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.

The stable isotopes of iron are ^{54}Fe, ^{56}Fe, ^{57}Fe, and ^{58}Fe. The mass spectrum of iron looks like the following:

Use the data on the mass spectrum to estimate the average atomic mass of iron, and compare it to the value given in the periodic table.

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Based on our data, we think this problem is relevant for Professor Norton's class at UGA.

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Our data indicates that this problem or a close variation was asked in Chemistry: An Atoms First Approach - Zumdahl Atoms 1st 2nd Edition. You can also practice Chemistry: An Atoms First Approach - Zumdahl Atoms 1st 2nd Edition practice problems.