Physical chemistry

Explanation

• Atomic models changed when new experimental evidence could not be explained by an earlier model.
• The accepted model has a very small nucleus containing protons and neutrons, with electrons occupying the surrounding space.
• A proton has relative charge $+1$ and relative mass $1$; a neutron has charge $0$ and relative mass $1$; an electron has charge $-1$ and relative mass about $1/1836$.
• A neutral atom has equal numbers of protons and electrons.
• An ion forms only by losing or gaining electrons, so its nucleus and element identity remain unchanged.

Explanation

• You must be able to determine the number of fundamental particles in atoms and ions from the mass number, the atomic number and the charge:
• neutrons are A minus Z, and electrons equal Z adjusted for the charge. The specification then builds the time-of-flight mass spectrometer on top of this, limited to five stages you should be able to name in order: ionisation, acceleration to give all ions the same kinetic energy, ion drift, ion detection, and data analysis.
• Ionisation is by electron impact or by electrospray; electron impact knocks an electron out, which is written as X(g) gives X⁺(g) + e⁻. Because acceleration gives every ion the same kinetic energy, heavier ions move more slowly and take longer to drift the flight tube, so time of flight increases with mass.
• The mass spectrometer gives accurate relative isotopic masses and relative abundances, which lets you calculate a relative atomic mass as a weighted mean: multiply each isotopic mass by its abundance, add, and divide by the total abundance.
• The same instrument identifies elements and determines relative molecular mass from the molecular ion peak.

Explanation

• For atoms up to $Z=36$, electrons fill $1s$, $2s$, $2p$, $3s$, $3p$, $4s$, $3d$, then $4p$. Chromium is $\mathrm{[Ar]3d^5\,4s^1}$ and copper is $\mathrm{[Ar]3d^{10}\,4s^1}$; transition-metal ions lose $4s$ electrons before $3d$.
• First ionisation energy removes one electron from each atom in one mole of gaseous atoms.
• Across Period 3, increasing nuclear charge generally raises first ionisation energy, with sub-shell and electron-pairing anomalies.
• Down Group 2, increasing distance and shielding lower it.
• A large successive-ionisation jump shows removal from a shell closer to the nucleus and provides evidence for shell structure.

Explanation

• Relative atomic mass, $A_r$, is the weighted mean mass of an atom of an element relative to one twelfth of the mass of a carbon-12 atom. Relative molecular mass, $M_r$, is the mean mass of a molecule on the same scale.
• Both are ratios and therefore have no units.
• Calculate $M_r$ by multiplying each element's $A_r$ by the number of its atoms and summing.
• For an ionic substance, the equivalent quantity is relative formula mass because the solid contains a lattice rather than separate molecules.
• Brackets and waters of crystallisation must be counted fully.

Explanation

• One mole contains the Avogadro constant, $N_A$, of specified entities, such as atoms, molecules, ions, electrons or formula units.
• Amount of substance links measurable quantities to particle number: $n=m/M$, $N=nN_A$, and for a solution $n=cV$ when $V$ is in $\mathrm{dm^3}$.
• A secure calculation first converts the given quantity to moles, applies the chemical formula or equation ratio, then converts to the requested quantity.
• The named entity matters: one mole of $\mathrm{MgCl_2}$ formula units contains two moles of chloride ions.

Explanation

• The ideal gas equation is $pV=nRT$.
• The gas constant $R$ has the value $8.31$ in units of $\mathrm{J\,K^{-1}\,mol^{-1}}$, so pressure must be in $\mathrm{Pa}$, volume in $\mathrm{m^3}$, temperature in $\mathrm{K}$ and amount in $\mathrm{mol}$.
• Convert $\mathrm{kPa}$ to $\mathrm{Pa}$ by multiplying by $10^3$, $\mathrm{cm^3}$ to $\mathrm{m^3}$ by multiplying by $10^{-6}$, and $\mathrm{dm^3}$ to $\mathrm{m^3}$ by multiplying by $10^{-3}$.
• Convert Celsius using $T/\mathrm{K}=\theta/^{\circ}\mathrm{C}+273$.
• Rearranging the equation before substitution reduces unit and algebra errors in multi-step calculations.

Explanation

• An empirical formula is the simplest whole-number ratio of atoms of each element; a molecular formula gives the actual numbers of atoms in one molecule.
• From masses or percentages, divide each value by the element's $A_r$ to obtain relative moles, divide every result by the smallest, then multiply all ratios by the same small integer if necessary.
• A ratio near $1.5$ requires doubling, not rounding.
• To obtain a molecular formula, calculate the empirical formula mass and find the integer $M_r/(\text{empirical formula mass})$.
• Combustion data must be converted from product moles to constituent atom moles.

Explanation

• Almost every quantitative question in 7405 routes through this point. You are expected to write balanced equations, full and ionic, for reactions you have studied, and to balance unfamiliar equations when the reactants and products are given.
• From a balanced equation you must be able to calculate masses, volumes of gases, percentage yields, percentage atom economies, and concentrations and volumes for reactions in solution. Percentage yield compares the mass actually obtained with the theoretical maximum from the limiting reactant, so the first step is always to identify which reactant limits the reaction.
• Atom economy, by contrast, is fixed by the equation alone and is unaffected by how well the reaction is carried out; a high atom economy means less waste, better use of raw materials and lower disposal costs, which is the economic, ethical and environmental argument the specification asks for.
• Gas volume questions use pV = nRT with SI units — pressure in Pa, volume in m³, temperature in K — and the commonest lost marks come from failing to convert cm³ to m³ or kPa to Pa.
• Report answers to the number of significant figures the least accurate datum allows.

Explanation

• Ionic bonding is the strong electrostatic attraction between oppositely charged ions throughout a giant lattice. Electron transfer forms the ions but is not itself the bond.
• Simple-ion charges can be predicted from Periodic Table groups, and the ionic formula uses the smallest whole-number ratio giving zero overall charge.
• Compound ions remain intact; for example, two $\mathrm{Al^{3+}}$ ions balance three $\mathrm{SO_4^{2-}}$ ions to give $\mathrm{Al_2(SO_4)_3}$.
• Brackets are required when more than one compound ion is present.
• Explanations of high melting point must refer to strong attractions acting in many directions through the lattice.

Explanation

• A covalent bond is a shared pair of electrons. Double and triple bonds contain two and three shared pairs.
• In an ordinary covalent bond, each bonded atom contributes one electron to the shared pair. In a co-ordinate, or dative covalent, bond, both electrons in the shared pair are supplied by one atom.
• The displayed-formula arrow starts at a lone pair on the electron-pair donor and points towards the electron-deficient acceptor.
• Once formed, the dative bond behaves like any other covalent bond.
• Dot-and-cross diagrams must show outer-shell electrons, charges and the origin of the bonding pair accurately.

Explanation

• Metallic bonding is the strong electrostatic attraction between a lattice of positive metal ions and delocalised electrons.
• The electrons are not attached to a particular ion and can move through the structure, so a metal conducts electricity when solid or liquid.
• Metallic attraction is non-directional, allowing layers of ions to slide while remaining attracted to the delocalised electrons; this explains malleability and ductility.
• Bond strength generally increases with greater ionic charge, more delocalised electrons and smaller ionic radius because attraction between the ions and electron density is stronger.
• Metals do not consist of separate molecules.

Explanation

• The four crystal types are ionic, metallic, macromolecular and molecular. Sodium chloride, magnesium, diamond or graphite, and iodine or ice are respective examples.
• Ionic and metallic solids contain giant lattices; macromolecular substances have many covalent bonds; molecular solids contain discrete molecules held by intermolecular forces. Melting a molecular solid overcomes intermolecular attractions, not covalent bonds inside molecules.
• Ionic substances conduct only when molten or dissolved because ions can move. Metals and graphite conduct through delocalised electrons.
• Structure is inferred from combined melting-point, solubility and conductivity evidence.
• A requested diagram must show the specified number and type of particles.

Explanation

• Electron charge clouds repel and arrange as far apart as possible around a central atom. Count each bonding region, including a multiple bond as one region, and each lone pair.
• Repulsion decreases in the order lone pair–lone pair $>$ lone pair–bond pair $>$ bond pair–bond pair, so lone pairs compress adjacent bond angles.
• Base geometries are linear $180^{\circ}$, trigonal planar $120^{\circ}$, tetrahedral $109.5^{\circ}$, trigonal bipyramidal $90^{\circ}$ and $120^{\circ}$, and octahedral $90^{\circ}$.
• In a trigonal bipyramid, lone pairs occupy equatorial positions first.
• Shape names describe atom positions, not lone pairs.

Explanation

• Electronegativity is the power of an atom to attract the bonding pair of electrons in a covalent bond.
• A difference in electronegativity produces an unequal electron distribution: the more electronegative atom is $\delta^-$ and the other is $\delta^+$.
• A bond dipole has both magnitude and direction.
• To decide whether a whole molecule has a permanent dipole, identify every polar bond and use the molecular shape to combine their dipoles.
• Equal bond dipoles arranged symmetrically cancel, so a molecule can contain polar bonds but have no overall permanent dipole.

Explanation

• Induced dipole–dipole forces act between all atoms and molecules. Their strength generally increases with electron number and with molecular surface contact.
• Permanent dipole–dipole attractions also act between polar molecules.
• Hydrogen bonding requires hydrogen covalently bonded to nitrogen, oxygen or fluorine and a lone pair on nitrogen, oxygen or fluorine in another molecule.
• Boiling-point comparisons must identify the forces present and compare their overall strength, not simply molecular mass.
• In ice, hydrogen bonds hold water molecules in an open arrangement, making ice less dense than liquid water; the bonds are not described as inherently stronger.

Explanation

• Enthalpy change, $\Delta H$, is the heat-energy change at constant pressure.
• An exothermic process transfers heat to the surroundings and has $\Delta H<0$; an endothermic process absorbs heat and has $\Delta H>0$.
• A standard enthalpy change uses substances in their standard states at $100\,\mathrm{kPa}$ and a stated temperature.
• Standard enthalpy of combustion is the enthalpy change when one mole of a substance is completely burned in oxygen under standard conditions.
• Standard enthalpy of formation is the change when one mole of a compound forms from its constituent elements in their standard states.

Explanation

• Calorimetry estimates heat transfer using $q=mc\Delta T$, where $m$ is the mass undergoing the temperature change, $c$ is its specific heat capacity and $\Delta T$ is the measured temperature change.
• In solution calorimetry, total solution volume and density commonly estimate mass, while the limiting reagent supplies the reacting amount.
• Convert heat to a molar enthalpy using $\Delta H=-q/n$ when a temperature rise shows that the reaction released heat.
• Required practical 2 measures an enthalpy change, for example by dissolution, neutralisation, displacement or alcohol combustion.
• Heat loss, incomplete combustion, evaporation and heating the apparatus create systematic differences that must be evaluated.