We’re being asked to **determine the atomic mass of ^{65}X. **

We can use the following equation:

$\overline{){\mathbf{Atomic}}{\mathbf{}}{\mathbf{Mass}}{\mathbf{=}}{{\mathbf{[}}{\mathbf{mass}}{\mathbf{\times}}{\mathbf{f}}{\mathbf{.}}{\mathbf{a}}{\mathbf{.}}{\mathbf{]}}}_{\mathbf{isotope}\mathbf{}\mathbf{1}}{\mathbf{+}}{{\mathbf{[}}{\mathbf{mass}}{\mathbf{\times}}{\mathbf{f}}{\mathbf{.}}{\mathbf{a}}{\mathbf{.}}{\mathbf{]}}}_{\mathbf{isotope}\mathbf{}\mathbf{2}}}$

where:

atomic mass = average atomic mass of the element

f.a. = fractional abundance of the isotope.

To get ** f.a.**, we simply need to divide the given percent abundance by 100.

We’re given the following values:

Average Atomic Mass = **63.55 amu**

^{63}X: Mass = **62.9396 amu**

** f.a. = 69.17/100 = 0.6917**

We don’t know the percent abundance of either isotope but we can use the fact that the **f.a. of all isotopes** of an element *add up to 1*.

^{65}X: Mass = **???**

** f.a. = 1 - 0.6917 = 0.3083**

The naturally occurring isotopes of element X consists of the isotopes ^{63}X, and ^{6}^{5}X. The atomic mass of ^{63}X is 62.9396 amu with a percent abundance of 69.17%. With an average atomic mass of X of 63.55 amu, the atomic mass of ^{65}X is:

A. 64.44 amu

B. 79.18 amu

C. 64.92 amu

D. 65.22 amu

E. 69.17 amu

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