9701 Chemistry · Topics 1 & 3 · AS + A Level
Atomic Structure and Chemical Bonding Cheat Sheet — A Level Chemistry 9701
Topics 1 and 3 are the foundation every other 9701 topic stands on — and the two places examiners most often catch students out with a definition they half-remember. This sheet condenses subatomic particles, isotopes, electron configuration and orbitals, successive ionisation energies, bond types, VSEPR shapes and bond angles, electronegativity and intermolecular forces into a single revision page you can read in about ten minutes.
✅ 22 worked sections
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What’s on this cheat sheet
Atomic Structure
01 · Subatomic particles
| Particle | Charge | Rel. mass |
|---|---|---|
| proton | +1 | 1 |
| neutron | 0 | 1 |
| electron | −1 | 1/1836 |
Notation ᴬZX — Z = protons (= atomic number), A = mass number = p + n, so n = A − Z. Nucleus holds nearly all the mass; charge = protons − electrons.
02 · Isotopes
Same element, same Z, different number of neutrons.
Same chemical properties (same electron configuration). Different physical properties — mass, density, rate of diffusion, boiling point, radioactivity.
Uses: C‑14 dating, Co‑60 radiotherapy, I‑131 thyroid imaging.
03 · Relative atomic mass
Mass spectrum: x‑axis = m/z, y‑axis = relative abundance. Weighted mean, so Ar sits nearer the more abundant isotope.
(35 × 75.8 + 37 × 24.2) ÷ 100 = 35.48
2 d.p., no units — Ar is a ratio.
04 · Electron configuration
Shell n holds max 2n². Subshell capacity: s 2, p 6, d 10, f 14.
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p
Aufbau lowest energy first · Pauli max 2 e⁻ per orbital, opposite spins · Hund singly fill degenerate orbitals before pairing.
Condensed: Fe = [Ar] 4s² 3d⁶. Exceptions: Cr = [Ar] 3d⁵ 4s¹, Cu = [Ar] 3d¹⁰ 4s¹.
Ions: remove the outermost (highest n) electrons first — Fe³⁺ = [Ar] 3d⁵, not 3d³4s².
05 · Orbitals A2
An orbital is a region of high probability of finding an electron; it holds up to 2 e⁻.
s spherical, one per shell. p two‑lobed, three per shell along x, y, z, equal in energy. Draw orbital diagrams as boxes with arrows, applying Hund.
06 · Ionisation energy
Always endothermic; species must be gaseous.
Across a period ↑ — greater nuclear charge, same shell, similar shielding. Down a group ↓ — outer electron further out and more shielded.
Dips: Be→B (electron enters a higher‑energy 2p) and N→O (paired‑electron repulsion in one 2p orbital). These dips are the evidence for sub‑shells.
Successive IEs always rise; a large jump marks the start of a new inner shell, so the number of electrons removed before the jump gives the group.
07 · Evidence for shells and sub‑shells
Plot log(IE) against the number of the electron removed. The large jumps mark the start of each new inner shell, and the number of electrons removed before the first jump gives the group.
Within a shell the smaller steps mark sub‑shells: for Al the 3p electron leaves first, then the two 3s electrons cost noticeably more.
Successive ionisation energies always rise — each electron is pulled from an increasingly positive ion.
08 · Periodicity
| Property | → period | ↓ group |
|---|---|---|
| Atomic radius | ↓ | ↑ |
| Ionisation energy | ↑ | ↓ |
| Electronegativity | ↑ | ↓ |
| Electron affinity | more −ve | less −ve |
| Metallic character | ↓ | ↑ |
Ionic radius: cations smaller than the atom, anions larger. Across an isoelectronic series (N³⁻ → Al³⁺) radius falls as charge rises.
Oxides across period 3: basic (Na₂O, MgO) → amphoteric (Al₂O₃) → acidic (P₄O₁₀, SO₃). Melting point peaks at Si (giant covalent) and collapses at the noble gas.
09 · Worked example — successive ionisation energies
The big jump comes between the 3rd and 4th values, so three electrons are removed easily.
→ Group 13, and with these values the element is aluminium.
Always quote the position of the jump, not just its size.
Definitions to quote exactly
First ionisation energy — the energy needed to remove one mole of electrons from one mole of gaseous atoms.
Electronegativity — the ability of an atom to attract a shared pair of electrons in a covalent bond.
Isotopes — atoms of the same element with the same number of protons but different numbers of neutrons.
Relative atomic mass — the weighted mean mass of an atom relative to carbon‑12 taken as exactly 12.
Marks lost here
— Omitting (g) or the electron in an ionisation equation.
— Writing 3d before 4s when 4s fills first, or removing 3d electrons before 4s when forming ions.
— Saying “shielding increases” across a period; it stays roughly constant — nuclear charge is doing the work.
— Explaining a trend by “more shells” when the ions are isoelectronic.
Chemical Bonding
10 · Which bond type?
| Δ electronegativity | Bond |
|---|---|
| < 0.4 | non‑polar covalent |
| 0.4 – 1.8 | polar covalent |
| > 1.8 | ionic |
Metal + non‑metal → ionic · non‑metal + non‑metal → covalent · metal + metal → metallic. Bonding is a continuum, so treat the cut‑offs as guides and use the bonding triangle when asked.
11 · Ionic bonding
Electrostatic attraction between oppositely charged ions in a giant lattice. Formula is empirical — the ratio, not a molecule.
Properties: high melting point, brittle, conducts when molten or aqueous but not solid (ions fixed), generally soluble in polar solvents.
Lattice enthalpy ↑ with larger ionic charges and smaller ionic radii — MgO ≫ NaCl.
12 · Covalent bonding
A shared pair of electrons attracted by both nuclei. As bond order rises, length falls and strength rises: C–C > C=C > C≡C in length, reversed in enthalpy.
Coordinate (dative) bond: both electrons from one atom — NH₄⁺, H₃O⁺, Al₂Cl₆. Identical to a normal covalent bond once formed.
13 · Lewis structures
1 · Count valence electrons (add for negative charge, subtract for positive). 2 · Least electronegative atom central. 3 · Single bonds, complete octets, then multiple‑bond the leftovers. 4 · Show all lone pairs and the charge in brackets.
Formal charge = valence − (lone‑pair e⁻ + ½ bonding e⁻). The best structure keeps formal charges nearest zero, negative on the most electronegative atom.
Resonance (O₃, NO₃⁻, CO₃²⁻, benzene): delocalised electrons, all bonds equal and intermediate in length and strength.
A2 Period‑3 atoms can exceed an octet: PCl₅, SF₆, ClF₃, XeF₄.
14 · VSEPR — shapes & angles
| Dom. | lp | Shape | Angle | Ex. |
|---|---|---|---|---|
| 2 | 0 | linear | 180° | CO₂ |
| 3 | 0 | trigonal planar | 120° | BF₃ |
| 3 | 1 | bent | <120° | SO₂ |
| 4 | 0 | tetrahedral | 109.5° | CH₄ |
| 4 | 1 | trigonal pyramidal | 107° | NH₃ |
| 4 | 2 | bent | 104.5° | H₂O |
| 5 | 0 | trigonal bipyramidal | 90 / 120° | PCl₅ |
| 5 | 1 | see‑saw | <90 / <120° | SF₄ |
| 5 | 2 | T‑shaped | <90° | ClF₃ |
| 5 | 3 | linear | 180° | XeF₂ |
| 6 | 0 | octahedral | 90° | SF₆ |
| 6 | 1 | square pyramidal | <90° | BrF₅ |
| 6 | 2 | square planar | 90° | XeF₄ |
Rows with 5 or 6 domains are A2. Repulsion lp–lp > lp–bp > bp–bp; each lone pair closes the angle by ~2.5°. Count domains, not atoms — a double bond is one domain.
15 · Polarity
A bond is polar if Δχ > 0 → dipole δ+ → δ−. A molecule is polar only if the bond dipoles do not cancel.
Non‑polar despite polar bonds (symmetrical): CO₂, CCl₄, BF₃, CH₄, SF₆. Polar: H₂O, NH₃, CHCl₃, SO₂.
Symmetry beats bond count — check the shape first, then add the vectors.
16 · Intermolecular forces
Weakest → strongest:
1 London (dispersion) — in everything; stronger with more electrons (higher Mr) and greater surface contact, so pentane boils above 2,2‑dimethylpropane.
2 Dipole–dipole — between permanent dipoles in polar molecules.
3 Hydrogen bonding — H bonded directly to N, O or F, plus a lone pair to accept. Explains the anomalous boiling points of H₂O, NH₃ and HF, and ice being less dense than water.
Melting and boiling points break intermolecular forces, never covalent bonds. “Like dissolves like” for solubility.
17 · Metallic & giant covalent
Metallic: lattice of cations in a sea of delocalised electrons. Strength ↑ with ionic charge and ↓ with radius (Na < Mg < Al). Conducts, malleable, high melting point; alloys are harder because different sizes disrupt the layers.
Giant covalent: diamond — tetrahedral, 4 bonds, very hard, non‑conductor. Graphite — hexagonal layers, 3 bonds + 1 delocalised e⁻, conducts along layers, soft lubricant. SiO₂ resembles diamond; C₆₀ is molecular, not giant.
18 · Hybridisation, σ & π A2
| Hybrid | Domains | Geometry | Ex. |
|---|---|---|---|
| sp³ | 4 | tetrahedral 109.5° | CH₄ |
| sp² | 3 | trigonal planar 120° | C₂H₄ |
| sp | 2 | linear 180° | C₂H₂ |
σ = head‑on overlap along the internuclear axis. π = side‑on overlap of parallel p orbitals. Double bond = 1σ + 1π; triple = 1σ + 2π. π bonds block rotation, giving cis/trans isomers.
Delocalisation: in benzene the six p orbitals overlap sideways into one π system above and below the ring, so all six carbon–carbon bonds are identical and intermediate between a single and a double bond.
19 · Worked example — deduce the shape
S has 6 valence electrons; four bond to F, leaving one lone pair.
5 electron domains → trigonal bipyramidal arrangement
The lone pair takes an equatorial position → see‑saw
Angles slightly under 90° and 120°.
Data booklet & marks lost
Look up, never memorise: electronegativity values, ionic and covalent radii, bond enthalpies and lengths, first ionisation energies, the electromagnetic spectrum, and the periodic table itself.
— Confusing electron domain geometry with molecular shape: NH₃ is tetrahedral in domains, trigonal pyramidal in shape.
— Calling a molecule with polar bonds polar without checking symmetry.
— Saying covalent bonds break on boiling a simple molecular substance.
— Drawing Lewis structures without lone pairs or the ion’s charge bracket.
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Atomic Structure and Chemical Bonding — Frequently Asked Questions
What is the order of filling for the 4s and 3d sub-shells?
4s fills before 3d because it is at a lower energy in the empty atom, but 4s electrons are removed first when the ion forms. So chromium is [Ar] 3d⁵ 4s¹ and Fe²⁺ is [Ar] 3d⁶, not [Ar] 3d⁴ 4s².
How do you explain the jumps in successive ionisation energies?
A large jump between two successive ionisation energies means the next electron is being removed from a shell closer to the nucleus with less shielding. Count the electrons removed before the jump — that number is the group number of the element.
What bond angles do I need to know for 9701?
Linear 180°, trigonal planar 120°, tetrahedral 109.5°, trigonal pyramidal 107°, bent 104.5°, trigonal bipyramidal 120°/90° and octahedral 90°. Each lone pair reduces the angle by roughly 2.5° because lone pair–bond pair repulsion is greater than bond pair–bond pair repulsion.
