(b) Why is the atomic weight of carbon reported as 12.011 in the table of elements and the periodic table in the front inside cover of this text?

Verified step by step guidance

1. Understand that the atomic weight of an element is a weighted average of the masses of its isotopes.
2. Recognize that carbon has two stable isotopes: Carbon-12 and Carbon-13.
3. Carbon-12 is the most abundant isotope, making up about 98.89% of carbon found in nature, while Carbon-13 makes up about 1.11%.
4. The atomic weight of carbon (12.011) reflects the average mass of these isotopes, weighted by their natural abundance.
5. The calculation involves multiplying the mass of each isotope by its relative abundance and summing these values to get the average atomic weight.

Key Concepts

Atomic Weight

Atomic weight, or atomic mass, is the weighted average mass of an element's isotopes, measured in atomic mass units (amu). It reflects the relative abundance of each isotope in nature, which is why it is not always a whole number. For carbon, the atomic weight of 12.011 accounts for the presence of both carbon-12 and carbon-13 isotopes.

Isotopes

Isotopes are variants of a chemical element that have the same number of protons but different numbers of neutrons, resulting in different atomic masses. For carbon, the most common isotopes are carbon-12 (with 6 neutrons) and carbon-13 (with 7 neutrons). The existence of these isotopes contributes to the average atomic weight reported in the periodic table.

Relative Abundance

Relative abundance refers to the proportion of each isotope of an element found in a natural sample. This concept is crucial for calculating the atomic weight, as it determines how much each isotope contributes to the overall average. For carbon, the relative abundance of carbon-12 and carbon-13 influences the reported atomic weight of 12.011.