Element m has two naturally occurring isotopes: m-10 (10.01 amu, 19.6% abundance) and m-11 (11.01 amu, 80.4% abundance). What is the average atomic mass (in amu) of element m?

• A. 11.01
• B. 10.01
• C. 10.81
• D. 10.50

Answer: (C)

Verified step by step guidance

1. Identify the isotopes and their respective masses and natural abundances. Here, isotope m-10 has a mass of 10.01 amu and an abundance of 19.6%, and isotope m-11 has a mass of 11.01 amu and an abundance of 80.4%.
2. Convert the percentage abundances into decimal form by dividing each by 100. For example, 19.6% becomes 0.196 and 80.4% becomes 0.804.
3. Use the formula for average atomic mass, which is the weighted average of the isotopes' masses based on their abundances: \[\text{Average atomic mass} = (\text{mass}_1 \times \text{abundance}_1) + (\text{mass}_2 \times \text{abundance}_2)\]
4. Substitute the values into the formula: \[\text{Average atomic mass} = (10.01 \times 0.196) + (11.01 \times 0.804)\]
5. Calculate each product separately and then add them together to find the average atomic mass of element m.