Chapter(s) 11
Name ____________________
Period _____
I. Things you should remember


From
o
o
o
From
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o
the Moles Unit:
Identify particles as atoms, molecules (mc), and formula units (fun)
1 mole = 6.02 x 1023 atoms, molecules, or formula units
1 mole atom = mass (in grams) from the periodic table
the Naming & Formulas Unit:
How to write a formula given a chemical name
How to count the number of atoms or ions in a given compound
Review Examples:
Example 1: Determine the number of atoms/moles and ions in aluminum carbonate
1 mol _________ = ___________ formula units _________
1 formula unit _________ =
1 mole ___________ =
Example 2: Determine the moles of magnesium there are in 1.23 X 1024 atoms of magnesium.
Example 3: Determine the mass of 3.50 mol of copper.
Practice
1. Determine the number of sulfate ions in of each of the following.
a. 1 FUN (NH4)2SO4
b. 1 FUN aluminum sulfate
2. Determine the number of moles in 11.9 kg of aluminum.
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II. Quantifying Chemical Compounds
 We will now do the same thing, but we will have questions that are more detailed about the compounds.
First we need to know how to find the molar mass of the whole compound.
A. Molar Mass
1. Definition
 Remember, molar mass is the mass (in ________) of ______ mole of an element or ________________.
2. Determining Molar Mass
 The molar mass of the ______________ is the _______ of the molar masses of ________ atom in the
compound.
 The units for molar mass are ________________________________________
To determine the molar mass of a compound:
1. Determine the formula for the compound.
2. Determine the mass of 1 mole of each element in the compound using a periodic table
3. Multiply the molar mass of each element by the total number of atoms (moles) present in the
compound.
4. Take the sum for all the elements in the compound. Show your answer using the sig figs.
Ex 1: Determine the molar mass of sulfur trioxide.
Ex 2: Determine the molar mass of calcium nitrate.
Worked Out Example: Determine the molar mass of Iron(III) sulfate.
Iron (III) Sulfate = Fe3+, SO4-2  Fe2(SO4)3
Fe =
S =
O=
2 X 55.845 g (111.69)
3 X 32.066 g (96.198)
12 x 15.999 g (191.99)
Molar Mass = 399.88 g/mol Fe2(SO4)3
Practice
1. Determine the molar mass of each of the following.
a. Li2S
c. magnesium hydroxide
b. (NH4)2CO3
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3. Using Molar Mass in Calculations
You would use the molar mass of a compound in dimensional analysis just like you did with elements.
Worked Out Example: Determine the mass of 48.6 moles of sodium chloride.
Step 1: formula for sodium chloride
Na+
Cl-  NaCl
Step 2: set up a dimensional analysis problem
G: 48.6 mole NaCl
W: mass of NaCl
R: _____ g = 1 mol NaCl
Step 3: find the molar mass for NaCl because the wanted is MASS
Na = 1 X 22.990 g = 22.990g
Cl = 1 X 35.453 g = 35.453g
58.443g NaCl
Step 4: Plug it all in
48.6 mol NaCl x 58.443g NaCl = 2.80 X 103 g NaCl
1
1 mole NaCl
Tips for working molar mass problems – the R’s (relationships)
1. If mass is involved in your problem, determine molar mass of compound or element.
1 mole X = (molar mass) g X
2. If you see the terms: atoms, molecules, or formula units, use Avogadro’s number.
1 mol X = 6.02 X 1023___*___ X
* Atoms, molecules or formula units
3. If you see 2 different substances in the G and W, you have to think a little harder.
Remember,
Ions make formula units
moles of elements make moles of compounds (Ex. 1 mole CO2 = 2 moles O)
Ex 2: Determine the number of moles in 582 g of magnesium nitrate.
Ex 3: Determine the number of molecules in 47.3 g of sulfuric acid.
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Ex. 4: Determine the number of oxide ions in 3.6 moles of TiO2.
Practice
1. Determine the mass of 9.45 mol of dinitrogen trioxide. (Answer = 7.18 x 102 g N2O3)
2. Determine the mass (kg) of 5.83 x 1023 molecules of HCl. (Answer = 3.53 x 10-2 kg HCl)
3. Determine the number of molecules in 782 g of N2O3. (Answer = 6.19 x 1024 molecules N2O3)
III. Percent Composition
 The relative amounts of each element in a compound are expressed as ________________.
 The percent by _____________ of an element in a compound is the total number of grams of the
_____________________ divided by the molar mass of the ___________________ multiplied by 100%.
 Identify all of the atoms in the compound and solve using the equation below to find the abundance of
each element.
% composition = (mass element x atoms) x 100
molar mass
Note: Percentage = part over whole.
Mass of one type of element in a compound / mass of the whole compound
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Worked Out Example: Calculate the % composition of propane (C3H8). (In other words, what percent of
propane is carbon and what percent is hydrogen?)
C = 3 x 12.011g
H = 8 x 1.008g
= 36.033g
= 8.064g
44.097g
%carbon = 36.033g x 100 = 81.71% carbon
44.097g
% hydrogen = 8.064g x 100 = 18.39% hydrogen
44.097g
Ex 2: Determine the percent nitrogen in zinc nitrate.
 You might need to set up a ratio for problems where the mass of an element is being compared the mass
of a compound.
Ratios:
Element mass
Mass of compound
=
Element mass
Mass of compound
Ex 3: Determine the mass of sodium in 450. g of sodium chloride
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Practice
1. Calculate the percent composition of
a. calcium chloride
b. potassium nitrate
2. Determine the percent oxygen in calcium carbonate
3. Calculate the mass of hydrogen in
a. 350 g of C2H6
b. 2.14 g NH4Cl
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IV. Empirical Formulas
 The formula that represents the smallest whole-number ratio of the elements in a compound is called the
_______________________ formula.
 The empirical formula ________ or ____________ be the same as the ______________ molecular formula
Ex. What is the empirical formula for S2O6?
Ex. What is the empirical formula for C9H18O6?
 To determine the empirical formula for an _________________ compound, the compound is analyzed to
find the ___________ of each element in the compound. The amount can be expressed in
_____________________ , ______________________ or _____________________.
Ex. Iron oxide is composed of 78% iron.
To determine an empirical formula:
1. Change the % composition to grams. (Remember, % composition is based on 100 g of the
compound.)
2. Determine the number of moles of each element (using molar mass).
3. Determine the lowest whole number ratio of moles by dividing the smallest number of moles
into the other numbers.
4. Write subscripts using these ratios.
5. If the numbers aren’t whole numbers, multiply by a factor you know will make them whole.
Ex: 0.25 x 4 = 1; 0.33 x 3 = 1; 0.50 x 2 = 1
Ex: Determine the empirical formula for a sulfur oxide that is 60% oxygen?
Ex. Determine the empirical formula for methyl acetate, which has the following chemical analysis: 48.64%
carbon, 8.16% hydrogen, and 43.20% oxygen.
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V. Molecular Formulas
 A molecular formula is the ___________________ chemical formula for a compound

A _______________ formula can be determined after determining the _________________ formula of a
compound.

The ____________________________ of a substance must be given to determine the molecular formula.

The term _________________ mass will refer to the mass of the ______________________ formula.

The molecular formula can be the same as the empirical formula.
 To go from the molecular formula to an empirical formula, divide by the ___________________________.
Ex. Determine the empirical formulas given the following molecular formulas
C12H22O12
S2O6
C3H6O2
Determining Molecular Formulas
1. Determine the empirical formula
2. Determine the formula mass of the empirical formula
3. molar mass = multiplying factor
formula mass
(mass empirical formula)
4. Multiply each of the subscripts by the multiplying factor to create the molecular formula
Ex. If the molar mass of succinic acid is 118.1 g and its empirical formula of C2H3O2, what is its molecular
formula?
Ex 2. A compound composed of hydrogen and oxygen is analyzed and a sample of the compound yields
0.590 g of hydrogen and 9.40 g of oxygen. The molecular mass of this compound is 34.0 g/mol.
Determine the empirical formula and the molecular formula for this compound.
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Practice
1. Calculate the molecular formula of a compound whose molar mass is 60.0 g/mol and empirical formula
is CH4N.
2. A colorless liquid composed of 46.68% nitrogen and 53.32% oxygen has a molar mass of 60.01 g/mol.
What is the molecular formula?
46.68 g N
46.68 g N x 1 mol N
= 3.333 mol N
14.007 g N
3.333 mol = 1 Nitrogen
3.333 mol
53.32 g O
53.32 g O x 1 mol O = 3.333 mol O
15.999 g O
3.333 mol = 1 Oxygen
3.333 mol
NO = Empirical Formula
NO has a formula mass of 30.006.
60.01 g/mol /30.006g/mol = ~2
2x
NO = N2O2
N2O2 = Molecular Formula
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