To solve for the percent abundances of the two isotopes of element A, we can use the following relationship for calculating the average atomic mass:

$\text{Average atomic mass} = (\text{fraction of isotope 1} \times \text{atomic mass of isotope 1}) + (\text{fraction of isotope 2} \times \text{atomic mass of isotope 2})$

Let:

• $x$ be the fraction of isotope 16A (which has a mass of 15.994915 amu)
• $(1 - x)$ be the fraction of isotope 17A (which has a mass of 16.999131 amu)
• The average atomic mass of element A is given as 16.00 amu

Now set up the equation:

$16.00 = (x \times 15.994915) + \left((1 - x) \times 16.999131\right)$

We can solve this equation to find $x$, which represents the fraction (or percent abundance) of isotope 16A. Then, $1 - x$ will give the fraction (or percent abundance) of isotope 17A.

Let's solve for $x$.The fraction of isotope 16A ($x$) is approximately 0.9949, or 99.49%.

Therefore, the fraction of isotope 17A ($1 - x$) is approximately $1 - 0.9949 = 0.0051$, or 0.51%.

Percent abundances:

• Isotope 16A: 99.49%
• Isotope 17A: 0.51%

Would you like further clarification or details on this?

Here are some related questions:

1. How do isotopes affect the chemical properties of an element?
2. What are the methods to experimentally determine isotope abundances?
3. How is atomic mass different from atomic number?
4. What are some common elements with multiple stable isotopes?
5. How do isotopic compositions vary in different environments?

Tip: The small difference in atomic masses of isotopes can significantly influence their nuclear stability, even if chemical behavior remains similar.