Silver has two naturally occurring isotopes. Ag- 107 with a mass of 106.905 amu and a natural abundance of 51.84 % and Ag-109. How do you use the atomic mass of silver listed in the periodic table to determine the mass of Ag-109?

1 Answer
Sep 12, 2016

Using my (not very precise) periodic table, I get the mass of Ag as 107.87.

This must be a combination of these two isotopes. So the % of the other isotope must be 48.16%.
To work with these percentages in an equation I'll use them as decimals.

So now I can start putting them into an equation:
107.87=(106.9050.5184)+(x0.4816)
Because both of the isotope masses, multiplied by their percentage, would give us the overall average mass.
107.87=55.419552(x0.4816)
Worked out the brackets that we can work out, and then rearrange to give:
107.8755.419552=(x0.4816)
Again rearrange:
52.450448=(x0.4816)
And final rearranging:
x=52.4504480.4816
x=108.909 (using same sig figs as given in question)

You probably want to work through that again using your periodic table value for Ag.