LC Chemistry common mistakes
LC Chemistry · Understand the traps and how to avoid them.
Writing electron configurations for transition-metal ions by removing electrons from the 4s subshell last (e.g. giving Fe2+ as [Ar]3d4 4s2 instead of [Ar]3d6), and filling 3d before 4s for neutral atoms.
The aufbau/energy-level diagram is taught as a fixed ladder, so students treat 4s as always lower than 3d. The 4s/3d crossover reverses once the 3d shell is occupied and again on ionisation, and this nuance is often glossed over.
What to do instead
Learn the rule explicitly: for neutral atoms fill 4s before 3d, but for cations remove 4s electrons first. Practise ten ion configurations (Fe2+, Fe3+, Cu+, Cr3+, Mn2+) and check each against the periodic-table position rather than the ladder diagram.
Treating a polar bond as automatically making a polar molecule, e.g. calling CO2 or CCl4 polar because the C=O or C-Cl bonds are polar.
Students memorise electronegativity differences but do not apply VSEPR geometry to add the bond dipoles as vectors, so they never test for cancellation by symmetry.
What to do instead
For every molecule, draw the VSEPR shape first, then add bond-dipole arrows and check whether they cancel. Drill the standard set: CO2, CCl4, BF3 (non-polar by symmetry) versus H2O, NH3, CHCl3 (polar). State the shape and the cancellation in one sentence each time.
Balancing an equation but then reading mole ratios from the unbalanced formula or from the coefficients of the wrong species, and failing to identify the limiting reagent before calculating product mass.
Under time pressure candidates jump from a mass straight to a product mass using a single conversion, skipping the balanced-equation step and the limiting-reagent comparison.
What to do instead
Adopt a fixed four-step template: balance, convert grams to moles, divide each reactant's moles by its coefficient to find the limiting reagent, then convert the limiting reagent's moles to product. Always write the limiting-reagent division explicitly, even when it looks obvious.
Calculating titration results with volumes in cm3 substituted directly into c = n/V, or mis-matching the acid:base mole ratio (e.g. using 1:1 for H2SO4 + NaOH).
The formula c = n/V assumes V in dm3; students mix cm3 and dm3, and they forget that the stoichiometric ratio comes from the balanced equation, not from the formula as written.
What to do instead
Convert every volume to dm3 (divide by 1000) as the first written step, and write the balanced neutralisation equation above every calculation to extract the ratio. For polyprotic acids, confirm the ratio with the equation (H2SO4 + 2NaOH gives 1:2) before substituting.
Using the ideal gas law with the wrong unit set, e.g. plugging pressure in kPa with R = 8.314 J mol-1 K-1 and volume in cm3, or forgetting to convert temperature to kelvin.
Multiple versions of R circulate and students pick one by memory without checking that the units of P, V and T match it; Celsius is used because it is the number on the thermometer.
What to do instead
Standardise on one unit set: P in Pa, V in m3, T in K, R = 8.314. Write the conversions (kPa to Pa times 1000, cm3 to m3 divide by 10^6, Celsius plus 273) as visible lines before substitution, and sanity-check the answer against an estimate.
Confusing the signs and definitions in thermochemistry: reporting an exothermic reaction with a positive delta-H, or mixing up bond-breaking (endothermic) and bond-forming (exothermic) when using average bond enthalpies.
The convention that delta-H is negative for exothermic reactions conflicts with everyday intuition that 'releasing energy' should be positive, and bond-enthalpy cycles are often memorised as a formula without meaning.
What to do instead
Anchor on the system: delta-H is negative when the system loses energy to the surroundings. For bond enthalpies, write 'energy in to break bonds minus energy out from forming bonds' in words before numbers, and check that a combustion always comes out negative.
Predicting the direction of an equilibrium shift by counting moles on each side only, ignoring the effect of a change in pressure on a system with equal moles of gas, or assuming a catalyst shifts the position of equilibrium.
Students apply 'increase pressure shifts to the side with fewer moles' mechanically, without checking whether the mole counts differ, and they conflate faster attainment with a changed position.
What to do instead
Before predicting, write the balanced equation and compare total gas moles on each side. If equal, pressure has no effect on position. State explicitly that a catalyst changes only the rate of approach, not the equilibrium position, and apply Le Chatelier's principle species by species.
Naming or drawing organic structures with the wrong suffix or locant, e.g. naming a carboxylic acid as an aldehyde, or numbering an alkyl chain from the end that gives the substituent the higher number.
Functional-group priority and IUPAC numbering rules are learned as lists rather than as a decision procedure, so candidates apply whichever rule they recall first.
What to do instead
Use a fixed naming procedure: identify the principal functional group (highest priority), choose the longest chain containing it, number to give that group the lowest locant, then name substituents alphabetically. Practise by naming and drawing the same ten molecules in both directions.
Balancing redox half-equations without adding water and H+ (or OH-) to balance oxygen and hydrogen, and mismatching the number of electrons transferred between the two half-equations.
The electron-balancing step is treated as optional, and students forget that oxygen and hydrogen must be balanced with H2O and H+ in acidic solution or OH- in basic solution.
What to do instead
Follow the half-equation checklist in order: balance atoms other than O and H, balance O with H2O, balance H with H+, balance charge with electrons, then multiply each half-equation so the electrons cancel before adding. Verify the final charge balance on both sides.
In practical questions, describing a test or procedure without the specific observable (e.g. 'test for chloride' with no mention of the white precipitate that dissolves in ammonia), or failing to state a control and a safety precaution.
Candidates recall the reagent but not the expected observation, and they treat practical questions as recall rather than as a report of what would be seen and done.
What to do instead
Learn each qualitative test as a triple: reagent, observation, and confirmatory step (AgNO3 gives a white precipitate soluble in dilute ammonia for Cl-, insoluble for Br- and I-). In every practical answer, name the observation, a control or blank, and one relevant safety measure.