Where everything starts. Electrons in orbitals, protons in nuclei, and a periodic table that explains chemistry from a few simple rules.
The required syllabus content for Unit 1, in order. Each card is one lesson-sized checkpoint.
Lesson 0 of Unit 1.
Lesson 1 of Unit 1.
Distinguish between the properties of elements, compounds and mixtures.
Conceptual understanding: The mass of an atom is concentrated in its minute, positively charged nucleus.
Lesson 4 of Unit 1.
1) Emission and absorption spectra provide both evidence for: A. the existence of neutrons B. the existence of isotopes C. the existence of atomic energy levels D. the nuclear model of an atom
Conceptual understanding: The electron configuration of an atom can be deduced from its atomic number.
Each lesson card below mirrors the original teacher deck — syllabus refs, content, worked examples and practice questions in order.
Chemistry begins with classifying matter. The opening lesson establishes the three-way taxonomy of elements, compounds and mixtures — and the laboratory techniques used to separate the last category.
Everything is made of matter — but not all matter is the same. Elements contain only one type of atom; compounds contain two or more elements chemically joined; mixtures are elements/compounds together in the same vessel without chemical bonding.
Pure substances (elements and compounds) have a uniform chemical composition. Mixtures can be homogeneous (same state, uniform composition — e.g. dissolved salt in water) or heterogeneous (different states/non-uniform — e.g. sand in water, oil and water).
Particles in a solid vibrate about fixed positions. Heating increases their kinetic energy — temperature is a measure of average kinetic energy. Adding enough energy lets particles overcome the forces holding them in place: solid → liquid → gas. The reverse: gas → liquid → solid. Direct s → g is sublimation; g → s is deposition.
Adding an impurity (any soluble solute) depresses the melting point and elevates the boiling point of a solvent. The magnitude depends on the number of solute particles, not their identity — that's what "colligative" means.
Every species in a chemical equation should carry a state symbol: (s) solid, (l) liquid, (g) gas, (aq) aqueous (dissolved in water).
Through the 19th and 20th centuries, models of the atom became progressively more refined: Dalton's indivisible spheres → Thomson's "plum pudding" → Rutherford's nuclear atom → Bohr's quantised shells → the modern quantum-mechanical model.
α-particles fired at a thin gold foil. Three observations and what each tells us:
| Particle | Location | Relative mass | Relative charge |
|---|---|---|---|
| Proton | Nucleus | 1 | +1 |
| Neutron | Nucleus | 1 | 0 |
| Electron | Surrounds nucleus | 1/1836 | −1 |
Atomic number Z = number of protons. Defines the element. Mass number A = protons + neutrons. Written in nuclear notation as AZX. For neutral atoms, # electrons = # protons. For an ion of charge q, # electrons = Z − q.
The periodic table lists chlorine with Ar = 35.45 — but you can't have half a proton. The value is a weighted average across all the naturally occurring isotopes of chlorine: 35Cl (≈ 75%) and 37Cl (≈ 25%).
Isotopes = atoms of the same element (same Z) with different numbers of neutrons (different A). They have identical chemical properties (same electrons, same bonding) but slightly different physical properties — mass, density, rates of effusion, and any property depending on inertia. This is why isotopes can only be separated by mass-based techniques (mass spec, centrifugation), never by chemistry.
The formula:
Ar = Σ (isotopic mass × fractional abundance)
where each fractional abundance is the % divided by 100. Use percentages straight from a mass spectrum if you prefer.
White light contains all visible wavelengths → a continuous rainbow. Hot atomic hydrogen emits only certain wavelengths → discrete bright lines on a black background. This line emission spectrum is the direct experimental evidence that electrons in atoms have only certain allowed energies — energy is quantised.
Heat or electrical excitation promotes electrons to higher energy levels. As they fall back down, photons are emitted. Photon energy = (Ehigh − Elow) = h·ν. Each gap is fixed, so each line is fixed.
c = ν λ · speed of light = frequency × wavelength. (c = 3.00 × 10⁸ m s⁻¹.)
E = h ν · photon energy = Planck constant × frequency. (h = 6.626 × 10⁻³⁴ J s.)
Combined: E = hc / λ.
Energy levels get closer as n increases. The spectral lines converge to a limit; beyond that limit, the electron is no longer bound. The energy at this convergence limit is the ionisation energy for that series.
The Bohr model only worked for one-electron atoms. The modern model treats electrons probabilistically — an atomic orbital is a region of space where an electron is likely (~90% probability) to be found. Shells subdivide into sublevels (s, p, d, f), each containing a fixed number of orbitals.
The three p-orbitals are degenerate (equal in energy). Similarly the five d's, the seven f's. Maximum electrons per shell = 2n² (so n=1→2, n=2→8, n=3→18, n=4→32).
Groups 1–2 fill s; groups 13–18 fill p; the d-block (groups 3–12, periods 4+) fills d; lanthanides/actinides fill f.
Replace inner-core electrons with the previous noble gas symbol in brackets. Na = [Ne] 3s¹. Fe = [Ar] 4s² 3d⁶.
Half-full and fully-full d-subshells are unusually stable, so the configurations are not what aufbau predicts:
To form a cation, remove electrons from the highest occupied energy level first. For transition metals this means remove 4s before 3d — counter-intuitive but a recurring exam catch.
Each box is one orbital; each arrow is one electron. ↑ for spin-up, ↓ for spin-down. Pauli: max 2 per box, opposite spin. Hund: in a degenerate sublevel, fill each box singly first.
If atoms had only shells (no sub-shells), first IE would rise smoothly across a period. It doesn't — there are two systematic dips in every period. Each dip exposes a sub-shell boundary.
HL pushes deeper: mass spectrometry data, successive ionisation energies, and the limit of convergence in emission spectra.
Lesson 5 of Unit 1.
Lesson 8 of Unit 1.
Practice: The first five successive ionisation energies of an unknown element are: 801, 2427, 3660, 25026 and 32827 kJ mol-1. Deduce the group of the periodic table in which this element is likely to be found.
A mass spectrometer ionises a sample, accelerates the ions through a magnetic field, and separates them by mass-to-charge ratio (m/z). The output is a "stick" plot with one line per ion species — height proportional to relative abundance.
For an atomic sample (e.g. Cl atoms), each m/z value corresponds to one isotope. Peak heights give the abundances directly. Ar is then the weighted mean.
If chlorine is fed in as Cl₂, the molecular ions are 35Cl−35Cl (mass 70), 35Cl−37Cl (mass 72) and 37Cl−37Cl (mass 74). With ~3:1 ratio of 35Cl:37Cl, the probabilities give relative peak heights of 9 : 6 : 1.
The first ionisation energy (IE₁) is the energy required to remove the most loosely held electron from one mole of gaseous atoms:
X(g) → X⁺(g) + e⁻ ΔH = IE₁
On a hydrogen emission spectrum (Lyman series), the spectral lines converge at high frequency. The line at the convergence limit corresponds to a transition from n = ∞ down to n = 1 — meaning the electron has come back from the just-ionised state. Therefore the energy of a photon at the convergence frequency, multiplied by Avogadro's number, equals the molar ionisation energy.
IE (J mol⁻¹) = h · νlimit · NA
Successive ionisation energies are the energies needed to remove the 1st, 2nd, 3rd … electrons in turn:
X⁺(g) → X²⁺(g) + e⁻ (IE₂) ; X²⁺(g) → X³⁺(g) + e⁻ (IE₃) …
Each successive removal is from an increasingly positive ion, with fewer electrons sharing the same nuclear pull. So IE₂ > IE₁ always — and the increases get steeper.
Removing the last electron of a given shell is much easier than removing the first electron of the next-inner shell, which is closer to the nucleus and less shielded. So a jump in successive IE plotted on a log scale signals a shell boundary. Counting electrons up to each jump reveals the electron configuration directly.
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