In a periodic table, the average atomic mass of magnesium is given as `24.312 u`. The average value is based on their relative natural abundance on earth. The three isotopes and their masses are `._12Mg^(24) (23.98504u)`, `._(12)Mg^(25) (24.98584)` and `._12Mg^(26) (25.98259 u)`. The natural abundance of `._12Mg^(24)` is `78.99%` by mass. Calculate the abundances of the other two isotopes.

Explanation

To solve the problem, we need to calculate the natural abundances of the isotopes \( \text{Mg}^{25} \) and \( \text{Mg}^{26} \) given the average atomic mass of magnesium and the abundance of \( \text{Mg}^{24} \).

### Step-by-Step Solution:

1. **Identify Given Data:**
- Average atomic mass of magnesium, \( A = 24.312 \, u \)
- Mass of \( \text{Mg}^{24} = 23.98504 \, u \)
- Mass of \( \text{Mg}^{25} = 24.98584 \, u \)
...