Key Concepts in Atomic Structure and Isotopes

A nice little medium-length practice set to get you pumped up for AP Chemistry!

1 amu or 1 u

Defined as 1/12 the mass of a carbon-12 atom

Atomic number

The number of protons in an atom; determines the element

Mass number

The total number of protons and neutrons in an atom

Nucleon

A particle in the nucleus; either a proton or a neutron

Proton

Positively charged subatomic particle found in the nucleus

Neutron

Neutral subatomic particle found in the nucleus

Electron

Negatively charged subatomic particle found outside the nucleus

Atomic mass

The weighted average of all naturally occurring isotopes of an element

Isotope

Atoms of the same element with different numbers of neutrons

Average atomic mass

(mass of isotope 1 × % abundance) + (mass of isotope 2 × % abundance) + …

Relative atomic mass

A unitless ratio comparing atomic mass to 1/12 of carbon-12

Naturally occurring isotopes

Isotopes found in nature without human synthesis

Synthetic isotope

A man-made isotope created in labs or nuclear reactions

AMU in kg

1 u ≈ 1.6605 × 10⁻²⁷ kg

Scientific notation

A way to express very small or large numbers using powers of 10

10⁻³

Equals 0.001

10⁻⁶

Equals 0.000001

Hydrogen isotopes

Protium (0n), Deuterium (1n), Tritium (2n)

How to calculate neutrons

Mass number – Atomic number

How to estimate atomic mass

Round periodic table value and subtract atomic number to get average neutrons

Why isotopes exist

Neutrons help stabilize the nucleus; different combinations occur in nature

Difference between mass number and atomic mass

Mass number is protons + neutrons (whole number); atomic mass is weighted average (decimal)

Chlorine's average atomic mass

About 35.5 u due to mix of Cl-35 and Cl-37 isotopes

Heaviest stable atom

Tin (Sn), with 10 naturally occurring isotopes

No stable isotopes

Elements like Technetium (Tc) and Promethium (Pm) have no stable natural isotopes