The correct option is **B. Atomic mass is the average mass of an element's isotopes, weighted by their natural abundance.** ### Detailed Explanation: 1. **Understanding Atomic Mass**: - Atomic mass is a measure of the mass of an atom, typically expressed in atomic mass units (amu). It reflects the average mass of an element's isotopes, taking into account how common each isotope is in nature. 2. **Isotopes**: - Isotopes are variants of a particular chemical element that have the same number of protons but different numbers of neutrons. For example, carbon has isotopes like Carbon-12 (6 protons, 6 neutrons) and Carbon-14 (6 protons, 8 neutrons). The atomic mass of carbon is not simply the mass of one isotope but an average that considers the relative abundance of each isotope. 3. **Weighted Average**: - The atomic mass is calculated as a weighted average. This means that if an isotope is more abundant, it contributes more to the average atomic mass than a less abundant isotope. The formula for calculating the atomic mass can be expressed as: \[ \text{Atomic Mass} = \sum \left( \text{mass of isotope} \times \text{fractional abundance} \right) \] - For example, if Carbon-12 makes up 98.89% of carbon atoms and Carbon-14 makes up 1.11%, the atomic mass of carbon would be calculated as: \[ \text{Atomic Mass} = (12 \, \text{amu} \times 0.9889) + (14 \, \text{amu} \times 0.0111) \approx 12.011 \, \text{amu} \] ### Why the Other Options are Incorrect: - **Option A: Atomic mass is the mass of an atom measured in grams.** - This statement is misleading. While atomic mass can be converted to grams, it is conventionally expressed in atomic mass units (amu). The atomic mass is not simply the mass of an atom in grams; it is a relative measure based on the mass of carbon-12, which is defined as exactly 12 amu. - **Option C: Atomic mass is equal to the number of protons in an atom's nucleus.** - This statement is incorrect because the atomic mass is not just the number of protons (which is the atomic number). The atomic mass also includes neutrons, which contribute significantly to the overall mass of the atom. For example, carbon has 6 protons and 6 neutrons in its most common isotope (Carbon-12), but its atomic mass is 12 amu, not just 6. - **Option D: Atomic mass is the total number of electrons in a neutral atom.** - This statement is also incorrect. The atomic mass does not relate to the number of electrons. In a neutral atom, the number of electrons equals the number of protons, but electrons have negligible mass compared to protons and neutrons. Therefore, they do not significantly contribute to the atomic mass. ### Common Pitfalls: - Confusing atomic mass with atomic number (the number of protons). - Misunderstanding the concept of isotopes and their contribution to atomic mass. - Forgetting that atomic mass is a weighted average based on natural abundance. ### Revision Summary: - Atomic mass is the average mass of an element's isotopes, weighted by their natural abundance. - Isotopes have the same number of protons but different numbers of neutrons. - Atomic mass is expressed in atomic mass units (amu), not grams. - The atomic mass calculation involves a weighted average based on the abundance of each isotope.