Let's break down the problem step by step:

(a) Percent Abundance of the Heaviest Isotope

• The table shows three isotopes.
• The heaviest isotope is No. 3 with an atomic mass of 77.01 u.
• According to the table, the percent abundance of isotope No. 3 is 23%.

(b) Calculation of Atomic Mass

The formula to calculate the atomic mass of a hypothetical element, based on isotopic masses and their abundances, is:

$\text{Atomic Mass} = (\text{Isotopic Mass}_1 \times \text{Abundance}_1) + (\text{Isotopic Mass}_2 \times \text{Abundance}_2) + (\text{Isotopic Mass}_3 \times \text{Abundance}_3)$

Substituting the given values from the table:

$\text{Atomic Mass} = (73.56 \times 0.12) + (75.98 \times 0.65) + (77.01 \times 0.23)$

Let's calculate each term:

1. $73.56 \times 0.12 = 8.8272$
2. $75.98 \times 0.65 = 49.387$
3. $77.01 \times 0.23 = 17.7123$

Adding these up:

$\text{Atomic Mass} = 8.8272 + 49.387 + 17.7123 = 75.9265 \, \text{u}$

Therefore, the atomic mass of this hypothetical element is approximately 75.93 u.

Would you like any further details or clarifications on this solution? Here are some related questions you might find useful:

1. How are atomic masses calculated in real-life elements?
2. What are isotopes, and why do their abundances affect atomic mass?
3. How do scientists determine the percent abundance of isotopes in nature?
4. Why does the atomic mass on the periodic table often appear as a decimal?
5. Can isotopic distributions change in different environments?

Tip: Always ensure to convert percent abundances to decimals when performing calculations related to atomic mass.