PAU Chemistry (Cantabria) common mistakes
PAU Chemistry (Cantabria) · Understand the traps and how to avoid them.
Filling orbitals by simply counting electrons into 1s, 2s, 2p, 3s, 3p, 4s in strict aufbau order and writing configurations such as 3d after 4s without testing the (n+l) rule or the anomalous cases (Cr, Cu, and their ions), and forgetting that 4s electrons are removed before 3d when forming cations.
Candidates memorize a single diagonal diagram and never internalize that the 3d subshell fills after 4s but lies lower in energy once occupied, so they cannot justify why Fe2+ is 3d6 and not 3d4 4s2, or why Cr is 3d5 4s1.
What to do instead
For every configuration, write the (n+l) ordering explicitly, then mark exceptions: half-filled d5 and full d10 stability for Cr, Cu, Mo, Ag. Practise ten cation configurations (Fe2+, Fe3+, Cu+, Zn2+, Cr3+) and state the rule 'electrons leave the highest n shell first' before writing the answer.
Ranking periodic properties (atomic radius, ionization energy, electronegativity) by moving down a group and across a period using only one trend at a time, producing answers such as 'Na has higher ionization energy than Mg' or forgetting that ionic radius comparisons require isoelectronic species.
The two trends (across a period: radius decreases, ionization energy increases; down a group: the reverse) are learned as slogans, so when a question mixes a group element with a period element the candidate applies the wrong axis or ignores effective nuclear charge and shielding.
What to do instead
For any comparison, first locate both elements on the periodic table, then state which factor dominates (Zeff or number of shells). For ions, group them by isoelectronic series (e.g., N3-, O2-, F-, Na+, Mg2+) and rank by nuclear charge. Write a one-line justification with each answer, as the exam requires reasoning, not just the symbol.
Drawing Lewis structures and predicting molecular geometry without counting lone pairs on the central atom, so that molecules such as H2O, NH3, SO2, and SF4 are assigned linear or trigonal-planar shapes, and then polarities and intermolecular forces are deduced from the wrong geometry.
Candidates count bonding pairs only, or they confuse the electron-pair geometry with the molecular geometry, and they do not systematically apply VSEPR repulsion ordering (lone pair–lone pair > lone pair–bond pair > bond pair–bond pair).
What to do instead
Use a fixed procedure: total valence electrons, central atom, bonding pairs, then remaining electrons as lone pairs, then count electron domains and apply the VSEPR table (2 linear, 3 trigonal planar, 4 tetrahedral with 0/1/2 lone pairs giving tetrahedral/bent/linear). Only after geometry is fixed, decide the dipole sum and hence the dominant intermolecular force.
Treating hydrogen bonds and dipole–dipole forces as the explanation for every boiling-point anomaly, and predicting that a larger molecule always has a higher boiling point, ignoring that H-bonds require H bonded to N, O, or F, and confusing boiling point (intermolecular forces) with volatility or with bond strength.
The phrase 'hydrogen bonding' is over-applied to any molecule containing H, and candidates do not separate intra- and intermolecular forces, so they explain HCl's low boiling point by H-bonds or claim CH4 boils above H2O because it is heavier.
What to do instead
Before answering, list the forces present with a checklist: London dispersion always, dipole–dipole if polar, H-bond only if H is directly bonded to N/O/F. Compare molecules by identifying the strongest force present, then by molar mass only within the same force class. Practise the H2O/H2S/H2Se and HF/HCl/HBr series and justify each ordering in one sentence.
Applying Hess's law by adding enthalpies without first writing and balancing each reaction, or reversing a reaction without changing the sign of ΔH, and multiplying ΔH by the coefficient when the equation is doubled.
Candidates manipulate numbers before fixing the target equation, so they lose track of which species must cancel, and they treat ΔH as a per-molecule quantity rather than an extensive quantity tied to the stoichiometric coefficients.
What to do instead
Write the target equation first, then label each given equation (1), (2), (3) and record the multiplier and direction needed. Apply the same operation to ΔH: reverse changes sign, multiply by n multiplies ΔH by n. Check that species not in the target cancel, and state units (kJ or kJ/mol) consistently.
Judging spontaneity from ΔH alone, or from ΔS alone, and mis-signing the entropy change for phase changes and for reactions that change the number of gas moles, e.g., declaring a reaction with ΔH > 0 always non-spontaneous.
The Gibbs relation ΔG = ΔH − TΔS is memorized but not used, and candidates forget that ΔS is positive when solid → liquid → gas or when the number of gas molecules increases, and that ΔG depends on temperature.
What to do instead
For each reaction, compute ΔS sign from the change in moles of gas and phase, then build the four-case table (ΔH<0/ΔS>0 always spontaneous; ΔH>0/ΔS<0 never; mixed cases spontaneous above or below T = ΔH/ΔS). Always state the temperature condition, and calculate ΔG° = ΔH° − TΔS° numerically when data are given.
Predicting the effect of a change on equilibrium by intuition rather than by Le Chatelier applied to the balanced equation, e.g., claiming that adding a catalyst shifts the equilibrium, that increasing pressure always shifts the equilibrium, or that adding an inert gas at constant volume shifts it.
The rule is recalled as a slogan without the condition that pressure changes matter only when the number of gas moles differs between sides, and the catalyst's role (kinetics only, no effect on K) is confused with a thermodynamic shift.
What to do instead
For each perturbation, write the balanced equation and count gas moles on each side. Then answer: concentration change shifts toward the side that consumes the added species; pressure increase shifts toward fewer gas moles (no shift if equal); temperature change shifts in the endothermic or exothermic direction and changes K; catalyst changes rate only. State K's dependence on T explicitly.
Calculating pH of a weak acid or base by using [H+] = c directly instead of solving the Ka equilibrium with an ICE table, or ignoring the hydrolysis of a salt and treating every salt solution as neutral.
Candidates apply the strong-acid formula to weak acids, forget to take the square root of Ka·c, and do not connect the conjugate acid/base of a weak species to the salt's pH (e.g., NH4Cl acidic, CH3COONa basic, NaCl neutral).
What to do instead
Use a fixed protocol: write the ionization equation, set up an ICE table, apply Ka = x²/(c − x) with the approximation x << c checked (x/c < 5%), then pH = −log x. For salts, identify the ions, decide which hydrolyses (conjugate of weak acid or weak base), write its hydrolysis equation and compute with the appropriate Ka or Kb = Kw/Ka. In titrations, compute moles before and after addition and use the limiting reagent.
Writing the Ksp expression with the wrong exponents, or forgetting that the stoichiometric coefficients become exponents, and failing to apply the common-ion effect by using the pure-water solubility formula when the solution already contains one of the ions.
The expression is written from memory as [ion][ion] regardless of formula, and the common-ion case is treated as a separate topic rather than as the same equilibrium with a non-zero initial concentration.
What to do instead
Write Ksp from the dissolution equation with each concentration raised to its coefficient (Ag2CrO4: Ksp = [Ag+]²[CrO4²-]). Then always set up an ICE table with the initial common-ion concentration included, and solve for s. Compare with pure water to show the decreased solubility, and check whether precipitation occurs by computing Q and comparing with Ksp.
Balancing redox equations in acidic or basic medium without separating the half-reactions, or forgetting to add H2O and H+ (or OH-) to balance O and H, and then mis-signing the cell potential when reversing an electrode reaction.
Candidates try to balance atoms and charges simultaneously in one line, and they do not apply the rule that E° of a half-reaction changes sign when reversed while E°cell = E°cathode − E°anode uses reduction potentials as tabulated.
What to do instead
Use the half-reaction method: separate oxidation and reduction, balance atoms other than O and H, balance O with H2O, H with H+, charge with electrons, equalize electrons, then add. In basic medium, neutralize H+ with OH- at the end. For cells, identify cathode (reduction) and anode (oxidation), use E°cell = E°cathode − E°anode with tabulated reduction potentials, and never multiply E° by stoichiometric coefficients.
Naming and formulating organic compounds by translating the structural formula word by word, missing the principal functional group priority, choosing the wrong parent chain, or assigning locants that do not give the lowest set, and confusing isomers (positional, functional, geometric) with each other.
The priority order of functional groups and the lowest-locant rule are not applied systematically, so a compound with both an alcohol and a ketone is named as a ketone, or cis/trans is asserted without checking that each carbon of the double bond bears two different groups.
What to do instead
Follow a fixed order: identify the principal group by priority (carboxylic acid > ester > amide > nitrile > aldehyde > ketone > alcohol > amine > alkene > alkyne), choose the longest chain containing it, number to give the principal group the lowest locant and then the lowest locant set to substituents, and assemble alphabetically. For isomerism, explicitly test each double bond carbon for two different substituents before writing cis/trans, and classify functional isomers by their different functional group.