S2

Mass number (A)

The total number of protons and neutrons in the nucleus of an atom

Atomic number (Z)

The number of protons in the nucleus of an atom.

Isotope

Atoms of the same element with the same number of protons but different numbers of neutrons.

Relative atomic mass (Aᵣ)

The weighted mean mass of the atoms of an element relative to 1/12 of the mass of a carbon-12 atom.

Line spectrum

A spectrum consisting of discrete lines at specific wavelengths corresponding to electron transitions between energy levels.

Continuous spectrum

A spectrum that contains all wavelengths of electromagnetic radiation with no gaps.

Ionisation energy (first ionisation energy)

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1⁺ ions.

What EM radiation is released when an electron goes transitions to the n=1 energy level in hydrogen?

UV

What EM radiation is released when an electron goes transitions to the n=2 energy level in hydrogen?

Visible light

What EM radiation is released when an electron goes transitions to the n=3 energy level in hydrogen?

IR

Radioisotope

An isotope with an unstable nucleus that undergoes radioactive decay.

Alpha particle (α)

A helium nucleus consisting of 2 protons and 2 neutrons, emitted during radioactive decay.

Beta particle (β⁻)

A high-speed electron emitted from the nucleus during radioactive decay.

Carbon-14 dating - application

A method used to determine the age of organic materials based on the decay of carbon-14.

Electromagnetic spectrum

The range of all types of electromagnetic radiation arranged according to frequency or wavelength.

Ground state vs excited state

• Ground state: the lowest energy level of an electron in an atom.

• Excited state: any higher energy level an electron occupies after absorbing energy.

Ion

A charged particle formed when an atom or group of atoms gains or loses electrons.

Cation

A positively charged ion formed by loss of electrons.

Anion

A negatively charged ion formed by gain of electrons.

Equation relating frequency and wavelength

c=λν

c (speed of light (in a vacuum)): 3.00 × 10⁸ m s⁻¹

λ (wavelength):

ν (frequency):

First ionisation energy

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1⁺ ions.

Sublevels (subshells) - and location periodic table

s, p, d, f

Atomic orbital

A region of space around the nucleus where there is a high probability of finding an electron.

Pauli exclusion principle

An orbital can hold a maximum of two electrons with opposite spins.

Aufbau principle

Electrons occupy orbitals of lowest energy first before filling higher energy orbitals.

Hund’s rule

Electrons occupy degenerate orbitals singly with parallel spins before pairing.

Electron spin

A property of electrons describing two possible orientations (↑ or ↓), giving rise to magnetic behavior.

Quantum numbers

Values that describe the position and energy of an electron in an atom (principal, angular momentum, magnetic, spin).

In what way do electrons have energy?

Energy is quantized; electrons occupy discrete energy levels rather than continuous values.

Stability and its relation to energy in particles

Lower energy states are more stable; systems tend to move toward minimum energy.

s orbital

A spherical orbital centered on the nucleus (holds a maximum of 2 electrons).

p orbital

A dumbbell-shaped orbital with two lobes oriented along axes (set of three orbitals, max 6 electrons).

Photon

A discrete packet (quantum) of electromagnetic radiation with energy proportional to frequency.

Bohr’s model

Electrons occupy fixed energy levels (shells) around the nucleus and emit/absorb energy when transitioning between levels.

Limitations of Bohr’s model

• Only works for hydrogen-like atoms (one electron).

• Does not explain fine spectral details or electron behavior in multi-electron atoms.

• Inconsistent with quantum mechanics (no electron probability).

Exceptions to the Aufbau principle - and why

• Chromium (Cr):
  Electron configuration: [Ar] 3d⁵ 4s¹ (instead of 3d⁴ 4s²)

• Copper (Cu):
  Electron configuration: [Ar] 3d¹⁰ 4s¹ (instead of 3d⁹ 4s²)

Half-filled and fully filled d sublevels have lower energy due to reduced electron–electron repulsion and increased stability.

Factors affecting ionisation energy

• Nuclear charge (more protons → higher attraction → higher IE)

• Distance from nucleus (greater distance → lower IE)

• Shielding effect (more inner shells → lower IE)

• Electron configuration (sublevel/repulsion effects)

Limit of convergence (don’t need to know definition, just understand concept)

The point at which spectral lines converge, corresponding to the ionisation energy (electron completely removed).

Successive ionisation energies

The energies required to remove electrons one at a time from gaseous atoms; large jumps indicate removal from a new (inner) shell.