# Isotope Relationships Flashcard Example #1282

Elaborate on how the isotopes and their relative abundances affect the average atomic mass of an element.
Isotope mass both varies in neutron numbers and is weighted by its relative abundance to give the average atomic mass.
note-Isotope mass both varies in neutron numbers and is weighted by its relative abundance to give the average atomic mass.

Isotopes are atoms with the same atomic number, or number of protons and electrons, but different masses. The different masses are due to different numbers of neutrons. The relative natural abundance of each isotope shows the predominance of each. When calculating the average atomic mass of the element, the mass of of each isotope is weighted by its relative abundance.

Elaborate on how the relative abundances of naturally occurring isotopes relate and can be used to determine the average atomic mass of an element.
Isotope masses are weighted by their relative abundances and then add to give the average atomic mass.
note-Isotope mass both varies in neutron numbers and is weighted by its relative abundance to give the average atomic mass.

Isotopes are atoms with the same atomic number, or number of protons and electrons, but different masses. The different masses are due to different numbers of neutrons. The relative natural abundance of each isotope shows the predominance of each. When calculating the average atomic mass of the element, the mass of of each isotope is weighted by its relative abundance.

Boron has two naturally-occurring isotopes. Boron-10 has an abundance of 19.8% and actual mass of 10.013 amu, and boron-11 has an abundance of 80.2% and actual mass of 11.009 amu. What is the average atomic mass for all isotopes of boron?
The correct answer is 10.812 amu. Since the two isotopes are found in different abundance, you cannot simply add them up and divide by 2.
First convert the percentages to a decimal, by dividing each by 100.
The correct setup is then:
(10.103 x .198) + (11.009 x .802) = 10.812
The naturally occurring isotopes of magnesium are magnesium-24. magnesium-25, and magnesium-26. Magnesium-24 has an abundance of 78.994% and a mass of 23.985 amu. Magnesium-25 has an abundance of 10.001% and a mass of 24.986 amu. Magnesium-26 has an abundance of 11.013% and a mass of 25.983 amu. Calculate the average atomic mass of magnesium.
24.307 amu
note-To calculate the average atomic mass first convert the abundances from percents to decimals. Then multiply the abundance times the mass for each isotope and sum these.
average atomic mass=(0.78994)(23.985 amu) + (0.10001)(24.986 amu) + (0.11013)*(25.983 amu)= 0.1248675 amu + 50.81664 amu= 24.30706855 amu= 24.307 amu
The naturally occurring isotopes of potassium are potassium-39, potassium-40, and potassium-41. Potassium-39 has an abundance of 93.258% and a mass of 38.964 amu. Potassium-40 has an abundance of 0.011710% and a mass of 39.964 amu. Potassium-41 has an abundance of 6.7302% and a mass of 40.962 amu. Calculate the average atomic mass of potassium.
39.099 amu
note-To calculate the average atomic mass first convert the abundances from percents to decimals. Then multiply the abundance times the mass for each isotope and sum these.
average atomic mass=(0.93258)(38.964 amu) + (0.00011710)(39.964 amu) + (0.067302)*(40.962 amu)= 0.1248675 amu + 50.81664 amu= 39.0985514284 amu= 39.099 amu