9. The Periodic Table: chemical periodicity

Across Period 3, properties such as atomic radius, melting point, boiling point and electrical conductivity change as nuclear charge, bonding and structure change.

The trend is not one uniform line: metallic bonding dominates Na–Al, giant covalent silicon produces a very high melting point, and simple molecular P₄/S₈/Cl₂/Ar are governed by intermolecular forces.

Silicon has a much higher melting point than phosphorus because it forms a giant covalent lattice, whereas phosphorus exists as discrete P₄ molecules.

Do not explain the whole period using only increasing proton number. The type of structure and bonding changes at the same time.

Melting points rise when stronger attractions must be overcome. Electrical conductivity requires mobile charged particles or delocalised electrons.

Na, Mg and Al generally strengthen metallic bonding as charge density increases; silicon forms a giant covalent network; molecular elements have lower melting points and limited conductivity.

Al conducts as a metal, silicon conducts less effectively as a semiconductor-like giant covalent solid, and sulfur does not conduct because its S₈ molecules have no mobile charge carriers.

A high melting point does not guarantee conductivity. Silicon and many ionic solids are strong lattices but lack freely moving charge carriers in the solid state.

Period 3 elements form oxides with different structures and acid–base behaviour. Metallic oxides are generally ionic/basic, while non-metal oxides are molecular/covalent and more acidic.

Use oxidation states and the balanced equation to compare products. The change from ionic to covalent bonding across the period explains the shift in oxide behaviour.

Na₂O reacts with water to form NaOH, whereas SO₂ dissolves to form an acidic solution. Silicon dioxide is a giant covalent oxide and does not behave like a simple molecular oxide.

Do not label every oxide simply “basic” or “acidic” without considering structure and whether it reacts with water. Amphoteric cases need their own evidence.

Oxidation number helps assign the formal oxidation state of the Period 3 element in its oxide. Oxygen is usually −2, so the total must match the neutral compound’s charge.

Use the formula and neutrality to calculate the oxidation number, then connect it to the element’s outer-electron arrangement and the possible oxide chemistry.

In SO₃, three oxygens contribute −6, so sulfur is +6. In Na₂O, oxygen is −2 and each sodium is +1.

Oxidation number is bookkeeping, not proof that covalent molecules contain isolated ions. Do not confuse formal oxidation state with actual charge distribution.

A chemical-periodicity comparison is only meaningful when the equations, states and products are specified. Balance atoms and charge before interpreting the trend.

Compare like reactions—such as element plus oxygen or oxide plus water—then use structure and oxidation state to explain why behaviour changes across the period.

4Na + O₂ → 2Na₂O and Si + O₂ → SiO₂ are both balanced, but their products differ from ionic/basic to giant covalent because the underlying bonding changes.

Do not infer a general reaction from one oxide or one state. Periodic trends describe families of related cases, not every possible condition.

Across Period 3, oxides and hydroxides generally change from basic on the left to acidic on the right as bonding becomes more covalent and the central atom attracts oxygen more strongly.

Basic oxides react with acids; acidic oxides react with bases or water to form acids. Amphoteric substances such as Al₂O₃ can react in both directions.

MgO reacts with HCl, while SO₃ reacts with water to form H₂SO₄. Al₂O₃ reacts with both acids and strong bases.

The trend is not simply “metal to non-metal”. Silicon dioxide is a giant covalent solid and its reaction behaviour must be considered separately from molecular oxides.

Period 3 chlorides range from ionic lattices to covalent molecules. Their reaction with water and the pH of the resulting solution depend on the chloride’s bonding and the polarising power of the central atom.

Ionic chlorides often dissolve to give neutral ions, while covalent chlorides hydrolyse to form acidic solutions and hydrogen chloride.

NaCl dissolves without hydrolysis, whereas PCl₃ reacts with water to form H₃PO₃ and HCl. Aluminium chloride shows strong polarisation and acidic aqueous behaviour.

Do not assume every chloride gives a neutral solution. Ask whether the chloride is ionic or covalent and whether hydrolysis occurs.

Periodicity questions require comparing like properties across related compounds: structure, physical state, hydrolysis, acid–base behaviour and bonding.

Start with the formula and bonding model, then write an equation for the stated reagent or water. The trend becomes meaningful only after the individual cases are identified.

Na₂O is ionic/basic, SiO₂ is giant covalent, and SO₂ is molecular/acidic. Their different structures explain why the same “oxide” label does not imply one behaviour.

Do not rank compounds from the element’s position alone. The actual oxide or chloride formula and structure control the observed reaction.

Period 3 compounds can contain ionic, simple covalent or giant covalent structures. The bonding type predicts melting point, conductivity, solubility and reaction behaviour when combined with particle mobility.

Name the structure first, then explain the property through the dominant attraction and whether charged particles or electrons can move.

NaCl has an ionic lattice and conducts when molten; SiO₂ is a giant covalent network with a high melting point; Cl₂ is a simple molecular substance with weak intermolecular forces.

The word “compound” does not imply a single structure type. Compare the actual particles and network before predicting a property.

The characteristic properties of an element arise from its electron configuration, bonding model and the structures it forms. Position in the periodic table is a starting clue, not a complete explanation.

Use effective nuclear charge and shell structure for atomic trends; use lattice, molecular or metallic models for bulk properties.

A Group 2 metal conducts because of delocalised electrons, while a molecular Group 17 element has weak intermolecular forces and a low boiling point.

Do not mix atomic trends with compound trends. A metal’s atomic radius and its oxide’s acidity are related through different levels of explanation.

An unknown element can be located by combining periodic-table position clues with measured properties such as ionisation energy, atomic radius, bonding or reactions.

First identify the likely group/period from the strongest evidence, then check whether the proposed element’s structure and reactions fit. Do not force one measurement to decide everything.

A large jump after two successive ionisation energies suggests two valence electrons; if the element also forms a basic oxide and conducts as a solid, a Group 2 metal is plausible.

A single property is rarely unique. Use multiple independent clues and state what each one supports or rules out.