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This lesson covers the chemistry of Period 3 elements and their oxides and chlorides, the reactions and trends of Group 2 elements, and the chemistry of Group 7 (the halogens). Understanding periodicity allows you to explain and predict the properties and reactions of elements from their position in the Periodic Table. This is a core inorganic chemistry topic examined across all A-Level specifications and frequently appears in synoptic questions linking structure, bonding, and reactivity.
The eight elements across Period 3 are sodium (Na), magnesium (Mg), aluminium (Al), silicon (Si), phosphorus (P), sulphur (S), chlorine (Cl), and argon (Ar). Their physical properties show clear trends that can be explained by changes in structure and bonding.
Key Definition: The melting point of a substance is the temperature at which it changes from solid to liquid at standard pressure. It depends on the type and strength of the bonding or intermolecular forces that must be overcome.
| Element | Na | Mg | Al | Si | P₄ | S₈ | Cl₂ | Ar |
|---|---|---|---|---|---|---|---|---|
| Structure | Giant metallic | Giant metallic | Giant metallic | Giant covalent | Simple molecular | Simple molecular | Simple molecular | Monatomic |
| Melting point / °C | 98 | 649 | 660 | 1414 | 44 | 115 | −101 | −189 |
Na, Mg, Al (metals with giant metallic structures): Melting point increases from Na to Al because:
Si (giant covalent structure): Silicon has the highest melting point in Period 3. Each silicon atom is covalently bonded to four others in a tetrahedral arrangement (like diamond). The very strong covalent bonds throughout the three-dimensional lattice require a great deal of energy to break.
P₄, S₈, Cl₂ (simple molecular structures): These elements exist as discrete molecules held together by weak London (dispersion) forces. The strength of London forces depends on the number of electrons (and hence the size of the molecule):
Ar (monatomic): Argon exists as individual atoms with very weak London forces. It has the lowest melting point.
Graphite (not in Period 3 but a useful comparison) conducts because of its delocalised electrons within the layers.
The oxides of the Period 3 elements show a clear trend from basic (metal oxides on the left) through amphoteric (aluminium oxide) to acidic (non-metal oxides on the right).
graph LR
A["Na₂O<br/>Strongly Basic"] --> B["MgO<br/>Basic"]
B --> C["Al₂O₃<br/>Amphoteric"]
C --> D["SiO₂<br/>Weakly Acidic"]
D --> E["P₄O₁₀<br/>Acidic"]
E --> F["SO₃<br/>Strongly Acidic"]
F --> G["Cl₂O₇<br/>Strongly Acidic"]
style A fill:#cce5ff
style B fill:#d6e9f8
style C fill:#fff3cd
style D fill:#f8e8d4
style E fill:#f8d7da
style F fill:#f5c6cb
style G fill:#f1b0b7
| Oxide | Formula | Bonding / structure | Acid–base character |
|---|---|---|---|
| Sodium oxide | Na₂O | Ionic, giant lattice | Strongly basic |
| Magnesium oxide | MgO | Ionic, giant lattice | Basic |
| Aluminium oxide | Al₂O₃ | Ionic with some covalent character, giant lattice | Amphoteric |
| Silicon dioxide | SiO₂ | Giant covalent (macromolecular) | Weakly acidic |
| Phosphorus(V) oxide | P₄O₁₀ | Simple molecular (covalent) | Acidic |
| Sulphur trioxide | SO₃ | Simple molecular (covalent) | Strongly acidic |
| Dichlorine heptoxide | Cl₂O₇ | Simple molecular (covalent) | Strongly acidic |
Na₂O: Dissolves readily to form a strongly alkaline solution.
Na₂O(s) + H₂O(l) → 2NaOH(aq) pH ≈ 13–14
MgO: Slightly soluble — forms a weakly alkaline suspension.
MgO(s) + H₂O(l) → Mg(OH)₂(aq) pH ≈ 9–10
MgO is only sparingly soluble because the lattice enthalpy of MgO is very high (Mg²⁺ and O²⁻ have high charges and small radii), so it is difficult to break apart.
Al₂O₃: Insoluble in water — does not react. However, it is amphoteric:
With acid: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O With base: Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄ (or 2Na[Al(OH)₄])
SiO₂: Insoluble in water (giant covalent structure — very strong Si–O bonds). However, SiO₂ is weakly acidic and reacts with hot concentrated NaOH:
SiO₂ + 2NaOH → Na₂SiO₃ + H₂O
P₄O₁₀: Reacts vigorously with water to form phosphoric(V) acid.
P₄O₁₀(s) + 6H₂O(l) → 4H₃PO₄(aq) pH ≈ 1
SO₃: Reacts vigorously (and exothermically) with water to form sulphuric acid.
SO₃(g) + H₂O(l) → H₂SO₄(aq) pH ≈ 0–1
This is a strongly acidic solution because H₂SO₄ is a strong acid that fully dissociates.
Cl₂O₇: Reacts with water to form perchloric acid.
Cl₂O₇(l) + H₂O(l) → 2HClO₄(aq) pH ≈ 0
Exam Tip: The trend in acid–base character of Period 3 oxides is one of the most commonly examined topics. Always link the trend to the type of bonding: ionic oxides are basic because O²⁻ ions react with water to form OH⁻; covalent oxides are acidic because the non-metal–oxygen bonds are polar and the oxide reacts with water to form an oxoacid (H⁺ donor). Al₂O₃ is amphoteric — remember it reacts with both acids and bases.
| Chloride | Formula | Bonding / structure | Melting point |
|---|---|---|---|
| Sodium chloride | NaCl | Ionic, giant lattice | High (801 °C) |
| Magnesium chloride | MgCl₂ | Ionic, giant lattice | High (714 °C) |
| Aluminium chloride | Al₂Cl₆ (AlCl₃ as monomer) | Ionic/covalent border — covalent in anhydrous form, forms a dimer | Sublimes at 178 °C |
| Silicon tetrachloride | SiCl₄ | Simple molecular (covalent) | Low (−68 °C) |
| Phosphorus(V) chloride | PCl₅ | Simple molecular (covalent) | Low (sublimes at 160 °C) |
NaCl: Dissolves to form a neutral solution (pH ≈ 7). No hydrolysis occurs.
NaCl(s) + water → Na⁺(aq) + Cl⁻(aq)
MgCl₂: Dissolves to form a slightly acidic solution (pH ≈ 6.5). The small, highly charged Mg²⁺ ion attracts water molecules strongly, and some hydrolysis occurs:
MgCl₂(s) + water → Mg²⁺(aq) + 2Cl⁻(aq) (slight hydrolysis gives weakly acidic solution)
AlCl₃: Reacts vigorously with water, producing white fumes of HCl and a strongly acidic solution.
AlCl₃(s) + 3H₂O(l) → Al(OH)₃(s) + 3HCl(g) pH ≈ 3
The Al³⁺ ion has a very high charge density, which strongly polarises the O–H bonds in surrounding water molecules, releasing H⁺ ions.
SiCl₄: Reacts vigorously with water, producing white fumes of HCl and a white precipitate of hydrated silicon dioxide.
SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(g) pH ≈ 1–2
This is a hydrolysis reaction — the Si–Cl bonds are broken and replaced by Si–O bonds.
PCl₅: Reacts vigorously with water in a hydrolysis reaction.
PCl₅(s) + 4H₂O(l) → H₃PO₄(aq) + 5HCl(g) pH ≈ 1
White fumes of HCl are observed. The reaction is very exothermic.
Exam Tip: The key trend is that ionic chlorides (NaCl, MgCl₂) dissolve in water to form neutral or slightly acidic solutions, whereas covalent chlorides (AlCl₃, SiCl₄, PCl₅) undergo hydrolysis with water, producing HCl fumes and acidic solutions. The reason for hydrolysis is that the non-metal–Cl bonds are polar and susceptible to attack by water acting as a nucleophile.
Group 2 metals react with water to form a metal hydroxide and hydrogen gas:
M(s) + 2H₂O(l) → M(OH)₂(aq) + H₂(g)
Reactivity increases down the group: Be does not react; Mg reacts very slowly with cold water (but vigorously with steam); Ca reacts steadily; Sr and Ba react vigorously.
The increase in reactivity is because ionisation energies decrease down the group (the outer electrons are further from the nucleus, in higher energy levels, with more shielding), so it becomes easier to form M²⁺ ions.
Group 2 metals burn in oxygen to form white solid oxides:
2M(s) + O₂(g) → 2MO(s)
For example: 2Mg(s) + O₂(g) → 2MgO(s) — brilliant white flame.
Key Definition: Solubility is the maximum amount of a substance that dissolves in a given volume of solvent at a specified temperature.
The solubility of Group 2 hydroxides increases down the group:
| Compound | Solubility |
|---|---|
| Mg(OH)₂ | Slightly soluble (sparingly soluble) |
| Ca(OH)₂ | Slightly soluble (limewater) |
| Sr(OH)₂ | More soluble |
| Ba(OH)₂ | Readily soluble |
Consequently, the pH of the saturated solutions increases down the group, because more OH⁻ ions are in solution.
Ca(OH)₂ solution is limewater — used as a test for carbon dioxide (turns milky when CO₂ is bubbled through, due to the formation of insoluble CaCO₃).
The solubility of Group 2 sulphates decreases down the group:
| Compound | Solubility |
|---|---|
| MgSO₄ | Soluble (Epsom salts) |
| CaSO₄ | Slightly soluble |
| SrSO₄ | Insoluble |
| BaSO₄ | Very insoluble |
The insolubility of BaSO₄ is exploited in the test for sulphate ions: add acidified barium chloride solution to the test sample. A white precipitate of BaSO₄ confirms the presence of sulphate ions. The acid (dilute HCl) is added first to prevent false positives from carbonate or sulphite ions, which would also form white precipitates with Ba²⁺ but dissolve in acid.
Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s) (white precipitate)
Group 2 carbonates decompose on heating to form the metal oxide and carbon dioxide:
MCO₃(s) → MO(s) + CO₂(g)
Thermal stability increases down the group: MgCO₃ decomposes at about 350 °C, CaCO₃ at about 840 °C, SrCO₃ at about 1290 °C, and BaCO₃ at about 1360 °C.
The reason is related to the polarising power of the M²⁺ cation. Smaller cations (e.g. Mg²⁺) have a higher charge density and polarise the electron cloud of the large CO₃²⁻ anion more strongly. This weakens a C–O bond within the carbonate ion, making decomposition easier (lower temperature required). Larger cations (e.g. Ba²⁺) have a lower charge density and polarise the carbonate less, so a higher temperature is needed.
Exam Tip: When explaining trends in thermal stability, always refer to the charge density (charge/radius ratio) of the cation and its ability to polarise the anion. A higher charge density means greater polarisation, which weakens the bonds in the anion and lowers the decomposition temperature.
The halogens (fluorine, chlorine, bromine, iodine) are reactive non-metals in Group 7 of the Periodic Table. They exist as diatomic molecules (F₂, Cl₂, Br₂, I₂).
| Property | F₂ | Cl₂ | Br₂ | I₂ |
|---|---|---|---|---|
| State at room temperature | Pale yellow gas | Green-yellow gas | Red-brown liquid | Dark grey/purple solid |
| Boiling point | −188 °C | −34 °C | 59 °C | 184 °C |
| Electronegativity | 4.0 | 3.0 | 2.8 | 2.5 |
Boiling point increases down the group because the molecules become larger with more electrons, leading to stronger London (dispersion) forces between molecules. More energy is required to overcome these stronger intermolecular forces.
Electronegativity decreases down the group because the atomic radius increases and there is more electron shielding, so the nucleus has less attraction for bonding electrons.
Key Definition: An oxidising agent is a substance that accepts electrons from another species, causing it to be oxidised. The halogen itself is reduced (gains electrons) to form the halide ion.
Oxidising ability decreases down the group: F₂ > Cl₂ > Br₂ > I₂
This is because the larger halogens have a greater atomic radius, so the incoming electron is further from the nucleus and experiences more shielding. The electron affinity becomes less exothermic, making it harder to gain an electron.
A more reactive halogen can displace a less reactive halide ion from solution. This demonstrates the trend in oxidising ability:
| Reaction mixture | Cl₂ added | Br₂ added | I₂ added |
|---|---|---|---|
| KCl(aq) | No reaction | No reaction | No reaction |
| KBr(aq) | Orange/brown colour (Br₂ displaced) | No reaction | No reaction |
| KI(aq) | Brown colour (I₂ displaced) | Brown colour (I₂ displaced) | No reaction |
Ionic equations:
Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq)
Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq)
Br₂(aq) + 2I⁻(aq) → 2Br⁻(aq) + I₂(aq)
In each case, the stronger oxidising agent (higher halogen) oxidises the halide ion (removes electrons from it).
graph TD
A["Halogen Displacement Reactions"] --> B["Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂<br/>Cl₂ displaces Br⁻ ✓"]
A --> C["Cl₂ + 2I⁻ → 2Cl⁻ + I₂<br/>Cl₂ displaces I⁻ ✓"]
A --> D["Br₂ + 2I⁻ → 2Br⁻ + I₂<br/>Br₂ displaces I⁻ ✓"]
A --> E["Br₂ + 2Cl⁻ → No reaction ✗<br/>Br₂ cannot displace Cl⁻"]
A --> F["I₂ + 2Cl⁻ → No reaction ✗<br/>I₂ + 2Br⁻ → No reaction ✗"]
G["Oxidising power:<br/>Cl₂ > Br₂ > I₂"]
When chlorine is added to water, it undergoes disproportionation — the same element is simultaneously oxidised and reduced:
Cl₂(g) + H₂O(l) ⇌ HClO(aq) + HCl(aq)
In this reaction, chlorine goes from oxidation state 0 in Cl₂ to +1 in HClO (oxidised) and −1 in HCl (reduced).
HClO (chloric(I) acid, also known as hypochlorous acid) is a powerful disinfectant. This reaction is the basis for water treatment — chlorine kills bacteria by oxidising cell membranes.
In alkaline solution (e.g. with NaOH):
Cl₂(g) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H₂O(l)
The ClO⁻ ion (chlorate(I), also called hypochlorite) is the active ingredient in bleach.
The identity of a halide ion in solution can be confirmed by adding dilute nitric acid followed by aqueous silver nitrate:
| Halide ion | Precipitate | Colour | Solubility in dilute NH₃ | Solubility in conc. NH₃ |
|---|---|---|---|---|
| Cl⁻ | AgCl | White | Soluble (dissolves) | Soluble |
| Br⁻ | AgBr | Cream | Insoluble | Soluble (dissolves) |
| I⁻ | AgI | Yellow | Insoluble | Insoluble |
Ionic equations:
Ag⁺(aq) + Cl⁻(aq) → AgCl(s) (white precipitate)
Ag⁺(aq) + Br⁻(aq) → AgBr(s) (cream precipitate)
Ag⁺(aq) + I⁻(aq) → AgI(s) (yellow precipitate)
The precipitates dissolve in ammonia by forming soluble complex ions, e.g.:
AgCl(s) + 2NH₃(aq) → [Ag(NH₃)₂]⁺(aq) + Cl⁻(aq)
The dilute nitric acid is added first to prevent interference from carbonate or hydroxide ions, which would form silver carbonate or silver hydroxide precipitates.
flowchart TD
A["Add dilute HNO₃<br/>then AgNO₃(aq)"] --> B{"Colour of<br/>precipitate?"}
B -->|"White"| C["AgCl → Cl⁻ present"]
B -->|"Cream"| D["AgBr → Br⁻ present"]
B -->|"Yellow"| E["AgI → I⁻ present"]
C --> F{"Add dilute NH₃"}
F -->|"Dissolves"| G["Confirmed Cl⁻"]
D --> H{"Add dilute NH₃"}
H -->|"Does NOT dissolve"| I{"Add conc. NH₃"}
I -->|"Dissolves"| J["Confirmed Br⁻"]
E --> K{"Add dilute NH₃"}
K -->|"Does NOT dissolve"| L{"Add conc. NH₃"}
L -->|"Does NOT dissolve"| M["Confirmed I⁻"]
Exam Tip: When describing the silver nitrate test, always state that dilute nitric acid is added first (to remove interfering ions), then silver nitrate solution is added, and finally the colour of the precipitate is noted. If the precipitate needs further identification, add dilute then concentrated ammonia in sequence. AgCl dissolves in dilute NH₃; AgBr dissolves only in concentrated NH₃; AgI does not dissolve in either.
Reducing ability increases down the group: I⁻ > Br⁻ > Cl⁻ > F⁻
This is because the larger halide ions hold their outer electrons less tightly (greater distance from the nucleus, more shielding), making it easier to lose an electron (be oxidised).
graph TD
A["NaCl + H₂SO₄"] -->|"Acid-base only<br/>No redox"| B["HCl gas<br/>(steamy white fumes)"]
C["NaBr + H₂SO₄"] -->|"Acid-base then<br/>partial reduction"| D["HBr initially<br/>then Br₂ + SO₂<br/>(orange fumes)"]
E["NaI + H₂SO₄"] -->|"Acid-base then<br/>progressive reduction"| F["HI initially<br/>then I₂ + SO₂ + S + H₂S<br/>(purple solid, yellow solid,<br/>rotten egg smell)"]
G["Reducing power:<br/>I⁻ > Br⁻ > Cl⁻"]
This trend is demonstrated by the reactions of solid sodium halides with concentrated sulphuric acid:
NaCl + H₂SO₄: Produces steamy white fumes of HCl. This is an acid–base reaction, not a redox reaction, because Cl⁻ is not a strong enough reducing agent to reduce H₂SO₄.
NaCl(s) + H₂SO₄(l) → NaHSO₄(s) + HCl(g)
NaBr + H₂SO₄: Initially produces HBr (steamy fumes), but Br⁻ is a stronger reducing agent than Cl⁻ and can partially reduce H₂SO₄:
2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O
Orange/brown fumes of Br₂ and colourless SO₂ gas are observed.
NaI + H₂SO₄: I⁻ is the strongest reducing agent and reduces H₂SO₄ further:
2HI + H₂SO₄ → I₂ + SO₂ + 2H₂O (initial reduction) 6HI + H₂SO₄ → 3I₂ + S + 4H₂O (further reduction to sulphur) 8HI + H₂SO₄ → 4I₂ + H₂S + 4H₂O (further reduction to hydrogen sulphide)
Observations: purple/black solid iodine, yellow solid sulphur, colourless SO₂ gas (pungent smell), and H₂S gas (rotten egg smell).
Exam Tip: This sequence of reactions is a favourite in exams. Remember: NaCl gives only HCl (no redox); NaBr gives HBr and then Br₂ + SO₂ (partial reduction); NaI gives HI and then I₂ + SO₂ + S + H₂S (progressive reduction). The reason for the trend is the increasing reducing power of the halide ions down the group.