Transcript Isotopes

Isotopes
Atoms with the same number of
protons but different numbers of
neutrons
Ex) Carbon 12 vs. Carbon 14
These atoms have a different mass
Chemically alike because still have the
same number of protons
Isotopes of Hydrogen
Hydrogen -1 simply called hydrogen
Hydrogen - 2 called deuterium
Hydrogen - 3 called tritium
Development of AMUs
Atomic Mass Units (AMUs)
Protons have a mass of 1 amu
Neutrona have a mass of 1 amu
Electrons have a mass of 0 amu
Atomic Mass
The weighted average mass of the isotopes in
a naturally occurring sample of the element
Don’t confuse with “mass number”
To calculate atomic mass you need 3 pieces
of information
1. The number of stable isotopes
2.The mass of each isotope
3.The natural percent abundance of each
isotope
Atomic Mass
Example Problem - Calculate the atomic mass
for element X. One isotope has a mass of 10
amus (10X) and is 20% abundant. The other
has a mass number of 11 amus (11X) and an
abundance of 80%.
To solve: Multiply the mass number times the
abundance than add them together.
Atomic Mass
10 x 0.20 = 2.0
11 x 0.80 = 8.8
Add 2.0 + 8.8 = 10.8
The atomic mass of element X is 10.8
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Atomic Mass
Your turn. Solve:
What is the atomic mass of Element Z?
The isotopes are 16Z, 17Z, 18Z; with percent
abundances of 99.759, 0.037, 0.204.
Atomic Mass
Answer
16 x 0.99759 = 15.961
17 x 0.00037 = 0.0063
18 x 0.00204 = 0.0367
15.961 + 0.0063 + 0.0367 = 16.004
Tha atomic mass of element Z is 16.004
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