Given isotopes 12C (98.89%) and 13C (1.11%) with masses 12.0000 amu and 13.00335 amu, respectively, what is the average atomic mass?

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Multiple Choice

Given isotopes 12C (98.89%) and 13C (1.11%) with masses 12.0000 amu and 13.00335 amu, respectively, what is the average atomic mass?

Explanation:
Average atomic mass is the weighted average of the isotopic masses, using their natural abundances as weights. Convert the percentages to decimals, multiply by each isotope’s mass, and add the results. For carbon here: 98.89% is 0.9889, and 1.11% is 0.0111. Do the math: 0.9889 × 12.0000 = 11.8668; 0.0111 × 13.00335 ≈ 0.1443. Add them: 11.8668 + 0.1443 ≈ 12.0111 amu. This gives an average atomic mass of about 12.011 amu, which aligns with the value typically listed for carbon on the periodic table. The heavier carbon-13 contribution, though small, pulls the average above 12.000.

Average atomic mass is the weighted average of the isotopic masses, using their natural abundances as weights. Convert the percentages to decimals, multiply by each isotope’s mass, and add the results.

For carbon here: 98.89% is 0.9889, and 1.11% is 0.0111. Do the math: 0.9889 × 12.0000 = 11.8668; 0.0111 × 13.00335 ≈ 0.1443. Add them: 11.8668 + 0.1443 ≈ 12.0111 amu. This gives an average atomic mass of about 12.011 amu, which aligns with the value typically listed for carbon on the periodic table. The heavier carbon-13 contribution, though small, pulls the average above 12.000.

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