A certain element consists of two stable isotopes. The first has a mass of $138 \mathrm{amu}$ and a percent natural abundance of $8.90 \times 10^{-2} \%$. The second has a mass of 139 amu and a percent natural abundance of $99.9 \%$. What is the atomic weight of the element? $\mathrm{amu}$

Step 1 ：Given that a certain element consists of two stable isotopes. The first has a mass of \(138 \, \mathrm{amu}\) and a percent natural abundance of \(8.90 \times 10^{-2} \%\). The second has a mass of \(139 \, \mathrm{amu}\) and a percent natural abundance of \(99.9 \%\).

Step 2 ：The atomic weight of an element is calculated by taking the weighted average of the atomic masses of its isotopes. The weight of each isotope is its abundance. So, we need to multiply the mass of each isotope by its abundance (in decimal form), and then add these products together.

Step 3 ：Let's denote the mass of the first isotope as \(\text{mass1} = 138\), and its abundance as \(\text{abundance1} = 0.0008900000000000001\).

Step 4 ：Similarly, let's denote the mass of the second isotope as \(\text{mass2} = 139\), and its abundance as \(\text{abundance2} = 0.9990000000000001\).

Step 5 ：Then, the atomic weight of the element can be calculated as \(\text{atomic_weight} = \text{mass1} \times \text{abundance1} + \text{mass2} \times \text{abundance2} = 138.98382\).

Step 6 ：Final Answer: The atomic weight of the element is \(\boxed{138.98382 \, \mathrm{amu}}\).