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 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 with the following masses and natural abundances:

Isotope | Mass (amu) | Abundance (%) |

Mg-24 | 23.9850 | 78.99 |

Mg-25 | 24.9858 | 10.00 |

Mg-26 | 25.9826 | 11.01 |

Calculate the atomic mass of magnesium.

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