Chapter 10 Worksheet Examples
Worksheet:
MOLAR MASS
Example: Calculate the molar mass (gram
molecular weight) of a mole of iodine, I2.
Unit for molar mass
I = 126.9(2) = 253.8
g
mol
From the formula of the compound
From the mass off of the periodic table.
Example: Calculate the molar mass of a
mole of aluminum sulfate (Al2(SO4)3).
From the masses off of the periodic table.
From the formula of the compound
Al = 26.98 (2) = 53.96
S = 32.07 (3) = 96.21
Unit for molar mass
O = 16.00 (12) = 192.00
---------------------------------------g
add together = 342.17 mol
Worksheet:
MOLE AS A UNIT OF MASS
Example: What is the mass of 5.00 moles
of water?
X
90.1 g
H2O
18.02 g/mol
5.00 mol
Molar Mass Calculation
H = 1.01 (2) = 2.02
O = 16.00 (1) = 16.00
-----------------------------= 18.02 g/mol
Molar mass
From the problem
Example: What is the mass of 0.50 moles
of calcium carbonate?
X
50.045 g
CaCO3
100.09 g/mol
0.50 mol
From the problem
Molar Mass Calculation
Ca = 40.08 (1) = 40.08
C = 12.01 (1) = 12.01
O = 16.00 (3) = 48.00
-----------------------------Molar mass
= 100.09 g/mol
Example: How many moles of calcium chloride
are in 333 grams of calcium chloride
From the problem
333 g
CaCl2
110.98 g/mol
3.00 mol
Molar Mass Calculation
Ca = 40.08 (1) = 40.08
Cl = 35.45 (2) = 70.9
-----------------------------= 110.98 g/mol
Molar mass
÷
Worksheet:
AVOGADRO’S NUMBER
Example: How many molecules of water
are there in 3.00 moles of water?
X
1.806x1024 mcl
H2O
6.02x1023 mcl/mol
3.00 mol
The number of
atoms/molecules in one
mole is always 6.02x1023
From the problem
Example: How many moles of neon are
there in 2.408x1024 atoms of neon?
From the problem
2.408x1024 atm
Ne
6.02x1023 atm/mol
3.99 mol
The number of
atoms/molecules in one
mole is always 6.02x1023
÷
Worksheet:
MOLAR VOLUME OF A GAS
Example: What is the volume, in liters, of a
2.00 mole sample of methane (CH4) at STP?
X
44.8 L
CH4
22.4 L/mol
2.00 mol
From the problem
The number of liters in one
mole is always 22.4
Example: How many moles of ethane
(C2H6) are there in 5.60 liters of ethane?
From the problem
5.6 L
C2H6
22.4 L/mol
0.25 mol
The number of liters in one
mole is always 22.4
÷
Worksheet:
MIXED MOLE PROBLEMS
Molar mass is g/mol and you find
grams from the periodic table.
We’ve done problems of
one step. Moles to a unit
or unit to moles. Now we
can do multi step
problems using this
picture to help us see
where we need to go
next.
This picture is a summary
of all of the problems we
have done up to this
point along with helpful
hints to units and
numbers. The better you
understand this the
easier mole conversions
will be.
Example: What would be the volume in
liters of 40.36 grams of neon at STP?
From the problem
40.36 g
Ne
20.18 g/mol
2.00 mol
From the
periodic table
The number of
liters in one mole
is always 22.4
We’re not to
liters yet…
44.8 L
Ne
22.4 L/mol
2.00 mol
X
÷
Example: How many molecules would there
be in 56 liters of carbon dioxide at STP?
From the problem
The number of
liters in one mole
is always 22.4
56 L
CO2
22.4 L/mol
2.5 mol
1.505x1024 mcl
CO2
6.02x1023 mcl/mol
2.5 mol
We’re not to
molecules yet…
The number of
atoms/molecules in one
mole is always 6.02x1023
X
÷
Worksheet:
PERCENT COMPOSITION &
FORMULAS
Example: Calculate the percent composition
of sodium hydrogen carbonate (NaHCO3)?
Step 1 – find molar mass of the
compound.
Step 2 – divide mass of each
element by molar mass.
Step 3 – multiply answers by 100%.
Molar Mass Calculation
Na = 22.99 (1) = 22.99
H = 1.01 (1) = 1.01
C = 12.01 (1) = 12.01
O = 16.00(3) = 48.00
----------------------------------------------
= 84.01 g/mol
22.99
 100%
84.01
 27.37% Na
1.01
100%
84.01
 1.20% H
12.01
100%
84.01
 14.30%C
48.00
 100%
84.01
 57.14%O
Example: Which pair of molecules has the
same empirical formula?
a) C2H4O2
CH2O
c) NaCrO4
b) C6H12O6
Molecular formula
CH2O
Empirical formula
d) Na2Cr2O7
These two can not be “reduced”.
An empirical formula is similar to a reduced fraction.
Example: Calculate the empirical formula for a
compound with 67.6% Hg, 10.8% S, 21.6% O?
Step 1 – change percents to grams (assume we have 100 gram sample then percents are the
number of grams).
Step 2 – change the grams to moles.
Step 3 – divide all answers by the smallest number (of the answers).
Step 4
Step 4 – the results become the subscripts for the formula.
Step 3
10.8 g
21.6 g
Hg
S
O
200.59 g
÷
Step 2
67.6 g
mol
0.34mol
1
0.34mol
32.07 g
mol
0.34mol
1
0.34mol
÷
Step 1
HgSO4
16.00 g
mol
1.35mol
4
0.34mol
Example: Find the molecular formula of
ethylene glycol, which is used as antifreeze.
The molecular mass is 62g/mol and the
empirical formula is CH3O?
Step 1 – Find the empirical formula mass (EFM)
Step 2 – Divide the molar mass (from problem) by the EFM to get the multiplier.
Step 3 – Use the multiplier to determine the subscripts by multiplying each subscript by the
multiplier.
Step 1 EFM Calculation
C = 12.01 (1) = 12.01
H = 1.01 (3) = 3.03
O = 16.00(1) = 16.00
----------------------------------------
=31.04 g/mol
Step 3
Step 2
62
2
32.01
MULTIPLIER
C2H6O2