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CHAPTER 7
Quantitative Composition of Compounds
7.1 THE MOLE
Used to count incredibly small objects, like atoms
 Avogadro’s Number


6.02 x 1023 = 1 Mole
Practice 7.1
 How many atoms are in 1.00 mole of:
a) Fe
b) H2
c) H2SO4

7.2 MOLAR MASS OF COMPOUNDS

Molar Mass – atomic mass of an element in
grams containing Avogadro's number of atoms


Practice 7.2


Can be used as a conversion factor
What is the mass of 2.50 moles of He?
Practice 7.3

How many atoms are present in 0.025 mol of iron?

Practice 7.4


Practice 7.5


Calculate the molar mass of KNO3
What is the mass of 0.150 mol Na2SO4?
Practice 7.6

How many moles and molecules are there in 500.0 g
of HC2H3O2?
7.3 PERCENT COMPOSITION OF COMPOUNDS

Mass percent of each element in a compound


Practice 7.7


Found from the formula or experimental data
Calculate the percent composition of Ca(NO3)2
Practice 7.8

Calculate the percent composition K2CrO4

Practice 7.9

Aluminum chloride is formed by reacting 13.43g
aluminum with 53.18g of chlorine. What is the
percent composition of the compound?
7.4 EMPIRICAL FORMULA VS.
MOLECULAR FORMULA


Empirical Formula – simplest formula – the
smallest whole number ratio of atoms present in
a compound
Molecular Formula – true formula – represents
the total number of atoms of each element
present in one molecule of a compound
7.5 CALCULATING EMPIRICAL FORMULAS
Assume a definite starting quantity (usually
100.0g) of the compound, if not given, and
express the mass of each element in grams.
 Convert the grams of each element into moles
using each element's molar mass. (This step will
usually not give whole numbers.)
 Divide each value by the smallest value obtained
in step 2. If these values aren’t whole numbers,
go on to the next step.
 Multiply by the smallest number needed to make
all of the numbers from step 3 whole numbers.


Practice 7.10

Calculate the empirical formula of a compound
containing 52.14% C, 13.12% H, and 34.73% O.

Practice 7.11

Calculate the empirical formula of a compound that
contains 43.7% P and 56.3% O by mass.
7.6 CALCULATING THE MOLECULAR FORMULA
FROM THE EMPIRICAL FORMULA
The molecular formula will be equal to, or some
multiple of, the empirical formula.
 The molar mass of the compound must be given
to solve for the molecular formula.
 Divide the molar mass of the compound by the
molar mass of the empirical formula to find the
factor to multiply the coefficients by.
 Example


If a compound has an empirical formula of NO2 and a
molar mass of 92.00g, determine the molecular
formula.

Practice 7.12

Calculate the empirical and molecular formulas of a
compound that contains 80.0% C, 20.0% H and has a
molar mass of 30.00g.
HOMEWORK

Required:


Paired Exercises 2-44, even
Suggested:

Paired Exercises 1-43, odd