Gaokao Chemistry: Structure, bonding, materials and inorganic chemistry

Atomic Structure and the Periodic Table

This independent foundation chapter develops how the quantum-mechanical model of the atom explains the layout of the periodic table and the trends you are asked to predict: electron configuration, ionisation energy, atomic and ionic radius, electronegativity, and the evidence available from mass spectra and successive ionisation energies.

Subatomic structure and isotopes

Atomic number Z = protons; mass number A = protons + neutrons; isotopes differ only in neutrons, ions only in electrons.
Relative atomic mass is the abundance-weighted mean of the isotope masses, not the mass number of any single isotope.
A mass spectrum gives m/z peaks with relative abundances; an identified intact singly charged molecular ion can support Mr. In a suitable small organic molecule, the M+1/M intensity ratio has an approximate contribution of 1.1 percentage points per carbon from carbon-13, but other isotopes and fragmentation must be considered.

Electron configuration

Fill by Aufbau order, obey the Pauli exclusion principle and Hund's rule (maximise unpaired spins within a subshell).
4s fills before 3d, but when transition metals ionise the 4s electrons are removed first.
Cr [Ar]3d5 4s1 and Cu [Ar]3d10 4s1 are the standard exceptions to expect.
Blocks are named for the subshell receiving the last electron.

Periodic trends

Across a period, nuclear charge rises while shielding is roughly constant, so effective nuclear charge increases: radius falls, ionisation energy and electronegativity rise.
Down a group, extra shells increase shielding and distance: radius rises, ionisation energy falls.
Successive ionisation energies jump sharply at shell boundaries; a large jump after the second value means two outer electrons, so Group 2.

Worked example

Rubidium is 72.2% 85Rb and 27.8% 87Rb. Ar = (85 × 0.722) + (87 × 0.278) = 61.37 + 24.19 = 85.56. Now an element with first ionisation energies 738, 1451 and 7733 kJ/mol: the third value is about five times the second, so two electrons are removed easily and the third comes from an inner shell - a Group 2 element.

Traps that cost marks

Quoting a periodic-table atomic mass as the mass number of a particular isotope, or forgetting to weight by abundance.
Defining first ionisation energy without "gaseous atoms" and "one mole".
Assuming Al has a higher first ionisation energy than Mg simply because it lies farther across the period; the measured Al value is lower, consistent with removing its higher-energy 3p electron.
Writing an electron configuration for a transition-metal ion without emptying 4s first.
Calling Zn a typical transition metal; its common ion is d10, so it lacks the characteristic variable oxidation states and colour.

Chemical Bonding and Molecular Structure

This independent foundation chapter develops how electrons are shared or transferred, how the resulting Lewis structure is converted into a three-dimensional shape by VSEPR theory, and how bonding and shape together explain polarity and physical properties.

Bond types

Ionic bonding transfers electrons to form a lattice; melting points are high and conduction requires mobile charge carriers, commonly in a melt or in aqueous solution when the compound dissolves. Lattice strength rises with ionic charge and falls with ion radius.
Covalent bonding shares electron pairs. Increasing bond order shortens the bond and increases bond enthalpy, so C≡C is shorter and stronger than C=C.
Metallic bonding uses delocalised electrons: malleable, ductile, good conductor. A dative bond (as in NH4+) becomes an ordinary covalent bond once formed.
Resonance delocalises electrons over equivalent positions, giving bond orders intermediate between the drawn structures, as in CO3^2− and benzene.

VSEPR shapes

Count electron domains around the central atom; double and triple bonds count as one domain, lone pairs count as one.
Repulsion order is lone pair–lone pair > lone pair–bond pair > bond pair–bond pair, so lone pairs compress adjacent angles.
2 domains linear (180°), 3 trigonal planar (120°), 4 tetrahedral (109.5°), 5 trigonal bipyramidal, 6 octahedral. In a trigonal bipyramid, lone pairs occupy equatorial positions.

Polarity

In the localized bond-dipole model, determine molecular polarity from the vector sum and actual charge distribution, not from the mere presence of polar bonds. CO2 and CCl4 cancel; H2O, NH3 and SO2 do not.
Electronegativity difference indicates bond polarity; molecular shape decides whether the dipoles add.

Worked example

SF4: sulfur has six valence electrons, four are used in S–F bonds leaving one lone pair, so there are 5 domains - trigonal bipyramidal electron geometry. The lone pair takes an equatorial site, giving a seesaw shape with F–S–F angles distorted from ideal trigonal-bipyramidal values. VSEPR predicts the seesaw arrangement qualitatively; exact angles require experimental or calculated structural data.

Traps that cost marks

Forgetting the lone pairs on the central atom when counting domains.
Expanding the octet on a period 2 atom; only period 3 and below can do this.
Confusing electron-pair geometry with molecular shape when a lone pair is present.
Assuming a molecule with polar bonds is automatically polar.
Describing ionic compounds as discrete molecules rather than a lattice.

States of Matter and Intermolecular Forces

This independent foundation chapter develops the gas laws, the assumptions and failures of the ideal-gas model, phase behaviour, and the ranking of melting and boiling points from the strength of intermolecular forces.

Gas behaviour

Ideal gas equation PV = nRT with P in Pa, V in m3, T in kelvin and R = 8.314 J mol−1 K−1.
Dalton's law: total pressure is the sum of partial pressures; each partial pressure is mole fraction × total pressure.
Kinetic theory assumes negligible molecular volume, no intermolecular forces, and elastic collisions. Real gases deviate most at high pressure and low temperature, when molecules are close together.
van der Waals equation (P + an²/V²)(V − nb) = nRT corrects for attraction (a) and molecular volume (b).

Intermolecular forces

London dispersion forces act between all particles and strengthen with more electrons and greater surface contact; they are the only force in non-polar molecules.
Dipole–dipole forces act between permanent dipoles. Hydrogen bonding is the strong special case requiring H directly bonded to N, O or F.
Ion–dipole forces operate in solutions of ionic compounds.
For otherwise comparable substances, stronger attractions often raise boiling point and enthalpy of vaporisation and lower vapour pressure. Melting also depends strongly on crystal packing and symmetry, so it is not ranked by attraction strength alone.

Phase behaviour

Vapour pressure rises with temperature; a liquid boils when its vapour pressure equals the external pressure.
A phase diagram shows the triple point and critical point; water's solid–liquid line slopes backwards, so ice melts under pressure.

Worked example

0.575 g of a volatile liquid gives 200 cm3 of vapour at 100 °C and 101 kPa. n = PV/RT = (101000 × 2.00 × 10−4)/(8.314 × 373) = 20.2/3101 = 6.51 × 10−3 mol. M = 0.575/6.51 × 10−3 = 88.3 g mol−1, consistent with a small ester such as C4H8O2.

Traps that cost marks

Substituting degrees Celsius or cubic centimetres directly into PV = nRT.
Claiming hydrogen bonding whenever hydrogen is present; it needs an N–H, O–H or F–H bond.
Forgetting that HF boils above HCl despite being lighter, because of hydrogen bonding.
Assuming ice is denser than water; the open hydrogen-bonded lattice makes it less dense.
Ignoring that dispersion forces increase with chain length and branching reduces them.

Inorganic Chemistry and Periodicity

This independent foundation chapter develops the descriptive chemistry that distinguishes groups and periods: s-block trends, halogen behaviour, the characteristic properties of transition metals, and the observations and equations you are expected to recall rather than derive.

Periodic trends

Group 2 reactivity with water increases down the group as ionisation energy falls. Hydroxide solubility increases down the group while sulfate solubility decreases, which is why BaSO4 is used as a precipitate.
Thermal stability of the usual Group 2 carbonates and nitrates generally increases down the group because the larger cation distorts the anion less.
Halogen oxidising power decreases down the group; halide reducing power increases. Hydrogen halides become less thermally stable and their aqueous solutions more acidic down the group.
Disproportionation, for example Cl2 + 2NaOH → NaCl + NaOCl + H2O.

Transition metals

Under the IUPAC Gold Book definition, a transition element has an atom with an incomplete d subshell or can form cations with one. Sc has an atomic 3d electron and qualifies under that definition; zinc is conventionally excluded because its atom and common ion retain a filled d subshell.
Characteristic properties: variable oxidation states (Mn from +2 to +7), catalytic activity, coloured complex ions and paramagnetism.
Ligands donate a lone pair to form coordinate bonds. Coordination numbers are commonly 6 (octahedral) and 4 (tetrahedral or square planar).
Ligand-field d–d transitions can produce colour where suitable d states are available. Simple aqueous Sc3+ and Zn2+ lack such transitions, but colour is not determined by d count alone: charge-transfer transitions can make other d0 or d10 complexes coloured.

Worked example

In [Fe(CN)6]3−, six cyanide ligands at −1 give −6; the overall charge is −3, so iron is +3. Fe3+ is d5. Cyanide is a strong-field ligand, so the configuration is low-spin t2g5 with one unpaired electron - the ion is paramagnetic and coloured.

Traps that cost marks

Reversing the hydroxide and sulfate solubility trends down Group 2.
Calling all d-block elements transition metals, including Zn without checking the atom-or-cation definition. Sc itself has an incomplete atomic d subshell.
Assuming every transition-metal ion is coloured.
Quoting a qualitative observation without a balanced equation.
Forgetting that halide reducing power increases, not decreases, down the group.