Atomic Structure

Atomic number (Z) = protons. Mass number (A) = protons + neutrons.

Isotopes: atoms with the same number of protons but different numbers of neutrons.

Mass spectrometry determines:

• Relative atomic mass from isotope abundances
• Molecular formula from molecular ion peak ($M^+$)

$A_r = \frac{\sum(\text{mass} \times \text{abundance})}{\sum \text{abundance}}$

Example: Cl has ⁣⁣⁣$^{35}$Cl (75.77%) and $^{37}$Cl (24.23%): $A_r = \frac{35 \times 75.77 + 37 \times 24.23}{100} = 35.45$

SL Level

Electrons occupy energy levels: 2, 8, 18, 32 maximum per shell.

HL Level — Subshells and Orbitals

Subshells: s (1 orbital, 2e⁻), p (3 orbitals, 6e⁻), d (5 orbitals, 10e⁻), f (7 orbitals, 14e⁻)

Filling order: $1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p...$

Rules: Aufbau (lowest energy first), Pauli (max 2e⁻ per orbital, opposite spins), Hund's (fill orbitals singly first)

Exceptions: Cr = $[\text{Ar}] 3d^5 4s^1$, Cu = $[\text{Ar}] 3d^{10} 4s^1$ (stability of half/fully-filled d)

When atoms are excited, electrons jump to higher energy levels. When they return, they emit photons of specific wavelengths → line emission spectrum.

$\Delta E = h\nu = \frac{hc}{\lambda}$

Evidence for quantised energy levels: discrete lines (not continuous spectrum)

Convergence: lines get closer together at higher frequencies → ionisation energy can be determined from the convergence limit.

$\text{X}(g) \rightarrow \text{X}^+(g) + e^-$

Trends:

• Increases across a period (more protons, same shielding)
• Decreases down a group (more shells, more shielding)
• Anomalies at Groups 13 and 16 (subshell and pairing effects)

Successive IEs show jumps that reveal the number of electrons in each shell.

• Isotopes: same protons, different neutrons
• Electron configuration: Aufbau, Pauli, Hund's rule; HL includes subshells
• Emission spectra: evidence for quantised energy levels; $\Delta E = h\nu$
• IE trends: increases across period, decreases down group