM
= g/mol

51. Step 2- Plug values into Formula
% H2O by mass = g X 100 = 63.0 %
286.1g

52. Practice:
What is the percent by mass of water in BaCl22H2O?

53. Challenge Question

54. Example
A gram sample of hydrated crystal is heated to a constant mass of 8.72 grams. This means all of the water has been driven out by the heat.
a) Calculate the mass of water that was driven out:
b) Calculate the %mass of water in the hydrate.
1.68 grams of water driven off by the heat

55. Types of Formulas
Empirical Formula- The simplest WHOLE NUMBER ratio of atoms in a compound or molecule.
Subscripts CANNOT be reduced any further
***Ionic compounds are already empirical formulas.

56. Molecular Formula = the ACTUAL FORMULA for a compound
whole # multiples of empirical formula
Subscripts can be reduced

57. Determining Empirical Formula
Divide subscripts by the greatest common factor
Example: molecular formula = C4H10
Divide by 2 (greatest common factor)
C2H5

58. Practice: Determine empirical formula from molecular formula.

59. Practice: Determine empirical formula from molecular formula.

60. Calculating Empirical Formula from % Composition
Step 1: Always assume you have a 100 g sample (The total % for the compound must = 100%, so just change the units from % to g)
Step 2: Convert grams to moles.
Step 3: Divide all mole numbers by the smallest mole number.

61. Ex. A compound is 46. 2 % mass carbon and 53. 8 % mass nitrogen
Ex. A compound is 46.2 % mass carbon and 53.8 % mass nitrogen What is its empirical formula?
Step 1: Assume a 100 g sample.
46.2 % C = 46.2 g C
53.8 % N = 53.8 g N

62. Step 2: Convert grams to moles (have grams, need moles)
Remember use the mole formula (Table T)
46.2 g C = 3.85 mol C
12.0 g C
53.8 g N = 3.84 mol N
14.0 g N
But we must have WHOLE NUMBERS for SUBSCRIPTS…….

63. Step 3: Divide each mole number by the smallest mole number (We will round in this step to the nearest integer if it’s super close).
The RATIO of C atoms to N atoms is 1:1 therefore, the empirical formula for our compound is CN

64. Practice: Determine the empirical formula from the percent composition for each of the following.

65. Practice
A compound contains 24.0 g C and g O. Calculate its empirical formula. (Hint: start with step 2)

66. A compound contains 14. 6% C and 85. 4% Cl by mass
A compound contains 14.6% C and 85.4% Cl by mass. Calculate the empirical formula of this compound.

67. A compound contains 32. 8% chromium and 67. 2% chlorine
A compound contains 32.8% chromium and 67.2% chlorine. Calculate the empirical formula of this compound.

68. Determining MOLECULAR Formula from Empirical:
Calculate Gram Formula Mass of the EMPIRICAL FORMULA
Divide the MASS of the molecular formula by the MASS of the empirical formula.
MULTIPLY the subscript of the empirical formula by the answer
in step 2.

69. Example Step 1. Determine the GFM of the empirical formula.
The empirical formula of a compound is C2H3, and the molecular mass is 54.0 grams. What is the molecular formula?
Step 1. Determine the GFM of the empirical formula.
C2H3 = (2 C X 12.0 g/mol) + (3 H X 1.0 g/mol)
= 27.0 g/mole

70. Step 2. Divide the molecular mass by the empirical mass
Step 2. Divide the molecular mass by the empirical mass. This will give you a whole-number multiple that tells you how many times larger the molecular formula is than the empirical formula
(54.0 g/mol) / (27.0 g/mol) = 2

71. Step 3. Multiply the whole number by the empirical formula
Step 3. Multiply the whole number by the empirical formula. This will give the molecular formula.
2 X C2H3 = C4H6

72. Practice:
What is the molecular formula of a compound that has an empirical formula of NO2 and molecular mass of 92.0 g?

73. A compound has an empirical formula of HCO2 and a molecular mass of 90 grams per mole. What is the molecular formula of this compound?

74. Challenge Question Determine Molecular Formula from % Composition

75. A compound is 50% sulfur and 50% oxygen by mass
A compound is 50% sulfur and 50% oxygen by mass. Calculate the empirical formula. If its molecular formula is g, determine its molecular formula.

76. Chemical Equations
subscripts

77. REACTANTS = the STARTING substances in a chemical reaction (found to the LEFT of the arrow)
PRODUCTS = a substance PRODUCED by a chemical reaction (found to the RIGHT of the arrow)

78. COEFFICIENT = the integer in front of an element or compound
indicates the number of moles present
SUBSCRIPT = the integer to the lower right of an element
indicates the number of atoms present
*COEFFICIENTS and SUBSCRIPTS tell us how many moles we have for each element

79. Chemical Symbols (states of matter)
(s) solid
(g) gas
(l) liquid
(aq) dissolved in water (aqueous)
Example:
2Na(s) + 2H2O(l)  → 2NaOH(aq)  + H2(g)

80. Catalyst A substance that SPEEDS UP a reaction Written ABOVE the arrow
Example: (elephants toothpaste)
2 H2O2(aq) -----> 2 H2O(l) + O2(g) KI

81. Balancing Equations:
In all chemical reactions there is a CONSERVATION of mass, energy, and charge
“what goes in must come out”
The number of ATOMS of each element must be EQUAL on BOTH SIDES of the equation.
**Matter and energy cannot be created nor destroyed, only changed from one form to another

82. Checking for Conservation
Consider the equation for the formation of water:
Does it show conservation?
H2 + O2 → H2O
≠ 18
Why doesn’t it show conservation?
Consider the drawing of the reaction below:
An oxygen atom is missing on the right!
H2 + O2 → H2O
NO!! H O O H

83. The only way to get an additional oxygen on the product side is to add another water molecule.
The only way to even things out is to put another hydrogen molecule on the reactant side.
This result can be shown in the equation with coefficients. This makes the equation balanced.
H O → H2O H O 2

84. An unbalanced equation does not show conservation of mass.
H2 + O2 → H2O
≠ 18
A balanced equation shows conservation of mass.
2H O2 → 2H2O
2(2) = 2(18)
= 36

85. An equation is balanced when:
The number of ATOMS of each type present are the SAME on both sides of the equation.

86. COEFFICIENTS are ONLY used to BALANCE equations
**NOTE: WE NEVER CHANGE THE SUBSCRIPTS IN A FORMULA!

87. Steps for Balancing Equations
Step 1: Find the most complex compound in the equation. Balance the elements found in that compound on the opposite side of the arrow by changing the coefficients for those atoms.

88. Step 2: Continue balancing until all atoms are balanced (save pure elements for last)
Step 3: Go back and check each atom to see if it is balanced on both sides of the equation.
Step 4: POLYATOMIC IONS may be balanced as a SINGLE UNIT rather than as separate elements as long as they stay intact during the reaction.

89. Guided Practice: __C + __O2 → __CO2 __N2 + __H2 → __NH3
__BaCl2 + __AgNO3 → __Ba(NO3)2 + __AgCl
__C3H8 + __O2 → __ CO2 + __ H2O 3 2 2 2 5 3 4

90. Practice

91. Practice

92. Practice

93. Calculating Mole Ratios
Step 1: Make sure the equation is BALANCED. Step 2: Set up a ratio of moles of SUBSTANCES in the balanced equation to the ACTUAL MOLES.

94. Example
How many moles of oxygen are consumed when 0.6 moles of hydrogen burns to produce water?
Step 1: Write a balanced equation and determine the mole ratios from the equation.
2 H2(g) + O2(g) ➝ 2 H2O
Step 2: Identify the known and the unknown
Step 3: Set up a proportion and solve for the unknown

95. Example
How many moles of hydrogen are consumed
when 0.5 moles of oxygen combine with it to produce water?
Step 1: Write a balanced equation and determine the mole ratios from the equation.
2 H2(g) + O2(g) ➝ 2 H2O
Step 2: Identify the known and the unknown
Step 3: Set up a proportion and solve for the unknown

96. Example Consider the following formula: N2 + 3H2 → 2NH3
How many moles of nitrogen gas (N2) would be needed to produce 10 moles of ammonia (NH3)?

97. Practice
C3H O2 → 3CO H2O
If 12 moles of C3H8 react completely, how many moles of H2O are formed in the reaction above?

98. Practice
C3H8 + 5O2 → 3CO2 + 4H2O If 20 moles of CO2 are formed, how many moles of O2 reacted?

99. Practice
C3H8 + 5O2 → 3CO2 + 4H2O If 8 moles of O2 react completely, how many moles of H2O are formed?

100. Practice
N H2 → 2NH3
If 2.5 moles of N2 react completely, how many moles of NH3 are formed?

101. Types of Chemical Reactions

102. Synthesis
During a SYNTHESIS or COMBINATION reaction, substances COMBINE to form a new compound with new chemical and physical properties.
General pattern: A + B → AB
Example:
2Mg(s) + O2(g) → 2MgO(s)

103. Decomposition
One reactant BREAKS APART into TWO or more elements or compounds.
General pattern: AB → A + B
Example:
2HgO(s) → 2Hg(ℓ) + O2(g)

104. Single Replacement
A more active METAL replaces a less active METAL from its compound
General Pattern: AB + C → CB + A
Use Table J for the activity of metals and nonmetals

105. Na + KCl ® K + NaCl Cl K Na

106. F2 + 2 LiCl ® 2 LiF + Cl2 Cl Li F Cl F Li

107. Zn(s) + CuSO4(aq) Cu(s) + ZnSO4(aq)
Example:
Zn(s) + CuSO4(aq) Cu(s) + ZnSO4(aq)
Zinc replaces copper because zinc is more active than copper (HIGHER on chart)
Cu(s) + ZnSO4(aq) No Reaction
Copper cannot replace zinc because it is LOWER than ZINC on chart)

108. B + D C A Double Replacement
METALS in two aqueous compounds SWITCH places.
General Pattern: AB + CD → CB + AD
Ex. AgNO3(aq) + NaCl(aq) → NaNO3(aq) + AgCl(s)
B + D C A

109. 3NaOH + FeCl3 ® Fe(OH)3 + 3NaCl

110. Check Table F to see if there is a PRECIPITATE during a double replacement reaction.
Precipitate- A solid that forms in solution

111. AgCl is a precipitate because it is insoluble in water.

112. Practice- Name the type of reaction
4Li + O2 2 LI2O Type of Reaction: _______________

113. Practice
Al + CuCl AlCl Cu
Type of Reaction: _______________

114. Practice
2AuBr Au Br2
Type of Reaction: _______________

115. Practice
BaCl2 + Na2SO4 2NaCl + BaSO4 Type of Reaction: _______________