Atomic Structure, Isotopes, and Electron Configurations
Grade 11 chemistry uses the electron structure of an atom to explain periodic patterns and predict chemical behavior.
Learning goals
- Relate atomic number and mass number to protons, neutrons, and electrons.
- Explain why isotopes are the same element but have different masses.
- Distinguish energy levels, sublevels, and atomic orbitals.
- Write electron configurations for neutral atoms and simple ions of the first four periods.
- Apply the Aufbau principle, Pauli exclusion principle, and Hund rule.
1. The nucleus and isotope notation
An atom contains a tiny positive nucleus and an electron cloud. The atomic number $Z$ equals the number of protons and identifies the element. The mass number $A$ equals protons plus neutrons, so the neutron count is $N=A-Z$. A neutral atom has $Z$ electrons; an ion has gained or lost electrons, but its nucleus and element identity remain unchanged.
Isotopes have the same $Z$ but different $N$. For example, carbon-12 and carbon-14 both have six protons, while they contain six and eight neutrons respectively. Isotopes usually have very similar chemical behavior because their electron configurations are the same, but their masses and nuclear stability can differ.
2. Levels, sublevels, and orbitals
Electrons occupy quantized energy states. A principal energy level is labelled $n=1,2,3,...$. Each level contains sublevels: $s$, $p$, $d$, and $f$. An orbital is a region of high probability of finding an electron, not a circular path. An $s$ sublevel has one orbital and can hold 2 electrons; a $p$ sublevel has three orbitals and can hold 6; a $d$ sublevel has five orbitals and can hold 10.