(a) What isotope is used as the standard in establishing the atomic mass scale? (b) The atomic weight of boron is reported as 10.81 , yet no atom of boron has the mass of 10.81 amu. Explain.

The isotope used as the standard in establishing the atomic mass scale is Carbon-12 (\(^{12}\)C), where 1 atomic mass unit (amu) is defined as \( \frac{1}{12} \) of the mass of a Carbon-12 atom.

The atomic weight (or relative atomic mass) of boron is reported as 10.81 amu. This is because the atomic weight is a weighted average of the masses of the isotopes of an element, taking into consideration their natural abundance. Boron has two stable isotopes, Boron-10 (\(^{10}\)B) and Boron-11 (\(^{11}\)B), which have atomic masses of approximately 10 amu and 11 amu, respectively. To calculate the atomic weight of boron, we consider the natural abundance of each isotope and their respective atomic masses: Atomic weight of Boron = (% abundance of Boron-10 × mass of Boron-10) + (% abundance of Boron-11 × mass of Boron-11) Boron-10 has a natural abundance of approximately 19.8% and Boron-11 has a natural abundance of approximately 80.2%. Therefore, we can calculate the atomic weight of boron as follows: Atomic weight of Boron = (0.198 × 10) + (0.802 × 11) = 1.98 + 8.82 = 10.80 amu (rounded to 10.81 amu) No atom of boron has a mass of exactly 10.81 amu, but this value represents the weighted average mass of boron atoms based on the natural abundances of its isotopes.