Question

(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 \(10.81 \mathrm{u}\). Explain.

Short answer

(a) Carbon-12 is the standard isotope. (b) Boron's atomic weight, 10.81 u, is an average of its isotopes' masses.

Step-by-step solution

Isotope Standard for Atomic Mass

The isotope used as the standard for establishing the atomic mass scale is carbon-12. This isotope of carbon is defined to have an atomic mass of exactly 12 atomic mass units (u). As a result, all other atomic masses are measured relative to this standard.

Understanding Atomic Weight

The atomic weight of an element such as boron is not a whole number because it is a weighted average of the masses of all its isotopes, taking into account their relative abundances. This means the atomic weight of boron is a composite value, calculated based on the existence and proportions of its isotopes.

Example with Boron

Boron has two stable isotopes: boron-10 and boron-11. These isotopes have atomic masses of about 10 u and 11 u respectively. The atomic weight of boron, 10.81 u, reflects a weighted average of these isotopic masses according to their natural abundance.

Key concepts

Isotope Standard

The isotope standard is an essential part of the atomic mass scale, which allows scientists to measure the mass of atoms. The chosen standard for this scale is carbon-12. Carbon-12 is a naturally occurring isotope of carbon, and its significance lies in its defined atomic mass of exactly 12 atomic mass units (u). This means that other atomic masses are compared and measured in relation to this fixed value of carbon-12.
By choosing carbon-12 as a reference point, consistency is maintained across scientific measurements, making it easier for scientists worldwide to compare and communicate findings. This standardization is crucial for various applications, from studying chemical reactions to developing new materials.

• Carbon-12 is abundant in nature, making it accessible and reliable.
• Its use as a baseline ensures uniformity in atomic mass measurements.

Atomic Weight

The concept of atomic weight can be puzzling, because it often doesn't match the mass of any individual atom. Rather than being a simple number, atomic weight is a sophisticated calculation representing an average of all the isotopes of an element, weighted by their abundance. For example, boron's atomic weight is approximately 10.81 u.
Why isn't there a boron atom with a mass of exactly 10.81 u? This discrepancy arises because atomic weight reflects the "average" condition of the element, rather than any one specific isotope. All isotopes are taken into account, and their masses are averaged according to how commonly they are found in nature.
This concept allows scientists to better understand and predict how the element will behave in different chemical contexts, where all isotopes can influence reactions.

• Atomic weight expresses the average mass of an element's atoms.
• It takes into consideration the relative masses of all isotopes.
• Isotope abundance significantly affects calculated atomic weights.

Weighted Average

Weighted average is a vital concept in the calculation of atomic weights. It's different from a simple average because it factors in the significance, or weight, of each entry. When determining atomic weight, scientists calculate the weighted average of isotopic masses, considering how abundant each isotope is.
For boron, two primary isotopes exist: boron-10 and boron-11. These isotopes don't exist in equal amounts. Hence, their contribution to boron's atomic weight is not uniform. The atomic weight is calculated by multiplying the mass of each isotope by its natural abundance and then adding these values together.
This process ensures that the atomic weight reflects the actual distribution of isotopes in nature. Understanding weighted averages is crucial for interpreting chemical data and predicting an element's properties in reactions.

• Weighted average accounts for different levels of importance or occurrence.
• Isotopes with higher abundance have a greater impact on atomic weight.
• It's used to provide an accurate representation of an element's typical atomic mass.