Question

What is wrong with the following statement? "The atomic mass of an atom of chlorine is \(35.453 \mathrm{u}\)."

Short answer

The statement 'The atomic mass of an atom of chlorine is 35.453 u.' is misleading because 35.453 u is the weighted average atomic mass of chlorine, taking into account its isotopes, and does not represent the mass of an individual chlorine atom.

Step-by-step solution

Understanding Atomic Mass

Atomic mass is the mass of an atom expressed in atomic mass units (u). It represents the average mass of atoms, taking into account the presence of isotopes and their relative abundance. For elements like chlorine that have multiple isotopes, the atomic mass reflects the weighted average of these isotopes rather than the mass of a single atom.

Identify Chlorine Isotopes

Chlorine has two stable isotopes: Chlorine-35 and Chlorine-37. Their natural abundances are about 75.77% and 24.23% respectively. The atomic mass of an element is a weighted average of the atomic masses of each isotope, based on their natural abundance.

Clarifying the Misconception

The statement is incorrect as it suggests that all atoms of chlorine have a mass of 35.453 u. However, this number is the average atomic mass of chlorine that takes into account the masses of the isotopes of chlorine and their relative abundance. Not every individual atom of chlorine will have this exact mass due to the presence of isotopes with different masses.

Correct Interpretation

The correct interpretation is that the atomic mass of chlorine, 35.453 u, is the average mass of chlorine atoms when accounting for the existence and natural abundance of its isotopes. It does not correspond to the mass of any specific atom of chlorine but is a weighted average.

Key concepts

Isotopes

The concept of isotopes is fundamental to understanding atomic structure and chemistry. Isotopes are variants of a particular chemical element that have the same number of protons (and hence the same atomic number) but different numbers of neutrons. As a result, while isotopes of an element have identical chemical properties, they have different atomic masses.

For example, carbon has three naturally occurring isotopes: Carbon-12, Carbon-13, and Carbon-14. All carbon atoms have six protons, but the number of neutrons can vary, which makes Carbon-12 lighter than Carbon-14. The presence of isotopes accounts for the fact that the atomic mass of an element is not a whole number but a decimal value, representing a weighted average of all the isotopes' masses.

Natural Abundance of Isotopes

The natural abundance of isotopes refers to the proportion of each isotope present in a naturally occurring sample of an element. This abundance is crucial when calculating the average atomic mass of an element since isotopes do not occur in equal amounts. For instance, chlorine consists primarily of the isotopes Chlorine-35 and Chlorine-37.

The natural abundance is typically expressed as a percentage. In the case of chlorine, natural samples roughly consist of 75.77% of Chlorine-35 and 24.23% of Chlorine-37. Knowing these percentages allows scientists to determine how much each isotope contributes to the average atomic mass of chlorine.

Average Atomic Mass

Average atomic mass is the weighted average mass of the atoms in a naturally occurring sample of the element. It essentially reflects the mass of an atom on a scale where Carbon-12 is exactly 12 atomic mass units (u).

To calculate the average atomic mass, the mass of each isotope is multiplied by its relative natural abundance, and these products are then summed. It is important to note that the average atomic mass is not the mass of a single atom, but rather a statistical representation of an element's isotopes. As seen in the case of chlorine, the average atomic mass (35.453 u) combines the contributions from both Chlorine-35 and Chlorine-37 based on their natural abundances and individual atomic masses.