How to calculate the molar masses of chemical
compounds
What is molar mass?
Molar mass is the weight of one mole (or 6.02 x 1023 molecules) of any chemical
compounds. Molar masses of common chemical compounds that you might find in the
chemistry laboratory can range between 18 grams/mole for compounds like water to
hundreds of grams per mole for more complex chemical compounds.
The lightest possible chemical that one can have under normal conditions is hydrogen gas, or
H2 . There is no limit to how heavy a chemical compound can be - it is not uncommon for
macromolecules (large organic or bioorganic compounds such as DNA) to weigh thousands
of grams per mole.
How can I find the molar mass of an element?
The molar mass of elements is found by looking at the atomic mass of the element on the
periodic table. For example, if you want to finqJ the molar mass of carbon, you would find the
atomic mass of carbon on the periodic table, and this is equal to the molar mass in grams per
mole. So, in our example, carbon has a molar mass of 12.01 grams per mole.
There are a few exceptions to this rule. In some cases, the element is usually found in a
different form than just one unbonded atom. In the case of hydrogen, nitrogen, oxygen,
fluorine, chlorine, bromine, and iodine, the element is diatomic, meaning that each molecule
of the element has two atoms of that element stuck together. As a result, the formula of
hydrogen is Hz, nitrogen is N 2 , etc...
This gets weirder for a couple of cases... phosphorus is normally found in clumps of four
atoms, P4' and sulfur is found in clumps of eight atoms, or S8·
Still, aside from the exceptions above, all elements have the same molar mass as the atomic
masses on the periodic table.
How can I find the molar mass of a chemical compound?
For any chemical compound that's not an element, we need to find the molar mass from the
chemical formula. To do this, we need to remember a few rules:
1. Molar masses of chemical compounds are equal to the sums of the molar masses of
all the atoms in one molecule of that compound. If we have a chemical compound like
NaCI, the molar mass will be equal to the molar mass of one atom of sodium plus the molar
mass of one atom of chlorine. If we write this as a calculation, it looks like this:
(1 atom x 23 grams/mole Na) + (1 atom x 35.5 grams/mole CI)
= 58.5 grams/mole NaCI
2. If you have a subscript in a chemical formula, then you multiply the number of
atoms of anything next to that subscript by the number of the sUbscript. For most
compounds, this is easy. For example, in iron (II) chloride, or FeCI2, you have one atom of
iron and two-atems of chlorine. The molar mass will be equal to (1 atom x 56 grams/mole Fe)
+ (2 atoms x 35.5 grams/mole of chlorine):: 127 grams/mole of iron (II) chloride.
For other compounds, this might get a little bit more complicated,' For example, take the
example of zinc nitrate, or Zn{N0 3)2' In this compound, we have one atom of zinc, two atoms
of nitrogen (one atom inside the brackets multiplied by the subscript two) and six atoms of
oxygen (three atoms in the brackets multiplied by the subscript two). The molar mass of zinc
nitrate will be equal to (1 atom x 65 grams/mole of zinc) + (two atoms x 14 grams/mole of
nitrogen) + (SIX atoms x 16 grams/mole of oxygen) = 189 grams/mole of zinc nitrate.
For all other compounds, the general idea is the same, Basically, you should know how to
find the molar masses of any chemical compound now. In the next and final section, I'll give
you some practice problems, followed by a solution key...
.,
Some sample problems:
These are the kind. of molar mass calculation probems I might ask you on a quiz. The
solutions are given at the end.
Give the molar masses of the following compounds:
1. sodium fluoride
Ivh F
2.. potassium hydroxide
3. copper (I) chloride
~+
1(04-
Cu.el
4. manganese (IV) oxide
5. calcium sulfate ,
"2.-
., '1
~'-\+
l.j
1.. j ("""
-I- \
10 "'- I ":.
S;,
M() 2.. 0 L/-
Co,.. So 4-
6. magnesium phosphate
1'1 ..
VI 0 +-
~
I
-S ~ , r""l " I
"l'fj I ..... J
'S$CL)
+-
sl.. ~ \~ ( u,')
j\\f) 3 (PDq.) L
\
'.to (41 -;; fI '1 .) / ...... , I
7,
13 ""; / .... "
,7.­
.L. Y
C:~> +-) \ ~ ( I ~,g')
~ L. C:.
.--,
Chemistry 7-1: Introduction to the Mole and Molar Mass
Worksheet: Mole Problems
Name.
Period
Part 1: Molar Mass
Use the periodic table to find the moiar maSSeS of the following.
HCI
K,C0 3
Ca(OH),
Part 2: Mole Conversions
Work each of the following problems. SHOW ALL WORK.
1. How many atoms are in 6.2 moles of aluminum?
2. Convert 5.3 x 10 25 molecules of C02 to moles.
3. How many formula units of sodium acetate are in 0.87 moles of
sodium acetate?
4. Convert 3.55 moles NaCI to formula units.
page 7-6
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All
R.itlUJl,..=ed, OPTe ~GO},'
_
_
5. Convert 3.00 mole As,S, to grams.
6. How many moles are represented by 11.5 9 of C,H50H?
7. What is the mass of 9.30 moles of SiH.?
8. Convert 8.00 x 10'0 molecules of Hz to moles.
9. How many atoms of tin are found in 3.50 moles of tin?
10. How many grams of tin are found in 3.50 moles of tin?
Bonus: How many atoms of hydrogen are found in 12.6 moles of water?
page 7-7
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Chemistry
. arK, ,arne
. ar t'IC Ie \\1\[ a le/P
N'
_
A sample of KI03 is found to contain 1.24 x 10 18 formula
units of potassium iodate. Determine the number of moles.
1)
2) You have 0.35 moles of phosphorus pentachloride, PCI5. \Vhat
mass is this?
3) How many moles would be contained in a sample of 8.045 g
of titanium (IV) fluoride, TiF4?
4) A lab assistant gives you a sample labeled as 1. 58 moles of
silicon dioxide, Si02. How many molecules are in the sample?
5)
You isolate 14.3 g of copper metal in lab. How many atoms
of copper are in this sample?
6) Consider a sample containing 4.25 x 10 10 formula units of
iron (III) oxide, Fe203. What mass would this sample have?
7) You have a sample which contains 8.78 x 1015 molecules of
solid carbon dioxide, C02 (dry ice). What mass is this sample?
8) A sample of Ca(N03)2 is found to have a mass of 148.9 grams.
How many formula units of calcium nitrate are in this sample?
9)
For the example in number 8, determine how many nitrogen
atoms would be in the same sample.
·--·10)·-Still using the example in number 8, determine the number of
oxygen atoms in the sample.
Hints: molar mass is grams per mole AND Avogadro's number gives
us 6.02 x 10 23 particles in 1 mole of any substance.
Percent Composition Handout
When we calculate percent composition, we're really trying to determine how
much of the compound's weight is due to each of the elements present. For
example, if I were to ask you to find the percent composition of oxygen in water,
I'm really asking you to find the percent of the total weight of water that's due to
oxygen atoms.
To solve this problem, use the following procedure:
1)
Assume you have one mole of the substance you're analyzing,
This assumption makes the problem easier to solve, In our case, we'll just
assume we have one mole of water,
2)
For each element, multiply the atomic weight of that element by the
number of atoms of that element in each molecule, Keep these
numbers separate from each other for now,
For hydrogen, multiply the atomic weight of hydrogen (1 g/mol) by the
number of atoms of hydrogen per water molecule (2) to find 2 glmol.
For oxygen, multiply the atomic weight of oxygen (16 g/mol) by the
number of atoms of oxygen per water molecule (1) to find 16 g/mol.
3)
Add these numbers together, This gives you the molar mass of the
compound,
2 g/mol + 16 g/mol = 18 glmole
4)
Divide the mass of the element you're looking for by the molar mass
of the compound, Multiply this number by 100 to find the percent
composition.
The mass of oxygen from step 2 is 16 grams/mole, and the moiar mass of
water from step 3 was 18 grams/mole. Dividing these numbers, we get:
(16 g/mol) / (18 g/mol) = 0.89
When we multiply 0.89 by 100, we find that the percent composition of
oxygen in water is 89%,
--_
..
85
Percent Composition Worksheet
dt sUlfur In sulfur dioxide (SO,)?
1)
What is the percent composition
2)·
What is the percent composition o(r"hon)n methane (CH,)?
3)
What is the percent composition QI'Oxygen.in lithium hydroxide (LiOH)?
4)
What is the percent composition of hydrogen in sulfuric acid (H,S04)?
5)
What is the percent composition of nitrogen in ammonium nitrate
(NH 4 NO,)?
89
Name
---------------------Empirical Formulas
I have provided you with copies of an online tutorial on Empirical Formulas. Read
through it and see if you can figure this concept out for yourself. I think you can!
(It's a very good website. Here's the address if you're interested:
http://dbhs.wvtlsdkl2.ca.lIs/webdocs/ChemTeamIndex. html Choose the section on the
mole and then click the section called Empirical Formulas. Of course there are other
topics on there too.)
I. What is the definition ofempirical formula?
2. Contrast the defmition of empirical formula with the definition of molecular
formula.
3. Give 2 examples ofempirical and molecular formulas.
4. Can an empirical formula and a molecular formula be the same?
5. Write the rhyme you can remember for converting % composition to an empirical
formula.
6. Using the sample problems within the reading, do the Empirical Formula Practice
Problems 1-5. Show your work. You can do them on this sheet.
Name
_
Date
_
Empirical and Molecular Formula Practice
1. Determine the empirical formula for a compound that is 79. 9'Yo
Copper and 20.1 'Yo Sulfur.
2. If 43.2 grams of carbon combine with 115.8 grams of oxygen
when burned, what is the empirical formula of the compound that
is formed?
3. Calculate the empirical formula for a compound which is 40.2'>'0
potassium, 26.9 '>'0 Chromium, and 32.9'>'0 Oxygen.
4. A compound is composed of 7.20 grams of carbon 1.20 grams of
hydrogen, and 9.60 grams of oxygen. The molar mass of the
compound is 180 grams. What are the empirical and molecular
formulas for this compound?
.­
5. In a 32.0 gram sample of Hydrazine (a chemical used to treat
waste water) there are found to be 28.0 grams of Nitrogen, and
4.0 grams of Hydrogen. The molar mass of hydrazine is 32.0
grams per mole. What is the empirical and molecular formulas
for hydrazine?
Mole Problems
Page 1 of 2
Guide Sheet for Moles Problems
I. Calculating Molar Mass
1. mUltiply atomic mass of each element by number of
atoms of that element in the formula (shown by the
sUbscript)
2. find the sum of all the atomic masses --this is formula
mass (unit is a.m.u.)
3. express formula mass in grams (unit is gjmol). This is
the Molar Mass.
II. Calculating % Composition (from
formula)
• calculate formula mass
• divide the total atomic mass of each element by the
formula mass and multiply by 100
III. Calculating % Composition (from
masses of each element)
1. divide the mass of each element by the total mass of
the:compound and multiply by 100
IV. Calculating Empirical Formula (from %
Composition)
1. convert % of each element to grams based on 100 grams of the compound
2. multiply grams of each element by ljmolar mass that element
3. compare ratio of moles of each element and divide each by the smallest
4. if result in step 3 gives a ratio with decimal equivalent to 1/4, 1/3,1/2, 2.3, 3/4 instead of
whole numbers, convert to the fraction and multiply all ratios by the denominator or the
fraction
V. Calculating Empirical Formula (from experimentally determined
masses)
1. multiply the mass of each element (in grams) by l/molar mass of that eiement
2. continue with steps 3 & 4 from IV above.
Example
VI. Finding Molecular Formulas (when molar mass is known)
1.
2.
3.
4.
calculate the empirical formula
use the equation: (empirical formula mass)x = molar mass
find value for x:
x = molar mass/empirical formula mass
multiply each subscript in empirical formula by value for x
http://www.rhem.vt.edu/RVGS/ACT/notes/Stuoy_Guide-Moles]roblems.html
10/18/200S
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----========­
Mole Relationships
,1
I. Give the atomic or molecular mass for each of the following.
a) NaOH
b) HN03
d) S8
e) Cal(P04)Z
2, Calculate the mass of I mole for each of the following.
a) CaCO l
b) SiOz
d) C12HZ2011
e) CaS04
c) NaHCO l
t) Pb(OH)2
3. How many GRAMS of each of the CONSTITUENT elements are in I MOLE of the following?
b) Fe203
c) CalPz
a) C~
4, Calculate the number of moles for each of the following.
a) 80,0 g NaOH
b) 24.5 g H 2S04
d) 4.19 g NH4CN
e) 107 g Ba(ClOl)z
c) 4.00 g Oz
f) 11.6 g NaCI
5, Calculate the percent of:
a) IRON in:
i) FeCOl
b) COPPER in:
i) CU20
c) NITROGEN in:
i) HN0 3
iii) Fel04
iii) CUCOl
iii) (N~)2S04
ii) Fe203
ii) CuFeSz
ii)~N03
6, Determine the % composition of each element for each of the following.
b) CaZPZ07
c) KZC03
a) AgzCr04
7. An iron ore is analyzed and found to be 70% iron and 30% oxygen by mass:
Find the empirical formula,
'
8. 88 g of a hydrocarbon is analyzed and found to contain 72 g of carbon and 16 g of hydrogen.
a) Calculate the % composition of each element in the compound.
b) Find the empirical formula.
'
c) 44 g of the hydrocarbon, in a gaseous state, occupies a volume of22.4 L at S.T.P..
What is the molecular formula?
9, The anaesthetic, chloroform, has a molar mass of 119.4 glmol. Analysis reveals it to consist of;
10.05 % C, 0.84 % H, and 89.10 % CI. What is the molecular formula?
10. Acetylene has molar mass of26 g/mo\. It has the following % composition: 92,25 % C and 7.74 % H,
What is the molecular formula?
11. A compound consists of 20.2 % P, 10.4 % 0, and 69.4 % C\.
a) What is its empirical formula?
b) The sample was subsequently vaporized, and, volume for volume, was found to be 76.75 times as
heavy as a sample of hydrogen gas. (ie, its molar mass = 76.75 X molar mass of Hz)
What is the molecular formula of this compound?
_
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_
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_
The Mole
Section 11.1
Measuring Matter
In your textbook, read about counting particles.
In Column B, rank the quantities from Column A from smallest to largest.
Column A
Column B
0.5 mol
1.
200
2.
5
3.
6000000 000
4.
6.02 X 1023
5.
dozen
6.
four moles
7.
gross
8.
pair
9.
10.
ream
In your textbook, read about converting moles to particles and particles to moles.
In the hoxes provided, write the conversion factor that correctly completes each
problem.
11. 1.20 mol Cu X
12.9.25 X
10~~ molecules CH, X
13. 1.54 X
1O~6 atoms Xe X
14. 3.01 mol F, X
Study Guide for Content Mastery
l
1
I
l
1
= 7.22 X 1()23 Cu atoms
1 = 1.54 X 10- 1 mol CHI
!
= 2.56 X 102 mol Xe
=
1.81 X 1024 molecules F2
Chemistry: Matter and Change' Chapter 11
61