The Periodic Table: chemical periodicity

📚The chemistry

Period 3 runs from sodium to argon. Across it the elements change from reactive metals, through a giant covalent metalloid, to simple molecular non-metals, and their oxides and chlorides change from ionic and basic to covalent and acidic. Almost every fact on this page follows from one idea: electronegativity increases across the period, so bonding to oxygen or chlorine becomes less ionic and more covalent.

9.1 Physical properties across period 3

Atomic and ionic radius

Atomic radius decreases from Na to Cl: nuclear charge rises by one each step, electrons go into the same (third) shell, and shielding hardly changes, so the outer electrons are pulled closer (topic 1).

Ionic radii fall into two groups. The cations Na⁺, Mg²⁺, Al³⁺ have lost the third shell entirely, so they are much smaller than their atoms, and shrink further across the series as the nuclear charge rises on the same (2,8) arrangement. The anions P³⁻, S²⁻, Cl⁻ keep three shells and gain electrons, so they are much larger, and they also shrink across the series. The biggest jump in the whole period is between Al³⁺ and P³⁻.

Melting point and conductivity

• Na → Al: metallic bonding gets stronger. Each atom contributes more delocalised electrons (1, 2, 3), and the ions are smaller and more highly charged, so the attraction between ions and the sea of electrons increases. Melting point rises, and so does electrical conductivity: more delocalised electrons per atom are available to carry charge. Aluminium conducts best.
• Si: giant covalent. Melting needs many strong covalent bonds broken, so its melting point is the highest in the period. Its electrons are held in bonds; it conducts only slightly — a semiconductor.
• P, S, Cl, Ar: simple molecular. Melting separates molecules held by weak id-id forces, so melting points are low. Their order follows the size of the molecule, because more electrons give stronger id-id forces: S₈ (128 electrons) > P₄ (60) > Cl₂ (34) > Ar (18, single atoms). None of them conducts.

9.2 Chemical properties across period 3

Reactions of the elements

With oxygen (heated in air or oxygen):

Sodium burns with a yellow-orange flame and magnesium with a brilliant white flame, both giving white solids. White phosphorus catches fire spontaneously in air, giving clouds of white P₄O₁₀. Sulfur burns with a blue flame to give colourless, choking SO₂; further oxidation to SO₃ needs a catalyst (the Contact process, topic 7).

With chlorine (heated):

Sodium and magnesium give white ionic solids. Aluminium chloride forms as a pale yellow-white solid that sublimes (it exists as Al₂Cl₆). SiCl₄ is a colourless liquid. PCl₅, a pale yellow solid, needs excess chlorine; with less, PCl₃ forms.

With water (sodium and magnesium only):

Sodium floats, melts into a ball and fizzes, giving a strongly alkaline solution (pH 13–14). Magnesium barely reacts with cold water, giving a weakly alkaline solution (about pH 9–10) over days; heated in steam it burns with a bright white glow.

Oxidation numbers in the oxides and chlorides

The highest oxidation number rises by one across the period because it equals the number of outer-shell (valence) electrons the element can use in bonding to the more electronegative oxygen or chlorine: sodium has one, sulfur six. Sulfur reaches +6 in SO₃ but only +4 in SO₂, where it uses four of its six. Period 3 elements can use all their outer electrons like this because they can expand their octet (topic 3).

The oxides with water

• Na₂O: Na₂O + H₂O → 2NaOH. Reacts readily; pH 13–14.
• MgO: MgO + H₂O → Mg(OH)₂. Only slightly soluble; pH about 9.
• Al₂O₃ and SiO₂: no reaction, insoluble; the water stays at pH 7. The ionic lattice of Al₂O₃ is too strong, and SiO₂ is giant covalent.
• P₄O₁₀: P₄O₁₀ + 6H₂O → 4H₃PO₄. Violent; pH about 0–1.
• SO₂: SO₂ + H₂O → H₂SO₃ (sulfurous acid), pH about 2–3.
• SO₃: SO₃ + H₂O → H₂SO₄, pH about 0–1.

Acid–base behaviour of the oxides and hydroxides

Basic — Na₂O, MgO, NaOH and Mg(OH)₂ react with acids to form salts and water:

Amphoteric — aluminium oxide and aluminium hydroxide react with both acids and bases:

The product with sodium hydroxide is sodium aluminate, containing the [Al(OH)₄]⁻ ion.

Acidic — P₄O₁₀, SO₂ and SO₃ react with bases:

Silicon dioxide is also acidic, but only reacts with hot, concentrated alkali: SiO₂ + 2NaOH → Na₂SiO₃ + H₂O.

The chlorides with water

• NaCl simply dissolves; the ions are hydrated. pH 7.
• MgCl₂ dissolves, and the small, 2+ Mg ion attracts water enough for very slight hydrolysis: pH about 6.5.
• AlCl₃ is hydrolysed. With a little water it fumes (HCl gas); in excess water it gives an acidic solution, pH about 3: AlCl₃ + 3H₂O → Al(OH)₃ + 3HCl. The very small, highly charged Al³⁺ ion pulls electron density from its water ligands so strongly that they release H⁺.
• SiCl₄ is hydrolysed completely, violently, with white fumes of HCl and a white solid: SiCl₄ + 2H₂O → SiO₂ + 4HCl. pH about 1–2.
• PCl₅ is hydrolysed violently: PCl₅ + 4H₂O → H₃PO₄ + 5HCl. pH about 1.

Explaining the trends

Across the period the difference in electronegativity between the element and oxygen or chlorine decreases, so the bonding changes from ionic to covalent:

• Ionic oxides (Na₂O, MgO) contain the O²⁻ ion, which is a strong base: O²⁻ + H₂O → 2OH⁻. So they are basic.
• Covalent oxides of non-metals (P₄O₁₀, SO₂, SO₃) react with water to form acids. So they are acidic.
• Aluminium oxide sits in between: ionic, but with its bonding given considerable covalent character by the small, highly charged Al³⁺ ion — hence amphoteric.
• Ionic chlorides (NaCl, MgCl₂) dissolve without reacting. Covalent chlorides (AlCl₃, SiCl₄, PCl₅) are hydrolysed by water, releasing HCl.

The same pattern works in reverse as evidence. A chloride with a high melting point that dissolves to a neutral solution and conducts when molten is ionic; one that is a volatile liquid, fumes in moist air and gives an acidic solution is covalent.

9.3 Periodicity beyond period 3

Elements in the same group have the same outer-electron arrangement, so they behave similarly, and the trends down a group (atoms larger, lower ionisation energy, lower electronegativity) let you predict how strongly. You may be asked to predict an element’s properties: germanium, below silicon in Group 14, should form GeO₂ and GeCl₄, with GeCl₄ a covalent liquid hydrolysed by water, and should be a semiconductor. Or to work the other way and deduce the identity of an unknown element from data — its oxide’s pH, its chloride’s volatility, the formulas of its compounds (which give its maximum oxidation number, and so its group), and its ionisation energies.