Ch 8: Chemical Composition
8.1
Counting by Weighing
1. How are atoms counted? Practically speaking, they are far too small to count directly. Therefore,
they must be counted indirectly.
2. Scenario: You work in a candy store. A customer comes in and wants to buy 1000 jellybeans. It is
not practical to count out 1000 beans. How do you sell the correct quantity?
a. First you need to know the average mass of a jelly bean. The masses of typical jelly beans are
added together and an average is computed.
b. Let’s say the average mass of a jelly bean is 5.0 grams. How can a customer quickly be given
1000 beans?
--You would simply weigh out 5000 grams of beans
5.0 grams/bean X 1000 beans = 5000 grams.
You never actually count out the beans; you count them by weighing.
3. We are not really talking about jellybeans here, but atoms. We count atoms by weighing them. All
we need to know is the average mass of an atom.
8.2
Atomic Masses: Counting by Weighing
1. Individual atoms weight far too little to use usual mass terms such as kilogram or even gram. A
typical atom only weighs about 5.00 x 10-23 grams.
2. To avoid using inconvenient numbers such as 10-23, a smaller unit of mass called the atomic mass unit
(amu) is used .
1 amu = 1.66 x 10-24 grams
3. To count atoms using their mass we have to know the average mass of an atom of a particular
element. As you recall from Ch 4, the average atomic mass of an element is a weighted average of all
isotope masses of the element.
Ex: Carbon has an average atomic mass of 12.01 amu (atomic mass units). It represents the weighted
average of carbon isotopes 12C , 13C , and 14C .
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(Remember, carbon-12 is by far the most common isotope of carbon.)
Therefore, even though each isotope of carbon has a slightly different mass, any sample of carbon in
nature can be treated as though each carbon atom in the sample has a mass of 12.01 amu (to 4 sig. figs.)
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The weighted average mass for all elements is listed on the periodic table under their symbols.
average atomic mass
What applies to carbon applies to all elements on the periodic table:
What is the average atomic mass in amu of an atom of nitrogen?
Therefore, any atom of nitrogen can be treated as though it has a mass of ________________ amu
Sample Problems:
Ex: What is the mass of 5000 carbon atoms in amu?
5000 atoms 
12 .01 amu
 60050 amu
C atom
Ex: What is the mass in grams of 5000 carbon atoms?
12 .01 amu 1.66  10 24 g
5000 atoms 

 9.97  10  20 grams
C atom
amu
Ex: How many sodium atoms are present in a sample that has a mass of 1172.49 amu?
What is the average atomic mass of Na?
1172 .49 amu 
22.99 amu from the periodic table
1 Na atom
 51 .00 Na atoms
22 .99 amu
Ex: How many iron atoms are present in an iron sample weighing 5.00 x 10-12 grams?
What is the average atomic mass of Fe? 55.85 amu
5.00  1012 g 
1 amu
1 Fe atom

 5.39  1010 atoms
 24
1.66  10 g 55.85 amu
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8.3
The Mole
Most chemical reactions are carried out with masses of materials measured in grams, not atomic mass
units.
Any sample of an element that has a mass in grams equal to its average atomic mass contains one mole of
atoms.
What is a mole? It is simply a number of something. A very, very big number, actually.
One mole of something contains 6.022 x 1023 units of that substance. This number is known as
Avogadro’s number.
A mole of marbles = 6.022 x 1023 marbles
A mole of eggs = 6.022 x 1023 eggs
A mole of atoms = 6.022 x 1023 atoms
A mole of molecules = 6.022 x 1023 molecules
So, a mole is simply a ___________________ of something.
In chemistry, a mole (Avogadro’s number) is defined as the number of atoms in a 12.01 gram sample
of carbon.
(12.01 amu/atom) (1.66 x 10-24 gram/amu) (6.022 x 1023 atoms) = 12.01 grams
In fact, if we weigh out samples of each element such that the sample has a mass equal to that of the
element’s average atomic mass in grams, these samples all contain the same number of atoms, that being
6.022 x 1023 atoms.
(63.55 amu/Cu atom) (1.66 x 10-24 gram/amu) (6.022 x 1023 atoms) = 63.55 grams Cu
Examples: see Table 8.2
A mole is simply an easy way of keeping track of large numbers of atoms.
The molar mass of an element is equal to its average atomic mass in grams. One molar mass of any
element contains 6.022 x 1023 atoms of that element
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