GCE A/L Chemistry common mistakes
GCE A/L Chemistry · Understand the traps and how to avoid them.
When writing electron configurations for transition metal ions, candidates remove electrons from the 4s orbital before the 3d orbital, leading to incorrect configurations such as Fe²⁺ as [Ar]3d⁴4s² instead of [Ar]3d⁶.
This error stems from a misunderstanding of the Aufbau principle and the relative energies of 4s and 3d orbitals. Students learn that 4s fills before 3d, so they assume 4s empties first. They overlook that once the 3d subshell is occupied, the 4s electrons are higher in energy and are removed first when forming cations.
What to do instead
Teach the rule explicitly: for transition metal cations, remove electrons from the highest n orbital first (4s before 3d). Provide practice with several ions (e.g., Fe²⁺, Fe³⁺, Cu⁺, Cr³⁺) and have students write both the atom and ion configurations, emphasizing that the 4s orbital is empty in most transition metal ions. Use orbital box diagrams to visualise the process.
When predicting the shape of a molecule using VSEPR theory, candidates ignore lone pair repulsion or misassign the number of lone pairs, leading to incorrect geometries (e.g., predicting H₂O as linear or NH₃ as trigonal planar).
Students often count only bonding pairs and forget to subtract bonding electrons from the total valence electrons to find lone pairs. They also may not appreciate that lone pairs occupy space and repel more strongly than bonding pairs, altering bond angles and shape.
What to do instead
Provide a step-by-step VSEPR method: (1) draw the Lewis structure, (2) count electron domains (bonding + lone pairs) around the central atom, (3) determine the electron-domain geometry, (4) adjust for lone pairs to get molecular geometry. Use a table of common geometries and have students practise with molecules like H₂O, NH₃, SF₄, and ClF₃, explicitly noting the effect of lone pairs on bond angles.
In calculations involving gases, candidates use the ideal gas equation without converting temperature to kelvin or pressure to pascals, or they mix units (e.g., using °C in pV = nRT or using mL without converting to dm³).
This error arises from rote memorisation of the formula without understanding the units of the gas constant R. Students often overlook that R has units J K⁻¹ mol⁻¹ and that temperature must be in kelvin and volume in cubic metres (or dm³ with appropriate R). They may also be careless with unit conversions under exam pressure.
What to do instead
Insist on a standard data extraction step: before substituting into any gas equation, write down the given values with units and convert them to SI (K, Pa, m³) or to a consistent set (e.g., dm³ and kPa with R = 8.314 J K⁻¹ mol⁻¹). Provide a checklist and practise with problems that require conversions, such as those giving temperature in °C and pressure in atm or mmHg. Emphasise that R = 8.31 J K⁻¹ mol⁻¹ requires volume in m³ and pressure in Pa.
When calculating enthalpy changes using Hess's law or bond enthalpies, candidates add or subtract values in the wrong direction, or they forget to multiply bond enthalpies by the number of bonds broken or formed.
Students often confuse the signs of ΔH for bond breaking (endothermic, positive) and bond forming (exothermic, negative). They may also lose track of stoichiometry when a bond appears multiple times in a structure. The multi-step nature of Hess's law problems can lead to errors in summing equations.
What to do instead
Teach a systematic approach: for bond enthalpies, write ΔH = Σ(bonds broken) – Σ(bonds formed), and explicitly list each bond type with its count and enthalpy. For Hess's law, use a cycle diagram and ensure that equations are reversed and multiplied correctly. Provide practice with complex cycles (e.g., formation of a compound from elements) and encourage students to check signs at each step.
When solving equilibrium problems, candidates assume that the initial concentrations can be used directly in the equilibrium constant expression without considering the change in concentration, or they omit the change row in an ICE table.
This mistake occurs because students memorise Kc = [products]/[reactants] but do not understand that equilibrium concentrations differ from initial ones. They may also struggle with the algebra of solving quadratic equations or making approximations without justification.
What to do instead
Require the use of ICE (Initial, Change, Equilibrium) tables for every equilibrium calculation. Emphasise that the change in concentration is proportional to the stoichiometric coefficients and that the equilibrium row is used in the Kc expression. Practise problems that require solving quadratic equations and teach when the approximation (x small) is valid. Provide worked examples and peer-checking exercises.
In acid-base titrations, candidates calculate pH using the concentration of the acid or base without accounting for dilution or the stoichiometry of the reaction, leading to incorrect pH values, especially at the equivalence point.
Students often apply the formula pH = -log[H⁺] directly to the initial concentration of the acid, ignoring that the acid may be weak or that the volume changes during titration. They may also forget that at the equivalence point, the pH is determined by the hydrolysis of the salt formed.
What to do instead
Teach a step-by-step titration calculation method: (1) calculate moles of acid and base, (2) determine the limiting reagent and moles of excess, (3) calculate total volume, (4) find concentration of excess H⁺ or OH⁻, (5) calculate pH. For weak acids, use Ka and set up an equilibrium. For equivalence points, identify the salt and consider hydrolysis. Provide practice with strong acid-strong base, weak acid-strong base, and weak base-strong acid titrations.
When determining the rate law from experimental data, candidates simply take the ratio of initial rates and concentrations without considering that the order with respect to a reactant must be deduced by keeping other concentrations constant, or they incorrectly assume the reaction is first order because it is elementary.
Students often confuse the stoichiometric coefficients with reaction orders. They may also not realise that the rate law must be determined experimentally and that the method of initial rates requires isolating one reactant at a time. The mathematical manipulation of ratios can be challenging.
What to do instead
Teach the method of initial rates explicitly: choose two experiments where only one reactant's concentration changes, set up the ratio (rate1/rate2) = ([A]1/[A]2)^n, and solve for n. Emphasise that orders are not necessarily related to stoichiometry. Provide datasets and have students determine the order for each reactant and then the overall rate law. Also discuss the effect of temperature on k.
When drawing or interpreting electrochemical cells, candidates reverse the anode and cathode, or they incorrectly assign the direction of electron flow and the sign of the electrodes in a galvanic cell.
Students often memorise 'anode is negative' without understanding that this applies to galvanic cells, while in electrolytic cells the anode is positive. They may also confuse oxidation and reduction definitions (OIL RIG) and apply them incorrectly to the cell notation.
What to do instead
Use the mnemonic 'An Ox' (anode oxidation) and 'Red Cat' (cathode reduction). For galvanic cells, emphasise that oxidation occurs at the anode (negative electrode) and reduction at the cathode (positive electrode). For electrolytic cells, the polarity is reversed. Practise with cell diagrams, writing half-equations, and calculating cell potentials. Have students label electron flow and ion migration in both types of cells.
In organic nomenclature, candidates misidentify the principal functional group or incorrectly number the carbon chain, leading to names like 2-methylpropan-3-ol instead of 2-methylpropan-1-ol, or they fail to use the lowest locant set for multiple substituents.
Students often prioritise the wrong functional group (e.g., treating an alcohol as a substituent when it should be the principal group) or number from the end that gives the first substituent rather than the lowest set of locants for all substituents. They may also forget that the principal functional group determines the suffix and numbering.
What to do instead
Teach a priority list of functional groups (e.g., carboxylic acid > ester > amide > nitrile > aldehyde > ketone > alcohol > amine > alkene > alkyne > alkane > ether > halide). Practise numbering the chain from the end that gives the lowest locant to the principal functional group, and then apply the lowest locant set rule for substituents. Use examples with multiple functional groups and have students explain their numbering choices.
In organic reaction mechanisms, candidates draw curly arrows starting from the wrong atom or in the wrong direction, such as showing a nucleophile attacking a carbon without indicating the movement of electrons from the nucleophile to the electrophilic carbon, or they omit the formation of intermediates like carbocations.
Students often lack a clear understanding of electron-rich and electron-poor species and the flow of electron pairs. They may also not realise that curly arrows represent electron pair movement, not atom movement, and that the arrow should start from the source of electrons (lone pair or bond) and point to the electron-deficient atom.
What to do instead
Reinforce the meaning of curly arrows: they show the movement of electron pairs from a nucleophile (or π bond) to an electrophile. Provide practice with common mechanisms (e.g., nucleophilic substitution, electrophilic addition) and have students annotate each step with the electron flow. Use coloured arrows to distinguish between bond formation and bond breaking, and emphasise the role of intermediates. Peer review of drawn mechanisms can help identify errors.