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

• Na and Mg react with water.
• Burning Period 3 elements gives the specified $\mathrm{Na_2O}$, MgO, $\mathrm{Al_2O_3}$, $\mathrm{SiO_2}$, $\mathrm{P_4O_{10}}$, $\mathrm{SO_2}$ and $\mathrm{SO_3}$ chemistry.
• Highest-oxide melting points reflect a change from giant ionic lattices through giant covalent $\mathrm{SiO_2}$ to molecular phosphorus and sulfur oxides.
• With water, ionic oxides form alkaline solutions while $\mathrm{P_4O_{10}}$, $\mathrm{SO_2}$ and $\mathrm{SO_3}$ form $\mathrm{H_3PO_4}$, $\mathrm{H_2SO_3}$ and $\mathrm{H_2SO_4}$ and their phosphate, sulfite and sulfate anions.
• Basic oxides react with acids, acidic oxides with bases, and amphoteric $\mathrm{Al_2O_3}$ with both.

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 specified aqua ions are $\mathrm{[Fe(H_2O)_6]^{2+}}$, $\mathrm{[Cu(H_2O)_6]^{2+}}$, $\mathrm{[Al(H_2O)_6]^{3+}}$ and $\mathrm{[Fe(H_2O)_6]^{3+}}$. Their hydroxide precipitates are green, blue, white and brown respectively.
• A $3+$ aqua ion is more acidic than a comparable $2+$ ion because greater charge-to-size ratio polarises ligand O–H bonds more strongly.
• Excess hydroxide dissolves amphoteric aluminium hydroxide; excess ammonia dissolves copper(II) hydroxide as a deep-blue ammine complex.
• Carbonate precipitates metal carbonates from the $2+$ ions but deprotonates acidic $3+$ aqua ions, giving hydroxide precipitate and $\mathrm{CO_2}$.
• Required practical 11 identifies transition-metal ions by these test-tube reactions.