To do these problems you need some information: the exact atomic weight for each naturally-occuring stable isotope and its percent abundance. These values can be looked up in a standard reference book such as the "Handbook of Chemistry and Physics."
This problem can also be reversed. Study the tutorial below and then look at the problems done in the reverse direction.
Example #1: Carbon
| mass number | exact weight | percent abundance |
| 12 | 12.000000 | 98.90 |
| 13 | 13.003355 | 1.10 |
To calculate the average atomic weight, each exact atomic weight is multiplied by its percent abundance (expressed as a decimal). Then, add the results together and round off to an appropriate number of significant figures.
This is the solution for carbon:
(12.000000) (0.9890) + (13.003355) (0.0110) = 12.011
Example #2: Nitrogen
| mass number | exact weight | percent abundance |
| 14 | 14.003074 | 99.63 |
| 15 | 15.000108 | 0.37 |
This is the solution for nitrogen:
(14.003074) (0.9963) + (15.000108) (0.0037) = 14.007
Video: How to Calculate an Average Atomic Weight.
| Example #3: Chlorine | Example #4: Silicon | |||||
| mass number | exact weight | percent abundance | mass number | exact weight | percent abundance | |
| 35 | 34.968852 | 75.77 | 28 | 27.976927 | 92.23 | |
| 37 | 36.965903 | 24.23 | 29 | 28.976495 | 4.67 | |
| 30 | 29.973770 | 3.10 | ||||
| The answer for chlorine: 35.453 | The answer for silicon: 28.086 | |||||
This type of calculation can be done in reverse, where the isotopic abundances can be calculated knowing the average atomic weight. Go to tutorial on reverse direction.
Calculate the average atomic weight for:
1) magnesium
| mass number | exact weight | percent abundance |
| 24 | 23.985042 | 78.99 |
| 25 | 24.985837 | 10.00 |
| 26 | 25.982593 | 11.01 |
2) molybdenum
| mass number | exact weight | percent abundance |
| 92 | 91.906808 | 14.84 |
| 94 | 93.905085 | 9.25 |
| 95 | 94.905840 | 15.92 |
| 96 | 95.904678 | 16.68 |
| 97 | 96.906020 | 9.55 |
| 98 | 97.905406 | 24.13 |
| 100 | 99.907477 | 9.63 |
3) tin (this one is optional!! Suggestion: set it up as a spreadsheet, take it into class and impress your teacher.)
| mass number | exact weight | percent abundance |
| 112 | 111.904826 | 0.97 |
| 114 | 113.902784 | 0.65 |
| 115 | 114.903348 | 0.36 |
| 116 | 115.901747 | 14.53 |
| 117 | 116.902956 | 7.68 |
| 118 | 117.901609 | 24.22 |
| 119 | 118.903310 | 8.58 |
| 120 | 119.902200 | 32.59 |
| 122 | 121.903440 | 4.63 |
| 124 | 123.905274 | 5.79 |
The answers? Look on a periodic table!! Remember that the above is the method by which the average atomic weight for the element is computed. No one single atom of the element has the given atomic weight because the atomic weight of the element is an average, specifically called a "weighted" average.