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Atoms of the same element can have different numbers of neutrons. This causes them to have different masses, even though they have the same number of protons. These different forms are called isotopes.
For example, chlorine has two main isotopes:

• extsuperscript{35}Cl, which has 18 neutrons.
• extsuperscript{37}Cl, which has 20 neutrons.

Each isotope has a unique atomic mass, because the total number of protons and neutrons in its nucleus differs. Importantly, they still behave the same way chemically because they have the same number of protons and electrons. Understanding isotopes is essential to comprehending how average atomic mass is determined based on their masses and abundances.

Relative abundance refers to how common each isotope of an element is relative to the total amount of the element found in nature. It is usually expressed as a percentage. For instance, for chlorine:

• extsuperscript{35}Cl has a relative abundance of 75.53%.
• extsuperscript{37}Cl has a relative abundance of 24.47%.

Relative abundance must be converted into decimal form when calculating the average atomic mass. You do this by dividing the percentage by 100. For example, 75.53% becomes 0.7553.
When calculating average atomic mass, each isotope's mass is multiplied by its relative abundance, and the results are summed. This helps effectively weigh each isotope's contribution.
Understanding relative abundance ensures that students accurately represent how much of each isotope naturally contributes to an element's overall atomic structure in their calculations.

The atomic mass unit (amu) is the standard unit for expressing atomic mass. It's used to express the mass of atoms and molecules in a way that is easier to manage, given the incredibly small scale at which atoms operate.
One amu is defined as one-twelfth the mass of a carbon-12 atom. Effectively, it serves as a baseline for determining atomic mass on a relative scale.
When calculating average atomic masses, the masses of isotopes are given in amu. This allows for a consistent and comparable way to express the mass contributions of isotopes. For example, the isotope extsuperscript{35}Cl has a mass of 34.968 amu.
Remember, the final average mass of an element, calculated from its isotopes and relative abundances, is also expressed in amu. This consistency simplifies understanding and communicating atomic masses.