1
Chapters
9 & 10
22
GOAL
Apply concepts of the
mole and Avogadro’s
number to conceptualize
and calculate empirical and
molecular formulas, mass,
moles and molecules
relationships.
3
Molar Mass
• The mass of one mole of an element or compound.
• Review Clip
• Examples:
– What is the
– What is the
– What is the
– What is the
– What is the
mass
mass
mass
mass
mass
of one
of one
of one
of one
of one
mole of Carbon?
mole of Bromine?
mole of KCl?
mole of NaOH?
mole of AgNO3?
4
1)
2)
3)
4)
5)
Determine the molar mass of the following compounds:
NaBr
PbSO4
Ca(OH)2
Na3PO4
(NH4)2CO3
1) 102.9 g/mol
2) 303.3 g/mol
3) 74.1 g/mol
4) 164.0 g/mol
5) 96.0 g/mol
5 Law
of Definite Proportions &
Law of Multiple Proportions
Law of Definite
Proportions:
Law of Multiple Proportions:
• In a sample of any • Whenever the same two
chemical
elements form more than one
compound, the
compound, the different
masses of the
masses of one element that
elements are
combine with the same mass
always in the
of the other elements are in
same proportions.
the ratio of small whole
numbers
6
Law of Definite Proportions
(sometimes called Proust's Law)
 A given compound always contains elements in a
certain proportion by mass. (Constant composition).
 Regardless of the amount, a compound is always
composed of the same elements in the same
proportion by mass.
 The proportions are found by calculating the
percent by mass.
MASS compound = sum of MASSES elements
% by mass = MASSelement x 100%
MASScompd
7
Law of Definite Proportions
Atoms combine in whole number ratios, so their
proportion by mass will always be the same.
Example: H2O is always made up of 2 atoms of
H and one atom of O.
What is the ratio of H to O?
H -1amu (x2)=2 and O-16 amu
The ratio of O to H in water is always 16:2 or 8:1.
MASS compound = sum of MASSES elements
% by mass = MASSelement x 100%
MASScompd
The make up of a compound can be expressed
as…
8
Percent Composition
• The relative
amounts of the
elements in a
compound are
expressed as
percent
composition.
9
Percent Composition
Ex. Problem
• When a 13.60g
sample of a
compound
containing only
magnesium and
oxygen is
decomposed, 5.40g
of oxygen is
obtained. What is
the percent
composition of this
compound?
10
1
% Composition Practice
2
Calculating %Comp. from the Chemical Formula
Example
• Propane
C3 H 8 .
Calculate
the percent
composition
of propane.
11
Calculating %Comp. from the Chemical Formula
12
13
• Determine
the
percentage
composition
of sodium
carbonate,
Na2CO3 .
Practice
More practice:
14
Determine the percentage composition of each of the
following compounds:
a. sodium oxalate, Na2C2O4
• ans: 34.31% Na, 17.93% C, 47.76% O
b. ethanol, C2H5OH
• ans: 52.13% C, 13.15% H, 34.72% O
c. aluminum oxide, Al2O3
• ans: 52.92% Al, 47.08% O
d. potassium sulfate, K2SO4
• ans: 44.87% K, 18.40%
15
Law of Multiple Proportions
(John Dalton, 1801)
 Mass percentages of elements in a compound
do NOT depend on amount.
Compounds with the same mass proportions
must be the same compound
Youtube
Clip
EX: Carbon combines with oxygen to form CO and CO2 .
1:1.33
2.66/1.33 = 2
1:2.66
16
Law of Multiple Proportions
(John Dalton)
Empirical Formulas
17
• Basic ratio of the elements contained in the
compound.
• The lowest whole-number ratio of the atoms of the
elements in a compound.
• May or may NOT be the same as the molecular
formula.
• Example: Hydrogen Peroxide
» Molecular Formula is H2O2
» Empirical Formula is HO
Determining
Empirical Formula
of a Compound
Ex: A compound is
analyzed and found
to contain 25.9%
nitrogen and 74.1%
oxygen. What is
the empirical
formula of the
compound?
18
#1: Convert into grams….
2) Find Mole Ratio
4) Multiply by smallest whole
number to get whole #
subscripts.
3) Divide by smallest # of mol
Practice: A
solid compound
is found to
contain K, S,
and O with the
percent
composition
listed below.
What is the
empirical
formula of this
compound?
1) Find Mole Ratio
2) Divide by smallest # of mol
K: 41.09%
S: 33.70%
O: 25.22%
19
3) Multiply by smallest whole
number to get whole #
subscripts.
Empirical Formula Practice
20
1) An oxide of chromium is found to have the
following % composition: 68.4% Cr and 31.6% O.
What is the empirical formula? (click to see worked out answer)
2) The % composition of a compound was found to be
63.5 % Ag, 8.2% N, and 28.3% O. What is the
empirical formula? (click to see worked out answer)
3) A 170g sample of an unidentified compound
contains 29.84g Na, 67.49g Cr, and 72.67g O. What
is the empirical formula? (hint: find %s 1st) (click to see worked out answer)
More problems with soln video: https://www.youtube.com/watch?v=MlaZnRbQN8g
Practice
21
22
Molecular Formulas
Molecular formula:
-whole-number multiple
of its empirical formula
(or the same as the empirical
formula)
• You can calculate the
molecular formula if
you know…
– #1 the empirical
formula AND
– #2 the molar mass of
the compound.
Steps
1.Calculate EFM (Empirical
Formula Mass)
2.Divide the Molar Mass by
the EMF
•
This tells you the # of
empirical formula units in a
molecule of the compound.
AND it is the multiplier to
convert the empirical formula
to the molecular formula
3.Multiple the subscripts in
the empirical by the answer
to step 2.
Example:
Steps
Hydrogen Peroxide
1.Calculate EFM (Empirical
Empirical Formula: HO
Molar Mass: 34.0 g/mol
2.Divide the Molar Mass
by the EMF
 Step 1
HO EFM= 17.0 g/mol
 Step 2
34.0/17.0 = 2
 Step 3
• H2O2
Formula Mass)
•
This tells you the # of
empirical formula units in a
molecule of the compound.
AND it is the multiplier to
convert the empirical
formula to the molecular
formula
3.Multiple the subscripts
in the empirical by the
answer to step 2.
23
Practice:
24
Calculate the molecular formula of a
compound whose molar mass is 60.0 g/mol
& the empirical formula is CH4N.
1. Calculate EFM
(Empirical Formula
Mass)
2. Divide the Molar
Mass by the EMF
3. Multiple the
subscripts in the
empirical by the
answer to step 2.
Practice:
25
(Determining Molecular formula from Empirical Formula)
1) What is the molecular formula of ethylene glycol,
which is used as antifreeze. The molar mass is 62
g/mol and the empirical formula is CH3O.
C2H6O2
1) What is the molecular formula of a compound with
the empirical formula CClN and a molar mass of
184.5 g/mol?
C Cl N
3
3
3
1) What is the molcular formula of a compound that is
56.6% K, 8.7% C, and 34.7% O (MM=138.21) ?
K2CO3
Hydrates
• A hydrate is an ionic
compound that contains
water molecules in its
structure
• An anhydrate is the
substance that remains
after the water is
• removed from a hydrate.
When a hydrate is heated
the water molecules are
driven off as steam, leaving
behind the water-free
anhydrate.
MgCO3 · 5H2O
Na2CO3 · 2H2O
BaCl2 · 2H2O
4 Steps to find formula of a
hydrate:
1.Determine mass of
water driven off:
2.Determine moles of
MgCO3 and water:
3.Find a whole number
molar ratio:
4.Write Formula
Hydrates Problems
#1: A 15.67 g
sample of a hydrate
of magnesium
carbonate was
heated, without
decomposing the
carbonate, to drive
off the water. The
mass was reduced to
7.58 g. What is the
formula of the
hydrate?
1) Determine mass of water driven off:
15.67 minus 7.58 = 8.09 g of water
2) Determine moles of MgCO3 and water:
MgCO3 ⇒ 7.58 g / 84.313 g/mol = 0.0899 mol
H2O ⇒ 8.09 g / 18.015 g/mol = 0.449 mol
3) Find a whole number molar ratio:
MgCO3 ⇒ 0.0899 mol / 0.0899 mol = 1
H2O ⇒ 0.449 mol / 0.0899 mol = 5
Formula:
MgCO3 · 5H2O
Hydrates Problems
#2: A hydrate of 1) Determine mass of water driven off:
4.31 minus 3.22 = 1.09 g of water
Na2CO3 has a
mass of 4.31 g
2) Determine moles of Na2CO3 and water:
Na2CO3 ⇒ 3.22 g / 105.988 g/mol = 0.0304 mol
before heating.
H2O ⇒ 1.09 g / 18.015 g/mol = 0.0605 mol
After heating,
the mass of the
3) Find a whole number molar ratio:
Na2CO3 ⇒ 0.0304 mol / 0.0304 mol = 1
anhydrous
H2O ⇒ 0.0605 mol / 0.0304 mol = 2
compound is
found to be 3.22 Formula:
g. Determine the
Na2CO3 · 2H2O
formula of the
hydrate.
Hydrate Starter Problems
1. A hydrate of magnesium sulfate has a mass of 13.52 g. This
sample is heated until no water remains. The MgSO4
anhydrate has a mass of 6.60 g. Find the formula and name
of the hydrate.
2. A sample of copper (II) sulfate hydrate has a mass of 3.97
g. After heating, the CuSO4 that remains has a mass of
2.54 g. Determine the correct formula and name of the
hydrate.
3. A sample of the hydrate of sodium carbonate has a mass
of 8.85 g. It loses 1.28 g when heated. Find the formula
and the name of the hydrate.
4. A 16.4 g sample of hydrated calcium sulfate is heated until
all the water is driven off. The calcium sulfate that
remains has a mass of 13.0 g. Find the formula of the
hydrate.
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