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Isotopes

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- 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

- Hydrogen -1 simply called hydrogen
- Hydrogen - 2 called deuterium
- Hydrogen - 3 called tritium

- 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

- 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

- 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.

- 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 amus

- 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.

- 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 amus