9. The Periodic Table: chemical periodicity
- Syllabus
- 9701–2028–2029
- Section
- 9
- Level
- AS
9.1 Periodicity of physical properties of the elements in Period 3
- Syllabus
- 9701–2028–2029
- Topic
- 9.1
- Level
- AS
Period 3 shows four distinct physical-property patterns
| Property | Qualitative variation across Na, Mg, Al, Si, P, S, Cl, Ar |
|---|---|
| atomic radius | decreases from Na to Cl as atoms use the same principal shell while nuclear charge increases; Ar has no directly comparable covalent radius in the usual dataset |
| ionic radius | decreases across Na⁺ → Mg²⁺ → Al³⁺ → Si⁴⁺; jumps to a much larger value at P³⁻; then decreases P³⁻ → S²⁻ → Cl⁻ |
| melting point | rises Na → Mg → Al → Si, with Si highest; drops sharply at P; rises at S; then falls through Cl to Ar |
| electrical conductivity | increases across the metals Na → Mg → Al; drops enormously at Si; P, S, Cl and Ar are effectively non-conducting |
The cation sequence Na⁺, Mg²⁺, Al³⁺ and Si⁴⁺ is isoelectronic with 10 electrons, so increasing proton number pulls the same electron arrangement inward. P³⁻, S²⁻ and Cl⁻ are isoelectronic with 18 electrons and also shrink as proton number rises. The jump occurs because the anions have an occupied third shell, whereas the cations have lost it.
These linked but non-uniform variations are periodicity: patterns recur because outer-shell configuration and resulting structure change systematically across a period.
Do not draw one smooth trend through the ionic radii: cations and anions belong to different electron-shell groups. Do not describe melting point as simply decreasing after Si—S melts above P because S₈ particles are larger than P₄ particles.
Structure controls what must move for melting and for conduction
| Elements | Structure and bonding | Melting-point explanation | Electrical-conductivity explanation |
|---|---|---|---|
| Na, Mg, Al | giant metallic lattice: positive ions attracted to delocalised electrons | generally rises Na → Al as ion charge, number of delocalised electrons and charge density increase, strengthening metallic bonding | all conduct through mobile delocalised electrons; conductivity increases overall toward Al as more electrons are contributed per atom |
| Si | giant covalent network | highest: many strong Si–Si covalent bonds must be broken to melt the network | semiconductor with conductivity far below the metals because only a small number of mobile charge carriers are available |
| P₄, S₈, Cl₂ | simple molecular; instantaneous dipole–induced dipole forces between molecules | much lower than Si because melting overcomes intermolecular forces, not internal covalent bonds; S₈ > P₄ > Cl₂ as molecular size and polarisability increase | essentially do not conduct because the molecules provide no mobile charged particles |
| Ar | monatomic particles with instantaneous dipole–induced dipole attractions | lowest: small Ar atoms have very weak attractions between them | does not conduct because there are no mobile charged particles |
For melting, identify the particles separated and the attractions between those particles. For conduction, ask whether the structure contains charge carriers that can move through the solid. These are separate tests: strong bonding can produce a high melting point without producing conductivity.
Silicon and aluminium are both extended lattices, but Al contains mobile delocalised electrons whereas pure Si has far fewer mobile carriers. Sulfur and phosphorus contain strong covalent bonds inside S₈ and P₄, yet their low melting points reflect the weaker forces between intact molecules.
Argon is monatomic, not a simple molecule. Do not say that melting P₄, S₈ or Cl₂ breaks covalent bonds, and do not infer electrical conduction solely from a high melting point.
9.2 Periodicity of chemical properties of the elements in Period 3
- Syllabus
- 9701–2028–2029
- Topic
- 9.2
- Level
- AS
Period 3 elements form a defined set of oxides and chlorides
| Element | Reaction with oxygen | Main required product |
|---|---|---|
| Na | 4Na + O₂ → 2Na₂O | sodium oxide |
| Mg | 2Mg + O₂ → 2MgO | magnesium oxide |
| Al | 4Al + 3O₂ → 2Al₂O₃ | aluminium oxide |
| P | P₄ + 5O₂ → P₄O₁₀ | phosphorus(V) oxide |
| S | S + O₂ → SO₂ | sulfur dioxide |
| Element | Reaction with chlorine | Main required product |
|---|---|---|
| Na | 2Na + Cl₂ → 2NaCl | sodium chloride |
| Mg | Mg + Cl₂ → MgCl₂ | magnesium chloride |
| Al | 2Al + 3Cl₂ → 2AlCl₃ | aluminium chloride |
| Si | Si + 2Cl₂ → SiCl₄ | silicon(IV) chloride |
| P | P₄ + 10Cl₂ → 4PCl₅ | phosphorus(V) chloride |
| Element | Reaction with water | What is observed |
|---|---|---|
| Na | 2Na + 2H₂O → 2NaOH + H₂ | vigorous reaction; alkaline solution and hydrogen form |
| Mg, cold water | Mg + 2H₂O → Mg(OH)₂ + H₂ | very slow; a little hydrogen and sparingly soluble Mg(OH)₂ form |
| Mg, steam | Mg + H₂O(g) → MgO + H₂ | heated magnesium reacts much more readily |
Learn the products named in the syllabus rather than every possible oxide or chloride. The required phosphorus chloride is PCl₅, and only sodium and magnesium reactions with water are assessed in this outcome.
Oxidation numbers rise as more outer-shell electrons are used
Assign O an oxidation number of −2 and Cl an oxidation number of −1 in these compounds. The oxidation numbers in a neutral formula must sum to zero, so the Period 3 element has the balancing positive value.
| Compound series | Oxidation number of the Period 3 element |
|---|---|
| Na₂O, MgO, Al₂O₃ | Na +1, Mg +2, Al +3 |
| P₄O₁₀ | P +5 |
| SO₂, SO₃ | S +4, S +6 |
| NaCl, MgCl₂, AlCl₃, SiCl₄, PCl₅ | Na +1, Mg +2, Al +3, Si +4, P +5 |
Across Na to P, the number of outer-shell electrons increases from one to five. In the listed highest oxides and chlorides, progressively more of these electrons are transferred or shared with the more electronegative O or Cl, so the maximum positive oxidation number rises from +1 to +5. Sulfur then shows both +4 in SO₂ and +6 in SO₃ because four or all six of its outer-shell electrons are involved in the oxidation-number accounting.
Oxidation number is electron bookkeeping, not the actual charge on an atom in a covalent molecule. Do not force SO₂ and SO₃ into one value: sulfur is +4 in SO₂ but +6 in SO₃.
Water reveals the basic-to-acidic oxide transition
| Oxide | Reaction with water, if any | Likely pH of resulting mixture |
|---|---|---|
| Na₂O | Na₂O + H₂O → 2NaOH | strongly alkaline, about 13–14 |
| MgO | MgO + H₂O → Mg(OH)₂ | mildly alkaline, about 9–10 |
| Al₂O₃ | no reaction | about 7 |
| SiO₂ | no reaction | about 7 |
| P₄O₁₀ | P₄O₁₀ + 6H₂O → 4H₃PO₄ | acidic, about 2 |
| SO₂ | SO₂ + H₂O ⇌ H₂SO₃ | acidic, about 2–3 |
| SO₃ | SO₃ + H₂O → H₂SO₄ | strongly acidic, about 1 |
The ionic oxides on the left supply O²⁻, which accepts protons from water and produces OH⁻. The covalent non-metal oxides on the right react with water to form oxoacids. Al₂O₃ and SiO₂ do not react with water, so their acid-base character must be tested with other reagents.
No reaction with water does not mean an oxide has no acid-base behaviour: Al₂O₃ is amphoteric and SiO₂ is acidic when tested with suitable acid or base reagents.
Oxides and hydroxides change from basic through amphoteric to acidic
| Species | Behaviour | Diagnostic reaction |
|---|---|---|
| Na₂O, MgO | basic oxides | Na₂O + 2HCl → 2NaCl + H₂O; MgO + 2HCl → MgCl₂ + H₂O |
| Al₂O₃ | amphoteric oxide | Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O; Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄ |
| SiO₂ | acidic oxide | SiO₂ + 2NaOH → Na₂SiO₃ + H₂O |
| P₄O₁₀ | acidic oxide | P₄O₁₀ + 12NaOH → 4Na₃PO₄ + 6H₂O |
| SO₂ | acidic oxide | SO₂ + 2NaOH → Na₂SO₃ + H₂O |
| SO₃ | acidic oxide | SO₃ + 2NaOH → Na₂SO₄ + H₂O |
| Hydroxide | Behaviour | Equation evidence |
|---|---|---|
| NaOH | soluble strong base | NaOH + HCl → NaCl + H₂O |
| Mg(OH)₂ | sparingly soluble base | Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O |
| Al(OH)₃ | amphoteric | Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O; Al(OH)₃ + NaOH → NaAl(OH)₄ |
A basic oxide or hydroxide reacts with acid; an acidic oxide reacts with sodium hydroxide; an amphoteric oxide or hydroxide does both. Across the period, the dominant behaviour therefore changes basic → amphoteric → acidic.
For the base reactions in this syllabus outcome, use sodium hydroxide as the base reagent. Amphoteric means reaction with both acids and bases; it does not mean neutral or unreactive.
Period 3 chlorides range from dissolution to hydrolysis
| Chloride added to water | Main change and equation | Likely pH |
|---|---|---|
| NaCl | dissolves: NaCl(s) → Na⁺(aq) + Cl⁻(aq) | about 7 |
| MgCl₂ | dissolves: MgCl₂(s) → Mg²⁺(aq) + 2Cl⁻(aq); hydrated Mg²⁺ makes the solution slightly acidic | about 6–7 |
| AlCl₃ | forms hydrated Al³⁺, which hydrolyses: [Al(H₂O)₆]³⁺ + H₂O ⇌ [Al(H₂O)₅(OH)]²⁺ + H₃O⁺ | about 3 |
| SiCl₄ | vigorous hydrolysis: SiCl₄ + 2H₂O → SiO₂ + 4HCl | about 2 |
| PCl₅ | complete hydrolysis: PCl₅ + 4H₂O → H₃PO₄ + 5HCl | about 2 |
NaCl and MgCl₂ simply form colourless solutions. AlCl₃ gives an acidic solution. SiCl₄ hydrolysis produces acidic hydrogen chloride and solid hydrated silica/SiO₂, often seen with steamy acidic fumes; PCl₅ also hydrolyses vigorously to acidic products.
Dissolving separates pre-existing ions; hydrolysis changes a species by reaction with water. Do not substitute PCl₃ for PCl₅: this outcome names PCl₅.
Electronegativity and bonding explain the chemical trends
Electronegativity increases across Period 3. The electronegativity difference between the Period 3 element and O or Cl therefore decreases, so bonding changes from predominantly ionic on the left to increasingly covalent toward the right. The small, highly charged Al³⁺ ion strongly polarises nearby electron clouds, giving aluminium compounds appreciable covalent character at the transition.
| Observed trend | Bonding explanation | Chemical consequence |
|---|---|---|
| Na₂O and MgO are basic | ionic lattices contain O²⁻ | O²⁻ accepts H⁺ and produces OH⁻ in water |
| Al₂O₃ and Al(OH)₃ are amphoteric | aluminium lies at the ionic–covalent transition | both acid and strong base can react |
| Si, P and S oxides are acidic | covalent central-atom–oxygen bonding; the central atom withdraws electron density | water forms oxoacids where reaction occurs, and the oxides react with NaOH |
| chloride solutions become more acidic | rising cation charge density, then covalent E–Cl bonds, increasingly polarise or react with water | hydrolysis generates H₃O⁺ or HCl |
For oxides of the same element, a higher positive oxidation number makes the central atom more electron-withdrawing. This is why SO₃, with S at +6, gives a more strongly acidic oxide than SO₂, with S at +4.
Do not explain every trend with electronegativity alone. Use electronegativity to infer ionic versus covalent character, then connect that bonding to the relevant particles, polarisation and reaction with water, acid or base.
Use a cluster of observations to infer bonding and structure
| Observation cluster | Best inference | Period 3 examples |
|---|---|---|
| high melting point; brittle solid; conducts when molten or in aqueous solution but not as a solid | giant ionic lattice with mobile ions only when free to move | NaCl, MgCl₂; Na₂O and MgO |
| very high melting point; hard; insoluble; no electrical conduction; no reaction with water | giant covalent network | SiO₂ |
| low melting/boiling point or volatility; no electrical conduction; hydrolysis with water | simple molecular covalent substance | SiCl₄ and molecular P/S oxides; covalent chlorides hydrolyse rather than merely dissociate |
| high-melting oxide that reacts with both acid and base | extended lattice at the ionic–covalent boundary | Al₂O₃ |
Infer in three steps: identify whether charged particles can move, decide whether melting separates ions/atoms or intact molecules, then use dissolution or hydrolysis as supporting chemical evidence. A conclusion is strongest when several observations point to the same model.
AlCl₃ has covalent character and often exists as Al₂Cl₆ molecules; PCl₅ has phase-dependent structure, so do not infer every chloride from one physical observation alone. State the model justified by the given data and use hydrolysis as additional evidence for covalent character.
9.3 Chemical periodicity of other elements
- Syllabus
- 9701–2028–2029
- Topic
- 9.3
- Level
- AS
Predict an element by transferring group patterns and periodic trends
Elements in the same group have the same number of outer-shell electrons, so they tend to form ions with the same charge and compounds with analogous formulae. Moving down the group adds an occupied shell: atomic radius and shielding increase, while the attraction between the nucleus and an outer electron generally weakens.
| Given clue | Prediction it can support | Required explanation |
|---|---|---|
| group number | outer-electron count, common ion charge, analogous oxide/chloride/hydride formulae | same outer-shell pattern |
| position lower in a group | larger atomic radius; generally lower first ionisation energy and electronegativity | extra shell and greater shielding outweigh increased nuclear charge |
| molecular size increases down a molecular group | stronger instantaneous dipole–induced dipole attractions; often higher melting/boiling points | more electrons and greater polarisability |
| known reaction trend within that group | likely relative reactivity of the new element | apply the group's established causal trend, not a universal down-group rule |
Example: let M be the element immediately below Mg in Group 2. Predict a +2 ion because both atoms have two outer-shell electrons; predict formulae MO and MCl₂ by charge balance. M should have a larger radius and lower first ionisation energy than Mg because its outer electrons occupy an additional, more shielded shell.
Write each prediction as evidence → trend → property. A shared group does not make numerical values identical, and reactivity does not always change in the same direction down every group; use the reaction mechanism or a stated group trend before predicting it.
Identify an unknown by intersecting independent clues
| Evidence supplied | What it can locate or classify |
|---|---|
| proton/atomic number | exact element identity |
| electron configuration or a large jump in successive ionisation energies | period from occupied shells; group from outer electrons or the number removed before the jump |
| common ion charge and compound formulae | likely group and oxidation states |
| conductivity, melting/boiling point and physical state | metal/non-metal character and possible metallic, ionic, giant covalent or molecular structure |
| oxide acidity/basicity and reactions with water, acid or base | approximate left-to-right position and metallic/non-metallic character |
| comparison with known neighbouring elements | relative position within a group or period |
- Translate every observation into a constraint. 2. Use the most diagnostic clues to propose a group, period and whether the element is metallic or non-metallic. 3. Intersect the candidates rather than allowing one clue to decide. 4. Test the proposed identity against every remaining physical and chemical observation; reject it if any reliable clue conflicts.
| Given information for E | Deduction |
|---|---|
| E is a Period 3 solid that conducts electricity | metallic candidate in the left/central part of Period 3 |
| a very large successive-ionisation-energy jump occurs after the third electron | three outer-shell electrons; Group 13 |
| E forms E₂O₃ and ECl₃ | oxidation state +3, consistent with the same group clue |
| E₂O₃ is amphoteric | matches the ionic–covalent transition in Period 3 |
| intersection | E is aluminium |
A property such as high melting point or electrical conductivity is not usually unique. State what each clue supports, distinguish a possible position from a confirmed identity, and reserve an exact identity for a unique identifier or a mutually consistent set of independent clues.