STPM Chemistry 962 common mistakes
STPM Chemistry 962 · Understand the traps and how to avoid them.
When asked for the electronic configuration of a transition-metal atom or ion, the candidate fills the 4s orbital after 3d (e.g. writing Fe as [Ar] 3d4 4s2 or Cu+ as [Ar] 3d10) and then cannot explain why the 4s electrons are lost first on ionisation.
Candidates memorise the aufbau filling order 4s-before-3d and apply it mechanically to ions as well as atoms, without realising that the relative energies of 3d and 4s change once the nuclear charge is increased and the 3d subshell contracts. They also do not check their configuration against the observed anomalous cases (Cr, Cu) and their ions.
What to do instead
Write configurations for atoms using the filling order but always verify with the two anomalies Cr ([Ar] 3d5 4s1) and Cu ([Ar] 3d10 4s1). For ions, remove electrons from the highest-n shell first, i.e. 4s before 3d, and then count d electrons to get the correct ion configuration (Fe2+ = [Ar] 3d6, not [Ar] 3d4 4s2). Practise a table of M, M2+ and M3+ for the first-row transition metals and check each against the number of unpaired electrons inferred from magnetic moment data.
When drawing a dot-and-cross diagram or predicting a shape, the candidate counts lone pairs on the central atom incorrectly (e.g. treating the double bond in CO2 as two separate electron domains, or forgetting the lone pair in NH3 and H2O) and then gives the wrong VSEPR geometry and bond angle.
The candidate applies VSEPR by counting atoms rather than electron domains, and does not systematically compute the number of valence electrons and hence the lone pairs. The difference between electron-pair geometry and molecular geometry is not distinguished, so a tetrahedral electron arrangement is reported as the molecular shape for water.
What to do instead
Adopt a fixed procedure: total valence electrons, subtract those used in sigma bonds, place the remainder on the central atom as lone pairs, then count domains (single, double and triple bonds each count as one domain). Use the domain count to get the electron-pair geometry and then subtract lone pairs to name the molecular shape. Memorise the standard angle trend (CH4 109.5, NH3 107, H2O 104.5) and explain it by lone-pair repulsion, and practise CO2 (linear, two domains), SO2 (bent, one lone pair) and SF6 (octahedral) as contrasting cases.
In a titration or gravimetric calculation, the candidate uses the given mass or volume directly in the mole ratio without converting to moles, or mixes up the stoichiometric coefficients when setting up the ratio (e.g. using 1:1 for a reaction that is 1:2).
The candidate jumps straight to a proportion between the two substances, treating the balanced equation as a direct mass or volume relationship. They also fail to identify the limiting reagent in problems where two reactant quantities are given, and do not convert cm3 to dm3 before using concentration.
What to do instead
Use a rigid four-step scaffold for every stoichiometry problem: (1) write and balance the equation; (2) convert every given quantity to moles (n = m/Mr, n = cV/1000, n = V/24 at r.t.p.); (3) apply the mole ratio from the balanced equation; (4) convert back to the requested quantity. When two reactant amounts are given, compute moles of each and divide by the stoichiometric coefficient to find the limiting reagent before proceeding. Always show units at each step and check the answer's order of magnitude.
When applying Le Chatelier's principle to a gaseous equilibrium, the candidate states that increasing total pressure always shifts the equilibrium to the side with fewer moles, even when the number of gaseous moles is equal on both sides, or claims that a catalyst increases the equilibrium yield.
The candidate learns the pressure rule as a slogan without checking whether Δn(gas) is zero, and confuses the effect of a catalyst on rate with an effect on position of equilibrium. They also fail to distinguish the effect of changing pressure (via volume) from adding an inert gas at constant volume.
What to do instead
Before applying the pressure rule, count gaseous moles on each side and state Δn explicitly; if Δn = 0, pressure has no effect on the position of equilibrium. State clearly that a catalyst changes only the rate of attainment, not the equilibrium constant or yield. For inert-gas questions, decide whether volume or pressure is held constant and reason from partial pressures. Practise writing a one-line justification with each prediction, naming the shift direction and the reason.
When calculating the pH of a weak acid or a buffer, the candidate uses [H+] = c directly (the strong-acid formula) or substitutes the analytical concentration of the acid into Ka without allowing for dissociation, and for buffers uses concentrations rather than the correct ratio in the Henderson-Hasselbalch expression.
The candidate has one formula for strong acids and applies it to all acids, and does not appreciate that Ka is small so the degree of dissociation is small. In buffer problems they also forget that the conjugate base concentration comes from the salt, not from the weak acid, and they neglect the effect of dilution on the ratio.
What to do instead
Always classify the acid first (strong, weak, buffer, salt of weak acid/base) and select the corresponding expression. For a weak acid use Ka = x^2/(c - x) and state the approximation c - x ≈ c only after checking x is under about 5% of c. For buffers use pH = pKa + log([salt]/[acid]) and confirm which species is the base. Check plausibility: a 0.1 mol dm-3 weak acid with Ka around 10^-5 should give pH near 3, not 1.
In an electrochemistry question the candidate reverses the sign convention or the direction of electron flow, writes the cell diagram with the anode on the right, or reports E°cell as a negative value for a spontaneous reaction.
The candidate memorises 'oxidation at the anode' but does not connect it to the IUPAC convention that the anode is the negative electrode in a galvanic cell, and confuses galvanic and electrolytic cells. Sign conventions for electrode potentials and for cell emf are mixed up under exam pressure.
What to do instead
Fix the conventions in writing: in a galvanic cell oxidation occurs at the anode, which is negative; reduction at the cathode, which is positive; electrons flow from anode to cathode through the external circuit. Write cell diagrams as anode | anode solution || cathode solution | cathode. Compute E°cell = E°cathode - E°anode and check that a spontaneous cell gives a positive value; if negative, the reaction as written is non-spontaneous. For electrolytic cells, state explicitly that the polarity is imposed by the external supply.
When drawing organic reaction mechanisms, the candidate uses a curly arrow from an atom rather than from a bond or lone pair, draws arrows that do not conserve charge, or omits the intermediate carbocation and its subsequent rearrangement or attack.
The candidate learns mechanisms as pictures to reproduce rather than as electron bookkeeping, so arrows are placed by imitation. They also do not check formal charges or the octet at each centre, and forget that a mechanism must show every bond made and broken.
What to do instead
For every arrow, start it at a lone pair or a bond (never at an atom symbol) and end it where the electrons go, then check that charge is conserved and no atom exceeds its octet. After drawing each step, label the intermediate (carbocation, oxonium ion, etc.) and ask what happens next. Practise the same mechanism (e.g. electrophilic addition to an alkene, nucleophilic substitution) on three different substrates and compare the intermediates explicitly.
In reaction-kinetics questions the candidate reads a rate-concentration graph or a table of initial rates and concludes the order from the wrong experiment, or confuses rate with rate constant, and states that doubling temperature doubles the rate.
The candidate does not hold other concentrations constant when comparing experiments, or compares experiments where two concentrations change simultaneously. The distinction between the rate equation, the rate constant and the temperature dependence via the Arrhenius equation is not secure, so temperature effects are treated as a simple proportionality.
What to do instead
When determining order from initial-rate data, select pairs of experiments in which only one concentration changes, and compute the ratio of rates to deduce the order for that reactant. Write the rate equation explicitly with units for k. For temperature, use the Arrhenius equation qualitatively: a modest temperature rise increases the fraction of molecules with energy above Ea, so the rate increases by a factor that depends on Ea, not by a fixed multiple. Confirm order by checking that the units of k match the overall order.
When interpreting an infrared or mass spectrum, the candidate assigns a peak without checking the characteristic wavenumber range, or reads the molecular ion as the base peak and so reports the wrong molecular mass.
The candidate memorises a few numbers without the ranges and without the logic of fragmentation. In mass spectra they do not identify the molecular ion as the highest m/z peak that corresponds to the whole molecule (often a small peak), and they confuse the base peak (most abundant fragment) with M+.
What to do instead
Learn infrared absorptions as ranges with the associated bond: O-H 3200-3600 (broad), C=O 1680-1750 (strong), C-H 2850-3100, N-H 3300-3500. For mass spectra, locate M+ as the highest m/z peak consistent with the molecule and check the nitrogen rule and isotope patterns (M+2 for Cl or Br). Then assign the base peak to a plausible fragment and confirm that the fragments sum to the molecular ion. Practise with a table of common fragments (15, 29, 43, 77).
In thermochemistry calculations the candidate adds or subtracts enthalpy values with the wrong sign, forgets that bond-breaking is endothermic and bond-making exothermic, or applies Hess's law without reversing the equations whose direction is changed.
The candidate treats ΔH as an arithmetic quantity to be combined without tracking the direction of each reaction. Sign conventions for formation, combustion and bond enthalpy are not distinguished, and the candidate does not check whether the final answer is plausible for the process described.
What to do instead
Before combining, write each equation with its ΔH and an arrow showing direction; when reversing an equation, reverse the sign; when multiplying, multiply ΔH by the same factor. For bond enthalpies use ΔH = Σ(bonds broken) - Σ(bonds formed) and state that bond breaking is positive. For formation and combustion, check that the sign matches the definition (formation from elements, combustion releasing heat). Finish by asking whether the magnitude and sign are physically sensible for the reaction.