Inorganic chemistry

Explanation

• An element is classified as s, p, d or f block from its position in the Periodic Table, and that position is determined by proton number.
• The block corresponds to the subshell receiving the differentiating electron: an outer $s$ subshell gives s block, an outer $p$ subshell gives p block, and filling a $(n-1)d$ subshell gives d block.
• Electron configuration therefore explains classification, but any filled inner subshell must not override the differentiating subshell.
• For example, $\mathrm{[Ar]3d^{10}4s^2 4p^4}$ is p block despite containing occupied d orbitals.

Explanation

• Across Period 3 from Na to Ar, atomic radius generally decreases because proton number rises while added electrons enter the same principal shell with similar shielding.
• First ionisation energy generally increases, but falls from Mg to Al because Al loses a higher-energy $3p$ electron, and from P to S because repulsion within a paired $3p$ orbital eases removal.
• Melting points require structure and bonding: metallic bonding strengthens from Na to Al, giant covalent silicon is very high, while $\mathrm{P_4}$, $\mathrm{S_8}$ and $\mathrm{Cl_2}$ molecules and monatomic Ar melt by overcoming London forces whose strength depends on electron-cloud size.

Explanation

• From Mg to Ba, atomic radius increases and first ionisation energy decreases as each element adds an electron shell. All have giant metallic structures; melting-point changes reflect metallic-bond strength and lattice packing rather than a simple monotonic trend.
• Reactions with water become more vigorous, forming $\mathrm{M(OH)_2}$ and $\mathrm{H_2}$, while Mg reacts with steam to form MgO.
• Magnesium extracts titanium from $\mathrm{TiCl_4}$.
• Hydroxide solubility increases down the group, whereas sulfate solubility decreases: $\mathrm{Mg(OH)_2}$ and $\mathrm{MgSO_4}$ have medicinal uses, $\mathrm{Ca(OH)_2}$ treats acidic soil, CaO or $\mathrm{CaCO_3}$ removes flue-gas $\mathrm{SO_2}$, and insoluble $\mathrm{BaSO_4}$ is a safe contrast medium.
• Sulfate testing uses acidified $\mathrm{BaCl_2}$.

Explanation

• Down Group 7, electronegativity decreases as atomic radius and shielding increase, while boiling point rises because larger $\mathrm{X_2}$ electron clouds form stronger London forces.
• Halogen oxidising ability decreases, so a halogen displaces halide ions below it; halide reducing ability increases.
• With concentrated sulfuric acid, chloride shows acid–base behaviour, bromide reduces sulfur mainly to $\mathrm{SO_2}$ and iodide can reduce it to sulfur or $\mathrm{H_2S}$.
• Acidified $\mathrm{AgNO_3}$ distinguishes halides: AgCl is white and dissolves in dilute ammonia, AgBr is cream and dissolves in concentrated ammonia, and AgI is yellow and insoluble.
• Nitric acid avoids adding halide ions.

Explanation

• Chlorine disproportionates in water to chloride and chlorate(I): $\mathrm{Cl_2+H_2O\rightleftharpoons2H^++Cl^-+ClO^-}$. In light, chlorine water can instead produce chloride ions and oxygen overall.
• Cold, dilute aqueous NaOH gives $\mathrm{NaCl}$ and $\mathrm{NaClO}$ for bleaching and disinfection.
• Chlorination kills pathogens and supplies residual protection; controlled health benefits outweigh toxic effects, though risks require assessment.
• Decisions about adding chemicals such as fluoride also balance evidence.
• Required practical 4 reliably identifies Group 2 and $\mathrm{NH_4^+}$ cations and halide, $\mathrm{OH^-}$, $\mathrm{CO_3^{2-}}$ and $\mathrm{SO_4^{2-}}$ anions by test-tube reactions.

Explanation

• Melting points must be explained by structure and bonding, never by 'stronger bonds' alone. Na₂O and MgO are ionic with giant lattices, so their melting points are high, and MgO is higher than Na₂O because Mg²⁺ has a greater charge and a smaller radius, giving stronger electrostatic attraction.
• Al₂O₃ is ionic with covalent character. SiO₂ is macromolecular, or giant covalent, so melting it breaks many strong covalent bonds and its melting point is very high.
• P₄O₁₀, SO₂ and SO₃ are simple molecular, so only weak van der Waals forces between molecules are overcome and their melting points are low. Sodium reacts vigorously with cold water to give sodium hydroxide and hydrogen; magnesium reacts very slowly with cold water but rapidly with steam to give magnesium oxide and hydrogen.
• Na₂O dissolves to give NaOH at about pH 13-14; MgO is sparingly soluble and gives roughly pH 9; Al₂O₃ and SiO₂ are insoluble and give pH 7, though Al₂O₃ is amphoteric, dissolving in both acids and bases; P₄O₁₀ gives phosphoric(V) acid H₃PO₄ at pH 0-1; and SO₂ and SO₃ give sulfurous and sulfuric acids, SO₃ giving the lower pH.
• The mark scheme is strict about attributing structures to the oxides rather than to the elements.

Explanation

• For Ti to Cu, transition-metal characteristics arise from an incomplete d sub-level in atoms or ions. These characteristics are complex formation, coloured ions, variable oxidation states and catalytic activity.
• A ligand is a molecule or ion that donates an electron pair to the central metal atom or ion, forming a co-ordinate bond.
• A complex consists of that central metal surrounded by ligands.
• Co-ordination number counts co-ordinate bonds to the centre, not the number of ligand particles: a bidentate ligand supplies two donor atoms.
• Applying the definition excludes any d-block element whose relevant ions are only $d^0$ or $d^{10}$.

Explanation

• $\mathrm{H_2O}$, $\mathrm{NH_3}$ and $\mathrm{Cl^-}$ are monodentate ligands. Similar-sized uncharged water and ammonia exchange without changing co-ordination number, although substitution may stop at $\mathrm{[Cu(NH_3)_4(H_2O)_2]^{2+}}$.
• Larger chloride can replace water with a co-ordination-number change.
• Ethane-1,2-diamine and ethanedioate are bidentate; EDTA is multidentate.
• Their replacement of monodentate ligands is the chelate effect: metal–ligand bond enthalpy changes broadly balance, while producing more particles makes $\Delta S$ positive.
• Haem is an Fe(II) multidentate complex that binds oxygen; carbon monoxide is toxic because it substitutes for oxygen and prevents transport.

Explanation

• Small ligands such as water and ammonia commonly form six-coordinate octahedral complexes; larger chloride ligands commonly form four-coordinate tetrahedral complexes.
• Four-coordinate complexes may instead be square planar, while $\mathrm{[Ag(NH_3)_2]^+}$ in Tollens' reagent is linear and two-coordinate.
• Octahedral complexes with suitable monodentate ligands and square-planar complexes can show cis–trans isomerism; cisplatin is the cis form of $\mathrm{[Pt(NH_3)_2Cl_2]}$.
• Octahedral complexes with bidentate ligands can show optical isomerism when their two arrangements are non-superimposable mirror images.
• Drawings must preserve three-dimensional geometry, not merely change page orientation.

Explanation

• A transition-metal ion appears coloured because it absorbs selected visible wavelengths and transmits or reflects the remainder. Absorption promotes a d electron from a ground state to an excited state, with the energy gap $\Delta E=h\nu=hc/\lambda$.
• Changing ligand, oxidation state or co-ordination number changes this gap and therefore the absorbed wavelength and observed colour.
• Spectroscopy uses visible absorption quantitatively.
• In colorimetry, a blank sets zero, a suitable filter or wavelength maximises useful absorption, and standards provide an absorbance–concentration calibration graph.
• An unknown concentration is interpolated only within the calibrated range, with any preparation dilution reversed afterwards.

Explanation

• Transition elements show variable oxidation states. Zinc reduces vanadate(V) in acid through vanadium(IV), vanadium(III) and vanadium(II), producing characteristic colour changes.
• A transition-metal redox potential depends on pH and ligand because these conditions change the relative stability of oxidation states.
• Tollens' reagent, $\mathrm{[Ag(NH_3)_2]^+}$, is reduced to metallic silver by an aldehyde and distinguishes aldehydes from ketones.
• Quantitative redox titrations include $\mathrm{Fe^{2+}}$ and $\mathrm{C_2O_4^{2-}}$ with acidified $\mathrm{MnO_4^-}$; balanced half-equations determine ratios used for iron-tablet, steel, hydrated-salt, ethanedioic-acid or hydrogen-peroxide calculations.
• Accurate aliquot scaling is essential.

Explanation

• A heterogeneous catalyst is in a different phase from reactants; reaction occurs at surface active sites through adsorption, reaction and desorption. A support maximises surface area and reduces cost, while poisons block sites and lower efficiency.
• Iron catalyses the Haber process and $\mathrm{V_2O_5}$ the Contact process.
• A homogeneous catalyst shares the reactant phase and forms an intermediate.
• Variable oxidation states allow $\mathrm{Fe^{2+}/Fe^{3+}}$ to catalyse iodide–peroxodisulfate through two electron-transfer steps and $\mathrm{Mn^{2+}}$ to autocatalyse ethanedioate–manganate(VII) through manganese intermediates.
• Catalyst equations must show the intermediate oxidation state and regenerate the catalyst so it cancels from the overall reaction.

Explanation

• The higher charge density of the 3+ ion polarises the O-H bonds of the coordinated water more strongly, so a proton is released more readily and the solution is more acidic. Test-tube reactions with three bases are examinable, and each behaves differently.
• Sodium hydroxide and ammonia both act as bases and remove protons, giving the neutral hydroxide precipitate M(H₂O)₃(OH)₃ or M(H₂O)₄(OH)₂: Fe(II) gives a green precipitate, Cu(II) a blue precipitate, Fe(III) a brown precipitate and Al(III) a white one. In excess sodium hydroxide only the aluminium hydroxide dissolves, because it is amphoteric and forms [Al(OH)₄]⁻;
• in excess ammonia only the copper hydroxide dissolves, forming the deep blue [Cu(NH₃)₄(H₂O)₂]²⁺ by ligand substitution.
• Sodium carbonate distinguishes the 2+ ions from the 3+ ions: the weakly acidic 2+ aqua ions merely form the metal carbonate precipitate, whereas the strongly acidic 3+ aqua ions are acidic enough to release carbon dioxide, so effervescence accompanies the hydroxide precipitate.
• Every equation must be balanced with the full complex ion formulae, and every answer needs a colour and a state. This is Required practical 11.