📚 Topic Summary
The average atomic mass is the weighted average of the masses of all the isotopes of an element. Isotopes are atoms of the same element that have different numbers of neutrons, leading to different mass numbers. To calculate the average atomic mass, you multiply the mass of each isotope by its relative abundance (expressed as a decimal) and then sum these values. Understanding this calculation is crucial for many AP Chemistry topics!
🧪 Part A: Vocabulary
Match the term with its correct definition:
- Term: Isotope
- Term: Atomic Mass Unit (amu)
- Term: Mass Number
- Term: Average Atomic Mass
- Term: Relative Abundance
- Definition: The total number of protons and neutrons in an atom's nucleus.
- Definition: The weighted average mass of the atoms of an element, considering all its isotopes.
- Definition: Atoms of the same element with different numbers of neutrons.
- Definition: A unit used to express the mass of atoms and molecules.
- Definition: The proportion of each isotope in a naturally occurring sample of an element.
Match the term to the appropriate definition.
🧩 Part B: Fill in the Blanks
Complete the following paragraph using the words provided below:
The (1) is calculated by taking the sum of the (2) of each isotope multiplied by its (3). (4) are atoms of the same element with different numbers of (5).
Words: neutrons, average atomic mass, isotopes, mass, relative abundance
🤔 Part C: Critical Thinking
Explain in your own words why the average atomic mass listed on the periodic table is not usually a whole number.