Electron Configuration Review

Flashcards covering key vocabulary related to electron configuration, quantum numbers, orbital filling rules, exceptions, ion formation, and magnetism.

Electron Configuration

Describes how electrons are arranged in energy levels (shells), subshells, and orbitals within an atom.

Principal quantum number (n)

Controls the shell level, size, and energy of an electron, roughly corresponding to the periodic table period number.

Azimuthal (subshell) number (l)

Controls the subshell type (s, p, d, f), related to the shape and angular momentum of an orbital.

Magnetic quantum number (ml)

Controls the specific orbital orientation within a subshell (e.g., px, py, pz).

Spin quantum number (ms)

Controls the electron spin direction, with values of +½ or -½.

s subshell

Has an azimuthal number (l) of 0, contains 1 orbital, and can hold a maximum of 2 electrons (spherical shape).

p subshell

Has an azimuthal number (l) of 1, contains 3 orbitals, and can hold a maximum of 6 electrons (dumbbell shape).

d subshell

Has an azimuthal number (l) of 2, contains 5 orbitals, and can hold a maximum of 10 electrons (clover/donut shapes).

f subshell

Has an azimuthal number (l) of 3, contains 7 orbitals, and can hold a maximum of 14 electrons (complex shapes).

Pauli Exclusion Principle

States that no two electrons in an atom can have the same set of four quantum numbers, meaning an orbital can hold at most 2 electrons with opposite spins.

Hund's Rule

States that within a subshell of degenerate (same-energy) orbitals, electrons occupy empty orbitals singly with parallel spins before pairing up.

Aufbau Principle

States that electrons fill the lowest available energy orbitals first.

(n + l) rule

Used to determine the energy ordering of orbitals: orbitals with a lower (n + l) value fill first; if tied, the orbital with the lower 'n' fills first.

s-block

Refers to elements in groups 1-2 and Helium on the periodic table, where electrons fill the s subshell.

p-block

Refers to elements in groups 13-18 on the periodic table, where electrons fill the p subshell.

d-block

Refers to transition metals (groups 3-12), where electrons fill the (n-1)d subshells.

f-block

Refers to lanthanides and actinides, where electrons fill the (n-2)f subshells.

Noble-gas shorthand

A condensed way to write electron configurations by replacing filled core shells with the symbol of the nearest noble gas in brackets.

Valence electrons

Electrons in the highest principal quantum number (n) shell for main-group elements, which determine an atom's reactivity.

Core electrons

All electrons in an atom that are not valence electrons, typically those in filled inner shells.

Orbital Diagram

A visual representation using boxes for orbitals and arrows (↑/↓) for electron spins to depict electron configuration.

Transition metal exceptions

Elements like Chromium, Copper, Molybdenum, Silver, Gold, and Palladium that exhibit unusual electron configurations (often shifting an s electron to a d subshell) to achieve half-filled (d⁵) or filled (d¹⁰) stability.

Main-group cations (ion formation)

Formed by removing electrons from the highest 'n' shell first (valence s/p electrons).

Main-group anions (ion formation)

Formed by adding electrons to fill the p subshell, often to reach a noble-gas configuration.

Transition-metal cations (ion formation)

Formed by removing s electrons before d electrons, even if the 'd' subshell was filled after 's'.

Paramagnetic

Describes a substance with at least one unpaired electron, causing it to be attracted to a magnet.

Diamagnetic

Describes a substance where all electrons are paired, causing it to be slightly repelled by a magnet.

Periodic table period

The row number on the periodic table, which approximately corresponds to the highest occupied principal shell number (n) for s/p elements.

Periodic table group

The column number for main-group elements, indicating the number of valence electrons (e.g., group 17 has 7 valence electrons).