Isotopes & Atomic Mass

Isotopes & Atomic Mass

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  Isotopes and AtomicMass
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  What are Isotopes? Twoatoms are isotopes if they have the same number of protons, but they have different numbers of neutrons. This means that: Isotopes are atoms of the same element. Isotopes have different atomic masses. Isotopes have different number of neutrons in their nuclei.
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  Comparing isotopes ofMagnesium  Magnesium has 3 isotopes, 24Mg, 25Mg, and 26Mg. Here is how they compare: Magnesium-24 Magesium-25 Magnesium-26 12 protons 12 protons 12 protons 12 neutrons 13 neutrons 14 neutrons 12 electrons 12 electrons 12 electrons Isotope mass 24 amu Isotope mass 25 amu Isotope mass 26 amu
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  Comparing the isotopesof Magnesium Similarities: Same number of protons. Same number of electrons. Same appearance and chemical properties. Differences: Different number of neutrons. Different atomic masses.
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  Isotopic Abundance (%Abundance) A sample of magnesium is a mixture of the three isotopes of magnesium. Each isotope is a fraction of the mixture, and has its own isotopic abundance (expressed as a percentage of the whole). The isotopic abundance is fixed so that every sample of the element (in the universe) has the same proportions of the isotopes.
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  Isotopic Abundance A sampleof magnesium is a mixture of three isotopes, present as: 79% 24Mg (or Mg-24) 10% 25Mg (or Mg-25) 11% 26Mg (or Mg-26)
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  Average Atomic Massand Isotope Abundance  The Average Atomic Mass seen on the Periodic Table is a weighted average of all of the isotope masses.  The weighted average takes into account the isotope masses and their percent abundances.  In a weighted average calculation, the isotope with the greatest % abundance has the biggest influence on the average atomic mass.
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  Average Atomic Mass Theaverage atomic mass is a weighted average of all the isotope masses for a particular element. When you calculate average atomic mass, you need three pieces of information: The number of isotopes The masses of each isotope The % abundance of each isotope
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  Average Atomic Mass Theaverage atomic mass for carbon on the Periodic table is 12.01 amu. This means:  Carbon has more than one isotope;  One of carbon’s isotopes has a mass of 12, another is greater than 12;  The most abundant isotope is Carbon-12.
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  Calculating Average AtomicMass Use the equation: AAM (amu) = %ab1 x mass1 + %ab2 x mass2 + … %abn x massn Where: AAM = Average Atomic Mass %ab1 = Percent abundance of isotope 1 mass1 = Mass of isotope 1
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  Sample Problem 1 Silver has two naturally occurring isotopes, Ag- 107 (m=106.9 amu) and Ag-109 (m=108.9 amu) with isotopic abundances of 51.8% and 48.2% respectively. Calculate the average atomic mass of silver. AAM = % ab1 x mass1 + % ab2 x mass2 AAM = (0.518) x 106.9 + (0.482) x 108.9 AAM = 55.4 + 52.5 AAM = 107.9 amu
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  Sample Problem 2 Boron exists as two naturally occurring isotopes: 10Boron (10.01 amu) and 11Boron (11.01 amu). Calculate the relative abundance of each isotope, if the average atomic mass of boron is 10.81 amu. To solve: The percent abundance of all the isotopes should add up to 1. The % abundance of boron-10 is (x), and the % abundance of boron-11 is (1-x).
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  Sample Problem 2(continued) Solution: AAM = %ab1 x mass1 + %ab2 x mass2 10.81 = x(10.01) + (1-x) (11.01) 10.81 = 10.01x + 11.01 – 11.01x 11.01x-10.01x = 11.01 -10.81 x = 0.2000 The abundance of 10Boron is 20.0% The abundance of 11Boron is (1-x) or 80.0%