HL IB Chemistry – Electronic Configurations & Spectra

30 vocabulary flashcards covering key terms on electromagnetic spectra, atomic structure, electronic configurations and ionisation energies for IB HL Chemistry.

Electromagnetic Spectrum

The complete range of electromagnetic radiation, showing the relationship between frequency, wavelength and energy from radio waves to gamma rays.

Frequency (f)

The number of wave crests passing a point per second; inversely proportional to wavelength.

Wavelength (λ)

The distance between two consecutive peaks of a wave; inversely proportional to frequency.

Speed of Light (c)

A constant value of 3.00 × 10^8 m s⁻¹ for all electromagnetic radiation; related to f and λ by c = fλ.

Continuous Spectrum

A spectrum that contains all wavelengths of visible light, like a rainbow produced by white light dispersion.

Line Spectrum

A spectrum displaying only specific wavelengths, indicating quantized energy transitions in atoms.

Quanta (Quantum)

Discrete packets of energy that electrons can absorb or emit; evidence of energy level quantization.

Emission Spectrum

The set of specific wavelengths emitted when excited electrons fall to lower energy levels.

Absorption Spectrum

The pattern of dark lines formed when electrons absorb specific wavelengths to move to higher energy levels.

Convergence Limit

The region in a line spectrum where spectral lines merge, corresponding to ionisation energy.

Ionisation Energy (IE)

The energy required to remove one mole of electrons from one mole of gaseous atoms.

Lyman Series

Ultraviolet spectral lines produced when electrons fall to the n = 1 energy level in hydrogen.

Balmer Series

Visible spectral lines produced when electrons fall to the n = 2 energy level in hydrogen.

Energy Level (Shell)

A region around the nucleus where electrons with similar energy are found, designated by principal quantum number n.

Principal Quantum Number (n)

An integer (1,2,3…) that labels energy levels; lower n is closer to the nucleus and lower in energy.

Subshell

A division within an energy level, labelled s, p, d, or f, each with characteristic energy.

Orbital

A 3-D region within a subshell that can hold a maximum of two electrons with opposite spins.

s Orbital

A spherical orbital; each s subshell contains one orbital that can hold two electrons.

p Orbital

A dumbbell-shaped orbital; each p subshell has three (px, py, pz) holding up to six electrons.

d Orbital

One of five orbitals in a d subshell; combined they hold up to ten electrons and have more complex shapes.

Aufbau Principle

Electrons fill orbitals of lowest available energy first to give the ground-state configuration.

Hund's Rule

Electrons occupy separate degenerate orbitals with parallel spins before pairing to minimise repulsion.

Pauli Exclusion Principle

No two electrons in the same atom can have identical sets of quantum numbers; an orbital holds max two electrons with opposite spins.

Full Electron Configuration

Notation listing every occupied subshell with superscripts showing electron numbers from 1s upwards.

Shorthand Electron Configuration

Abbreviated notation using the previous noble gas in brackets followed by remaining subshells.

Transition Metal

Element filling 3d or 4d subshells; loses 4s electrons first when forming positive ions.

Successive Ionisation Energies

Series of energies required to remove electrons one by one from the same atom, increasing each time.

Shielding Effect

Reduction of nuclear attraction on outer electrons due to repulsion by inner-shell electrons.

Nuclear Charge

Total positive charge of the nucleus (number of protons) influencing electrostatic attraction on electrons.

Spin-Pair Repulsion

Extra repulsion experienced by two electrons sharing the same orbital, slightly lowering ionisation energy.