OSAT Chemistry common mistakes
OSAT Chemistry · Understand the traps and how to avoid them.
Balancing an equation or answering a stoichiometry item with moles of the reactant as the answer, never converting back to grams, liters of gas at STP, or molarity of the product.
The mole is the natural unit of the calculation, so candidates stop at the mole result, especially when the item asks for mass and the arithmetic already feels finished.
What to do instead
Before computing, write the target unit at the end of the setup line (mol → g, mol → L at STP, mol → M). After obtaining moles, perform the final conversion and check that units cancel to the requested unit. Practise with mixed-unit items where the requested unit changes each time.
Treating the limiting reactant as the one with the smaller starting number of moles (or smaller mass) rather than dividing moles by the stoichiometric coefficient.
Early instruction emphasizes 'smallest amount runs out', and candidates carry that shortcut into items where coefficients differ from 1:1.
What to do instead
Always divide the moles of each reactant by its coefficient in the balanced equation; the smallest quotient is limiting. Re-solve two items where the smaller-mole reactant is actually in excess and note the reversal.
Predicting electron geometry instead of molecular geometry for a molecule with lone pairs, so H2O is called linear or 180 degrees and NH3 is called trigonal planar.
VSEPR is learned as a counting exercise, and candidates report the arrangement of electron domains without subtracting the lone-pair influence on bond angles.
What to do instead
Use a two-step routine: count domains for electron geometry, then count lone pairs on the central atom to derive molecular shape and approximate angle. Build a table for AX2, AX3, AX2E, AX3E, AX4, AX2E2, AX4E, AX5, AX6 and recite each shape and angle before practising items.
Writing a net ionic equation that keeps spectator ions, or fails to split a soluble strong electrolyte while incorrectly splitting a weak acid or a precipitate.
Solubility rules and the strong/weak electrolyte distinction are memorized as lists but not applied to the specific species in the item.
What to do instead
Run three checks: split only soluble strong electrolytes (strong acids, strong bases, soluble salts); keep weak acids, weak bases, and precipitates intact; cancel identical ions from both sides. Practise by classifying each species as strong, weak, or nonelectrolyte before writing the equation.
Assigning oxidation numbers that ignore the sum rule, e.g., giving oxygen −2 in peroxides or failing to make the sum equal the ion charge in a polyatomic ion such as MnO4−.
Candidates memorize individual rules without the overriding constraint that the sum of oxidation numbers equals the overall charge.
What to do instead
Set up an algebraic equation: sum of oxidation numbers = charge (0 for neutral species). Flag exceptions (peroxide O −1, hydride H −1, F always −1) and solve for the unknown element. Work through one ion and one compound per practice set.
Confusing the rate of disappearance of a reactant with the rate of the reaction, or omitting the 1/coefficient factor when relating rates of different species.
Rate is introduced as change in concentration over time, and candidates apply that definition to whichever species appears first in the equation.
What to do instead
Use rate = −(1/a)(Δ[A]/Δt) = (1/c)(Δ[C]/Δt) for aA → cC, and check that all expressions give the same value. Practise items where the coefficients are not all 1.
Treating a weak acid like a strong acid in pH calculations, i.e., taking pH = −log(initial acid concentration) without using the Ka expression or an ICE table.
The strong-acid formula is simple and overlearned; candidates apply it automatically whenever a pH is requested, ignoring the Ka or the percent-ionization data given.
What to do instead
Before computing pH, classify the acid as strong or weak using the given Ka or the acid identity. For weak acids, set up an ICE table with x = [H3O+], solve Ka = x²/(C − x), and only drop x if x < 5% of C. Check the answer with percent ionization.
Assuming a salt solution is neutral, or predicting the wrong direction of pH shift for a buffer after adding strong acid or base, without considering conjugate acid/base strength.
Neutrality is the default expectation for salts, and buffer behavior is memorized as 'resists change' without tracking which component is consumed.
What to do instead
For salts, identify the parent acid and base: strong/strong gives neutral, strong acid/weak base gives acidic, weak acid/strong base gives basic. For buffers, write the neutralization reaction first, subtract from the buffer components, then use Henderson–Hasselbalch with the new amounts.
Confusing the sign of ΔG, ΔH, and ΔS, or calling a reaction spontaneous because it is exothermic, when the entropy term determines the outcome at the given temperature.
Exothermic reactions feel 'favorable', and candidates collapse the Gibbs equation into a single-rule heuristic instead of evaluating ΔG = ΔH − TΔS at the stated temperature.
What to do instead
Build the four-case table (ΔH and ΔS signs) and state spontaneity for low and high T. For each item, compute or estimate ΔG at the given temperature and check that the sign matches the conclusion. Never infer spontaneity from ΔH alone.
Drawing a Lewis structure that satisfies octets but omits formal-charge minimization or ignores resonance, so the reported structure is not the preferred one and the bond order or polarity answer is wrong.
Candidates stop when every atom has eight electrons, treating octet satisfaction as the sole criterion.
What to do instead
After drawing, compute formal charge on each atom (valence − nonbonding − ½ bonding) and shift electrons to minimize charges, placing negative charge on the more electronegative atom. Then identify equivalent resonance structures and average bond order. Practise on ions such as NO3−, SO4²−, and CO3²−.