1. Isotopes and Atomic Mass
SCH 3U
Matter, Chemical Trends and Chemical Bonding
B3.2,3.1 l

2. What are Isotopes?
Two atoms 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.

3. Comparing isotopes of Magnesium
Magnesium has 3 isotopes, here is how they compare:
Magnesium-24
Magesium-25
Magnesium-26
12 protons
12 neutrons
13 neutrons
14 neutrons
12 electrons
Isotope mass 24u
Isotope mass
25u
26u

4. Comparing the isotopes of Magnesium
Similarities:
Same number of protons.
Same number of electrons.
Same appearance and chemical properties.
Differences:
Different number of neutrons.
Different atomic masses.

5. 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.

6. Isotopic Abundance
A sample of magnesium is a mixture of three isotopes, present as:
79% Mg-24
10% Mg-25
11% Mg-26

7. Average Atomic Mass and 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.

8. Average Atomic Mass
The average 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

9. Average Atomic Mass
The average atomic mass for carbon on the Periodic table is 12.01u.
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.

10. Calculating Average Atomic Mass
Use the equation:
AAM (u) = % 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

11. Sample Problem (p 167)
Silver has two naturally occuring isotopes, Ag-107 (m=106.9u) and Ag-109 (m=108.9u) with isotopic abundances of 51.8% and 48.2% respectively. Calculate the average atomic mass of silver.
AAM (u) = % ab1 x mass1 + % ab2 x mass2
AAM (u) = (0.518)x 106.9u + (0.482u) x 108.9u
AAM (u) = 55.4u u
AAM (u) = 107.9u

12. Sample Problem (p 169)
Boron exists as two naturally occuring isotopes: Boron-10 (10.01u) and Boron-11(11.01). Calculate the relative abundance of each isotope, if the average atomic mass of boron is 10.81u.
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).

13. Sample Problem (p 169 – continued)
Solution:
AAM (u) = % ab1 x mass1 + % ab2 x mass2
10.81 = x(10.01) + (1-x) (11.01)
10.81 = 10.01x – 11.01x
11.01x-10.01x =
x =
The abundance of Boron-10 is 20.0%
The abundance of Boron-11 is (1-x) or 80.0%