Cambridge Pre-U Chemistry (Principal Subject) Syllabus

Write the expression for $K_{c}$ for the reaction $\ce{N2 + 3H2 <=> 2NH3}$.

$$K_{c} = \frac{[\ce{NH3}]^{2}}{[\ce{N2}][\ce{H2}]^{3}}$$

Define $K_{p}$ and write it for $\ce{2SO2 + O2 <=> 2SO3}$.

$K_{p}$ is the equilibrium constant in terms of partial pressures: $$K_{p} = \frac{p_{\ce{SO3}}^{2}}{p_{\ce{SO2}}^{2}\,p_{\ce{O2}}}$$

How does increasing temperature affect $K_{c}$ for an exothermic reaction?

For an exothermic forward reaction, increasing temperature shifts equilibrium backwards and decreases $K_{c}$ (only temperature changes the value of $K$).

Define a Brønsted–Lowry acid and base.

An acid is a proton ($\ce{H+}$) donor; a base is a proton acceptor.

Give the expression for $\text{pH}$ and for $K_{w}$ at $298\,\text{K}$.

$$\text{pH} = -\log_{10}[\ce{H+}]$$ and $K_{w} = [\ce{H+}][\ce{OH-}] = 1.0\times10^{-14}\,\text{mol}^{2}\,\text{dm}^{-6}$.

Write the Henderson–Hasselbalch equation for a buffer.

$$\text{pH} = \text{p}K_{a} + \log_{10}\frac{[\text{A}^{-}]}{[\text{HA}]}$$

Explain how an acidic buffer of ethanoic acid and sodium ethanoate resists pH change on adding acid.

Added $\ce{H+}$ is removed by the conjugate base: $\ce{CH3COO- + H+ -> CH3COOH}$. The large reservoir of $\ce{CH3COO-}$ keeps $[\ce{H+}]$ and hence pH almost constant.

Describe the trend in atomic radius across Period 3 and explain it.

Atomic radius decreases across the period because nuclear charge increases while electrons enter the same shell, increasing effective nuclear attraction with little change in shielding.

Describe and explain the trend in melting point across Period 3 ($\ce{Na}$ to $\ce{Ar}$).

Rises $\ce{Na->Al}$ (stronger metallic bonding), peaks at $\ce{Si}$ (giant covalent lattice), then falls sharply to small molecular/atomic species ($\ce{P4, S8, Cl2, Ar}$) held by weak London forces.

What is observed when $\ce{Na}$ and $\ce{Mg}$ react with water?

$\ce{Na}$ reacts vigorously with cold water giving $\ce{NaOH}$ + $\ce{H2}$; $\ce{Mg}$ reacts only very slowly with cold water but readily with steam to give $\ce{MgO}$ + $\ce{H2}$.

Describe the trend in thermal stability of Group 2 carbonates down the group and explain it.

Thermal stability increases down the group: larger cations have lower charge density, polarising the carbonate ion less, so more heat is needed to decompose it to the oxide and $\ce{CO2}$.

State the trend in solubility of Group 2 hydroxides and sulfates down the group.

Hydroxides become more soluble down the group; sulfates become less soluble down the group (e.g. $\ce{BaSO4}$ is insoluble).

Describe the trend in oxidising power of the Group 17 halogens down the group.

Oxidising power decreases down the group ($\ce{F2 > Cl2 > Br2 > I2}$) because atomic radius increases and the ability to gain an electron weakens.

Write the equation and observation for the displacement of bromide by chlorine.

$$\ce{Cl2 + 2Br- -> 2Cl- + Br2}$$ The solution turns orange (or organic layer orange/brown) as bromine is displaced.

Describe the disproportionation of chlorine in cold dilute $\ce{NaOH}$.

$$\ce{Cl2 + 2NaOH -> NaCl + NaClO + H2O}$$ Chlorine is both oxidised (to $\ce{ClO-}$, $+1$) and reduced (to $\ce{Cl-}$, $-1$).

Define a transition element.

A $d$-block element that forms at least one stable ion with a partially filled $d$ subshell.

List four characteristic properties of transition elements.

Variable oxidation states; formation of coloured ions/complexes; catalytic activity; ability to form complex ions with ligands.

What is a ligand, and what is the coordination number of $\ce{[Cu(H2O)6]^{2+}}$?

A ligand is a species with a lone pair that forms a dative bond to a central metal ion. The coordination number of $\ce{[Cu(H2O)6]^{2+}}$ is $6$ (octahedral).

Explain why transition metal complexes are coloured.

Ligands split the $d$ orbitals into two energy levels of gap $\Delta E$; $d$–$d$ electron transitions absorb visible light of frequency $\Delta E = hf$, and the complementary colour is transmitted/observed.

Write the half-equation and colour change when acidified $\ce{MnO4-}$ acts as an oxidant.

$$\ce{MnO4- + 8H+ + 5e- -> Mn^{2+} + 4H2O}$$ Colour changes from purple ($\ce{MnO4-}$) to very pale pink/colourless ($\ce{Mn^{2+}}$).

Define oxidation and reduction in terms of electron transfer and oxidation number.

Oxidation is loss of electrons / an increase in oxidation number; reduction is gain of electrons / a decrease in oxidation number (OIL RIG).

Define standard electrode potential, $E^{\ominus}$.

The EMF of a half-cell relative to the standard hydrogen electrode under standard conditions ($298\,\text{K}$, $1\,\text{mol dm}^{-3}$ solutions, $100\,\text{kPa}$ gases), with the SHE defined as $0\,\text{V}$.

How is the standard cell EMF calculated from two electrode potentials?

$$E^{\ominus}_{\text{cell}} = E^{\ominus}_{\text{cathode (reduction)}} - E^{\ominus}_{\text{anode (oxidation)}}$$ A positive $E^{\ominus}_{\text{cell}}$ indicates a feasible reaction.

Define the term 'homologous series' and give the general formula of the alkanes.

A family of organic compounds with the same functional group and general formula, differing by $\ce{CH2}$, with similar chemical properties and a gradual trend in physical properties. Alkanes: $C_{n}H_{2n+2}$.

Describe the mechanism type and propagation steps for the free-radical chlorination of methane.

Free-radical substitution (initiated by UV). Propagation: $\ce{Cl. + CH4 -> .CH3 + HCl}$ then $\ce{.CH3 + Cl2 -> CH3Cl + Cl.}$.

Define electrophile and state the mechanism for $\ce{Br2}$ adding to ethene.

An electrophile is an electron-pair acceptor (electron-deficient species). Ethene + $\ce{Br2}$ undergoes electrophilic addition to give 1,2-dibromoethane: $\ce{CH2=CH2 + Br2 -> CH2BrCH2Br}$.

State Markovnikov's rule for addition of $\ce{HBr}$ to an unsymmetrical alkene.

The hydrogen adds to the carbon already bearing the most hydrogens, so the more stable (more substituted) carbocation forms, giving the major product (e.g. 2-bromopropane from propene).

Give reagents/conditions for oxidising a primary alcohol to an aldehyde and then to a carboxylic acid.

Acidified $\ce{K2Cr2O7}$ (or $\ce{KMnO4}$). Aldehyde: distil off as formed (limited oxidant). Carboxylic acid: heat under reflux with excess oxidant. Orange dichromate turns green ($\ce{Cr^{3+}}$).

How do Tollens' and Fehling's reagents distinguish aldehydes from ketones?

Aldehydes are oxidised and give positive tests (Tollens': silver mirror; Fehling's: brick-red $\ce{Cu2O}$ precipitate); ketones give no reaction with either.

Describe the electrophilic substitution mechanism for nitration of benzene, including the electrophile.

Benzene reacts with the nitronium ion $\ce{NO2+}$ (from $\ce{HNO3 + H2SO4}$) by electrophilic substitution: $\ce{NO2+}$ adds to form an unstable intermediate, then $\ce{H+}$ is lost to restore aromaticity, giving nitrobenzene.

Distinguish between structural isomerism and stereoisomerism, naming the two types of stereoisomerism.

Structural isomers have the same molecular formula but different atom connectivity. Stereoisomers have the same connectivity but different spatial arrangement: (i) geometric (cis–trans / $E$–$Z$) and (ii) optical (enantiomers).

What structural feature gives rise to optical isomerism, and what is a key physical property of enantiomers?

A chiral centre — a carbon bonded to four different groups. Enantiomers are non-superimposable mirror images that rotate the plane of plane-polarised light in equal but opposite directions.

Distinguish addition and condensation polymerisation, giving the repeat-unit feature of each.

Addition: monomers with C=C join with no loss of atoms (e.g. poly(ethene)). Condensation: monomers with two functional groups join with loss of a small molecule (e.g. $\ce{H2O}$ or $\ce{HCl}$), forming polyesters/polyamides.

Identify the bond responsible for an absorption near $1700\,\text{cm}^{-1}$ in an infrared spectrum.

The carbonyl $\ce{C=O}$ stretch (found in aldehydes, ketones, carboxylic acids and esters), typically $1680$–$1750\,\text{cm}^{-1}$.

In $^{1}\text{H}$ NMR, what do (a) the number of peaks, (b) the integration, and (c) the splitting tell you?

(a) Number of peaks = number of chemically distinct hydrogen environments; (b) integration (area) = relative number of H in each environment; (c) splitting follows the $n+1$ rule, indicating the number of H on adjacent carbons.

What information does the molecular ion peak (M$^{+}$) provide in mass spectrometry?

The mass/charge ratio of the molecular ion gives the relative molecular mass ($M_{r}$) of the compound; the highest $m/z$ peak (ignoring isotope peaks) corresponds to $M^{+}$.

Give the reagents and positive observations for testing a halide ion in solution.

Add dilute $\ce{HNO3}$ then $\ce{AgNO3(aq)}$: $\ce{Cl-}$ white ppt (soluble in dilute $\ce{NH3}$), $\ce{Br-}$ cream ppt (soluble in conc. $\ce{NH3}$), $\ce{I-}$ yellow ppt (insoluble in $\ce{NH3}$).

Describe the test for an ammonium ion ($\ce{NH4+}$).

Warm with $\ce{NaOH(aq)}$: ammonia gas is released which turns damp red litmus paper blue: $\ce{NH4+ + OH- -> NH3 + H2O}$.

State the formula for percentage yield and percentage atom economy.

$$\%\,\text{yield} = \frac{\text{actual moles of product}}{\text{theoretical moles of product}}\times100$$ $$\%\,\text{atom economy} = \frac{M_{r}\,\text{of desired product}}{\text{total }M_{r}\,\text{of all products}}\times100$$

How is the end point in an acid–base titration found, and how is the concentration of an unknown calculated?

An indicator (or pH meter) marks the end point at the equivalence volume. Using $n = c\times V$ and the balanced mole ratio: $$c_{1} = \frac{c_{2}V_{2}}{V_{1}}\times\frac{\text{ratio}_{1}}{\text{ratio}_{2}}$$

Distinguish between random and systematic errors in an experiment, with an example of each.

Random errors scatter readings unpredictably and are reduced by repeats/averaging (e.g. judging a colour change). Systematic errors shift all readings the same way and are not reduced by repeats (e.g. an uncalibrated balance giving a fixed offset).