🤓 Based on our data, we think this question is relevant for Professor Snaddon's class at IU.

We’re being asked to **calculate the mass of the third isotope of magnesium**. We can use the following equation:

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

where atomic mass = average atomic mass of the element and f.a. = fractional abundance of the isotope. To get ** f.a.**, we simply need to divide the given percent abundance by 100.

Magnesium has three naturally occurring isotopes: ^{24} Mg (23.985 amu) with 78.99% abundance, ^{25} Mg (24.986 amu) with 10.00% abundance, and a third with 11.01% abundance.

Look up the atomic weight of magnesium, and then calculate the mass of the third isotope.