Atomic Structure and Electron Configuration

Charge, nucleus and electron count

An atom has a tiny positive nucleus containing protons of charge +1 and neutrons of charge 0, surrounded by electrons of charge −1; these charges are measured relative to the elementary charge. Nearly all atomic mass is in the nucleus, but most atomic volume is occupied by the electron distribution. Proton count defines the element. Isotope differences change neutron count; ordinary ion formation changes electron count. A neutral atom has equal proton and electron counts. For Mg²⁺, keep 12 protons and remove two electrons, leaving 10. Positive charge does not mean adding protons.

An orbital describes an electron’s probability distribution, not a planet-style path. We organize occupied orbitals into shells and subshells. Each s subshell contains one orbital, p contains three, d contains five and f contains seven; at most two opposite-spin electrons fit in each orbital. The AP exam does not require assigning four quantum numbers.

Read configuration notation one piece at a time. In 2p⁴, the 2 names the principal shell, p names a subshell containing three orbitals, and the superscript 4 gives the four electrons occupying that subshell. It does not mean four p orbitals. In 1s², the single 1s orbital contains two electrons. A ground-state configuration is the lowest-total-energy arrangement for the species. A bracket such as [Ne] abbreviates neon’s filled core, 1s²2s²2p⁶, or ten electrons. Thus [Ne]3s² means those ten core electrons plus two 3s electrons. Decode and total the symbols before using them to compare species.

Build the ground-state configuration

Use the Aufbau pattern as a filling guide: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p for these early elements. Hund’s rule puts electrons singly with parallel spin into equal-energy orbitals before pairing. Pauli’s principle allows at most two electrons in one orbital with opposite spin. Those are different rules. Oxygen has eight total electrons but six valence electrons: 1s²2s²2p⁴. The 2p boxes contain one pair and two singles.

For main-group atoms, outermost-shell electrons are valence electrons and inner electrons are core electrons. For ions, first establish electron count and then remove or add electrons in the appropriate subshells. When a transition metal is ionized, remove the ns electrons before the (n−1)d electrons: Fe is [Ar]4s²3d⁶, but Fe²⁺ is [Ar]3d⁶. You cannot generally give a transition-metal ion the neutral configuration of the element a few boxes to its left. Exception configurations are outside the assessed AP requirement.

Why some electrons are harder to remove

The magnitude of Coulombic force varies as $F\propto\frac{|q_1q_2|}{r^2}$. Opposite charges attract. Doubling separation cuts this pairwise force to one quarter. Electrons in a many-electron atom experience both nuclear attraction and electron–electron repulsion; core electrons shield outer electrons. Effective nuclear charge is the net attractive influence, not simply proton count.

Ionization energy is an energy needed to remove an electron, not itself the Coulombic force. Stronger effective attraction and shorter average distance generally increase that energy. Core electrons are harder to remove than valence electrons. As shells and subshells are emptied, the remaining electrons experience different attractions, so successive ionization energies increase and a large jump can reveal when removal reaches a core shell.