Electron Configurations

An electron configuration tells you exactly where every electron in an atom lives. It uses spectroscopic notation - the principal quantum number (n) followed by the subshell letter, with a superscript showing how many electrons are in that subshell. For example, 1s²2s²2p⁶3s² describes magnesium (12 electrons total).

Writing electron configurations is a skill the MCAT expects you to have, and it boils down to three rules.

The Three Rules

1. Aufbau Principle (“Building Up”)

Electrons fill from the lowest energy subshell to the highest. “Aufbau” is German for “building up.” Each subshell fills completely before the next one begins to fill.

2. Pauli Exclusion Principle

No two electrons in the same atom can have the same four quantum numbers. In practice, this means each orbital holds at most 2 electrons, and those two electrons must have opposite spins.

3. Hund’s Rule

Within a subshell that has multiple orbitals (p, d, or f), electrons fill each orbital singly with parallel spins before any orbital gets a second electron.

The n + l Rule (Filling Order)

The energy of a subshell is determined by the sum n + l. Lower n + l = lower energy = fills first. If two subshells have the same n + l value, the one with the lower n fills first.

| Subshell | n | l | n + l | Filling order |
|----------|---|---|-------|--------------|
| 1s | 1 | 0 | 1 | 1st |
| 2s | 2 | 0 | 2 | 2nd |
| 2p | 2 | 1 | 3 | 3rd |
| 3s | 3 | 0 | 3 | 4th (same n+l as 2p, but lower n) |
| 3p | 3 | 1 | 4 | 5th |
| 4s | 4 | 0 | 4 | 6th (same n+l as 3p, lower n) |
| 3d | 3 | 2 | 5 | 7th |
| 4p | 4 | 1 | 5 | 8th |

This gives the familiar filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

You can also determine this from the diagonal rule diagram - write out the subshells in rows and draw diagonal arrows from upper right to lower left.

Noble Gas (Shorthand) Notation

Writing out the full electron configuration for heavier elements gets tedious. The shorthand uses the symbol of the preceding noble gas in brackets to represent all the core electrons:

• Sodium (Z = 11): Full = 1s²2s²2p⁶3s¹; Shorthand = [Ne] 3s¹
• Iron (Z = 26): Full = 1s²2s²2p⁶3s²3p⁶4s²3d⁶; Shorthand = [Ar] 4s²3d⁶

Orbital Diagrams

An orbital diagram shows each orbital as a line or box with arrows representing electrons. Spin-up electrons are shown as upward arrows and spin-down electrons as downward arrows.

For nitrogen (Z = 7): 1s²2s²2p³

The 2p subshell has 3 orbitals. By Hund’s rule, each gets one electron (all spin-up) before any pairing occurs. You would see three half-filled p orbitals with parallel spins.

The Chromium and Copper Exceptions

Two elements have electron configurations that violate the expected filling order because half-filled and fully filled d subshells are extra stable:

• Chromium (Z = 24): Expected [Ar] 4s²3d⁴, actual [Ar] 4s¹3d⁵ - promotes one 4s electron to achieve a half-filled d subshell
• Copper (Z = 29): Expected [Ar] 4s²3d⁹, actual [Ar] 4s¹3d¹⁰ - promotes one 4s electron to achieve a fully filled d subshell

Paramagnetic vs. Diamagnetic

The presence of unpaired electrons determines an atom’s magnetic behavior:

• Paramagnetic: Has at least one unpaired electron. Weakly attracted to magnetic fields. Unpaired spins align with the field.
• Diamagnetic: All electrons are paired. Weakly repelled by magnetic fields. No unpaired spins to align.

To determine if an atom is paramagnetic, write out its electron configuration and check for unpaired electrons using Hund’s rule.