The Periodic Table

Across this period the inner [Ne] core stays the same while nuclear charge increases and electrons enter the third shell. Down Group 2 from Mg to Ba, the outer configuration remains ns². Down Group 17 from Cl to I, it remains ns²np⁵. The value of n increases down each group, so the outer electrons occupy more distant shells.

Use these patterns to predict typical ions: Group 2 loses two electrons to give M²⁺; Group 17 gains one to give X^-. This does not mean that every compound consists of ions. Period-3 non-metals also share electrons in covalent compounds.

The ionisation-energy decrease from Mg to Al occurs because Al loses a higher-energy 3p electron, whereas Mg loses a 3s electron. The decrease from P to S occurs because S has a paired 3p electron; repulsion within that pair makes removal easier than from P's three singly occupied p orbitals. A general trend is not a monotonic rule.

Cations are usually smaller than their parent atoms: electron loss reduces repulsion and may remove the outer shell entirely. Anions are larger than their parent atoms because added electrons increase repulsion while nuclear charge is unchanged. Comparing isoelectronic ions isolates nuclear charge: Na⁺, Mg²⁺ and Al³⁺ each have ten electrons, and radius decreases in that order.

Down Group 2, M²⁺ radii increase because the ions have progressively more occupied shells. The same applies to Cl^-, Br^- and I^-. Always state whether a graph shows atomic radius or the radius of a particular ion.

Among the specified molecular forms, S₈ has a larger, more polarizable electron cloud than P₄ or Cl₂ and generally stronger instantaneous dipole-induced dipole attractions. Hence its melting point is higher. State the allotrope: red or black phosphorus cannot be explained by pretending it consists of white-phosphorus P₄ molecules.

The highest positive oxidation numbers in these oxides rise with the number of valence electrons available for bonding. These are formal oxidation numbers, not necessarily actual ionic charges: SiO₂ is not a lattice of bare Si⁴⁺ and O^2- ions.

Across the period, increasing element electronegativity reduces the electronegativity difference from oxygen. Bonding changes from predominantly ionic towards covalent. Oxygen in an ionic basic oxide can accept protons; covalent non-metal oxides instead commonly form oxoacids or react with bases. Silicon dioxide remains insoluble because its extended network is difficult to disrupt, despite being classified as acidic.

NaOH is a soluble strong base. Mg(OH)₂ is basic but sparingly soluble: limited solubility should not be confused with partial dissociation of the small amount that dissolves. Both react with acids. Al(OH)₃ is amphoteric and can dissolve in either acid or excess aqueous sodium hydroxide.

The familiar acidic classification of SiO₂ refers to its reaction with bases, not to dissolution in water. Likewise, no observable reaction with dilute hydrochloric acid is not evidence that a solid is neutral. Distinguish solubility, speed of reaction and acid-base character.

The general ionic-to-covalent trend follows decreasing electronegativity difference between the period-3 element and chlorine. AlCl₃ requires special care: the small, highly charged aluminium centre strongly polarises the chloride electron cloud, giving substantial covalent character. A simple electronegativity threshold is not a complete explanation.

When AlCl₃ enters excess water, aluminium becomes hydrated. The highly charged central ion withdraws electron density from the O-H bonds of coordinated water, allowing proton release: [Al(H₂O)₆]³⁺ + H₂O ⇌ [Al(H₂O)₅(OH)]²⁺ + H₃O⁺. Chloride is not the source of the proton. Do not assume that dissolution automatically precipitates all aluminium as Al(OH)₃.

In SiCl₄ and PCl₅, chlorine attracts the bonding electrons, leaving an electron-poor central atom. Water can donate an oxygen lone pair to that centre; replacement of chlorine by oxygen-containing groups and proton transfer release HCl. This is chemical change of the chloride, whereas dissolving NaCl separates existing ions. Mg²⁺ has lower charge density than Al³⁺, so it weakens coordinated O-H bonds less strongly and its hydrated ion releases protons less readily.

Low boiling temperature or volatility supports a discrete molecular structure, whereas high melting temperature and molten conduction support an ionic lattice. Water reactions provide different evidence: vigorous hydrolysis of a molecular chloride does not mean it was an ionic salt before water was added. Name the state when discussing PCl₅ or AlCl₃.

For M²⁺(aq) + 2e^- ⇌ M(s), a more negative standard reduction potential means the metal is more readily oxidised in a suitable paired reaction. Group 2 metals are reducing agents because they lose electrons. The standard potentials become generally more negative from Mg towards Ba, supporting greater reducing power down the group.

For X₂ + 2e^- ⇌ 2X^-(aq), a more positive standard reduction potential means the halogen is a stronger oxidising agent. The usual values are approximately +1.36 V for chlorine, +1.07 V for bromine and +0.54 V for iodine, with the standard physical states specified in the data. Oxidising strength therefore decreases Cl₂ > Br₂ > I₂.

Group 2 carbonates decompose on heating: MCO₃(s) → MO(s) + CO₂(g). Down Mg to Ba, the cation radius increases while its charge remains +2. Charge density and polarising power decrease. The larger cation distorts the large carbonate ion less, so the carbonate is harder to decompose and its thermal stability increases.

For hydrogen halides, the relevant reaction is 2HX(g) ⇌ H₂(g) + X₂(g). From HCl to HBr to HI, the halogen atom becomes larger, the H-X bond becomes longer and the bond energy decreases. Thermal stability therefore decreases in that order: HI decomposes most readily.

When practical evidence is supplied, distinguish observation from interpretation. "A white precipitate dissolves in excess alkali" is an observation; "the hydroxide is amphoteric" is an interpretation. Use the stated temperature, physical state and reagent excess. If two identities fit, name the additional evidence needed instead of forcing a unique answer.

Remember the exceptions and conditions: Mg/Al and P/S ionisation-energy dips; the ionic-radius jump between cation and anion series; white-phosphorus P₄; AlCl₃ covalent character; phase-dependent PCl₅; partial hydrated-ion hydrolysis rather than automatic hydroxide precipitation.