1. Isotopes Atoms of the same element that different mass numbers
(which means they have different numbers of neutrons)
They can be written two ways:
element name - mass number
Symbol
mass #
atomic #
carbon - 12
6 p+ and 6 n0 C 12 6 C 13 6
carbon - 13
6 p+ and 7 n0 C 14 6
carbon - 14
6 p+ and 8 n0
Remember:
mass number = p+ + n0 X 35 17
(a) X 35 18
(b) X 37 17
(c) X 38 19
(d)
Which of the following are
isotopes of the same element?

2. Atomic Mass x 75.77% = x 24.23% = 896.51 + = = 35 amu 2651.95 37 amu
A weighted average of the masses of all the isotopes of an element.
Step 1: Multiply the mass by its percent for each isotope
Step 2: Add them together
Step 3: Divide by 100
Use "amu" for its unit of measure
Example: Chlorine has two naturally occurring isotopes, chlorine-35 and chlorine-37. Their relative abundances are 75.77% and 24.23%, respectively. Calculate the atomic mass for chlorine.
35 amu x
75.77% =
37 amu x
24.23% =
896.51 + = =
35.5 amu
100

3. Atomic Mass x 19.91% = x 80.09% = + = = 10.012 amu 199.34 11.009 amu
Example: Element X has two natural isotopes. The isotope with a 
mass of amu has a relative abundance of 19.91%. The 
isotope with a mass of amu has a relative abundance of 
80.09%. Calculate the atomic mass of this element.
amu x
19.91% =
199.34
amu x
80.09% =
881.71 + = =
10.8 amu
100