AP Chemistry

What people get wrong

1. Forgetting to convert percent abundances to decimals before computing a weighted average atomic mass — always divide by 100 first, or the result will be off by roughly a factor of 100.
2. Filling 3d before 4s when forming a transition metal cation — electrons are removed from the highest principal quantum number (4s) first, so Fe3+ is $[Ar]3d^5$, not $[Ar]4s^2 3d^3$.
3. Confusing PES peak height with peak position — height tells you the number of electrons in a subshell, position tells you binding energy; mixing these up leads to misidentifying elements.
4. Assuming ionization energy rises smoothly across every period — the Group 2→13 (e.g., Mg to Al) and Group 15→16 (e.g., N to O) dips are frequently tested and must be explained by subshell stability, not treated as errors.

What people get wrong

1. Forgetting to add or subtract electrons for the ion's overall charge when counting valence electrons for a polyatomic ion; always adjust the total electron count before drawing the skeleton.
2. Treating one resonance structure as 'the' correct structure and assigning fixed bond lengths to it; instead recognize the resonance hybrid has bond lengths and orders intermediate between the individual structures.
3. Confusing electron-domain geometry with molecular geometry when lone pairs are present on the central atom; state both separately, since lone pairs change the molecular shape (e.g., bent vs. linear) but not the electron-domain arrangement.
4. Assuming formal charge of zero everywhere is always achievable; some structures require formal charges, and the goal is minimizing magnitude and placing negative charge appropriately, not forcing all charges to zero.

What people get wrong

1. Labeling any polar bond as 'hydrogen bonding' — hydrogen bonding only occurs when H is directly bonded to N, O, or F. Check for these specific bonds before invoking hydrogen bonding as the explanation.
2. Assuming a heavier molecule always has stronger intermolecular forces — dispersion forces do scale with molar mass, but a lighter polar or hydrogen-bonding molecule can still have a higher boiling point than a heavier nonpolar one. Compare the strongest force present, not just mass.
3. Using R with mismatched pressure units (plugging kPa into an equation set up for atm, or forgetting to convert Celsius to Kelvin) in PV=nRT. Always convert temperature to Kelvin and match R's units to the given pressure unit.
4. Believing gas solubility increases with temperature the way solid solubility usually does. Gas solubility in liquids decreases as temperature rises because dissolved gas molecules gain enough kinetic energy to escape the solution.

What people get wrong

1. Assuming every titration ratio is 1:1 — check the balanced equation first; polyprotic acids or bases with more than one reactive OH- change the ratio, as in H2SO4 + 2 NaOH.
2. Forgetting to convert volumes from mL to L before using M = mol/L, which produces an answer off by a factor of 1000.
3. Writing a net ionic equation without first checking solubility rules, so a spectator ion is left in (or a species that should dissociate is kept molecular) — always confirm which compounds are strong electrolytes before canceling ions.
4. Confusing the oxidizing agent with the species that is oxidized — the oxidizing agent is the substance that gets reduced (gains electrons), even though it 'causes' oxidation in the other species.

What people get wrong

1. Assuming reaction orders match the coefficients of the overall balanced equation — this is only true for elementary steps. Always determine orders from experimental data for the overall reaction.
2. Leaving a reaction intermediate in the final predicted rate law after using the rate-determining step. Instead, use the fast pre-equilibrium (or the step that produced the intermediate) to rewrite it in terms of original reactants.
3. Confusing which linear plot corresponds to which order (mixing up ln[A] vs t with 1/[A] vs t). Instead, memorize: zero order is linear as [A] vs t, first order as ln[A] vs t, second order as 1/[A] vs t.
4. Believing a catalyst changes ΔH or shifts the equilibrium position. Instead, remember a catalyst only lowers Ea for both directions and speeds up reaching the same equilibrium, without changing K or ΔH.

What people get wrong

1. Writing ΔH as positive for an exothermic process because energy is 'released.' Instead, remember the sign describes the system: exothermic reactions lose enthalpy, so ΔH is negative, even though the surroundings gain heat.
2. In calorimetry, forgetting the sign flip between the two components (treating q_metal and q_water as equal instead of opposite). Instead, always start from $q_{lost} = -q_{gained}$ before substituting numbers.
3. In bond-enthalpy calculations, computing (bonds formed) − (bonds broken) instead of the reverse. Instead, remember breaking bonds always costs energy (positive) and forming bonds always releases energy (subtracted), so it's always broken minus formed.
4. When reversing or scaling an equation for Hess's law, changing the equation but forgetting to apply the identical operation to its ΔH. Instead, treat the equation and its ΔH as a single unit — whatever you do to one, do to the other.

What people get wrong

1. Including pure solids or liquids in the K or Ksp expression. Instead, cross them out immediately after writing the balanced equation — only gases and aqueous species appear.
2. Assuming the small-x approximation always works. Instead, check whether x is under about 5% of the initial concentration after solving; if not, redo the problem with the quadratic formula.
3. Confusing Q with K by plugging in initial concentrations when the problem actually wants equilibrium concentrations, or vice versa. Instead, always identify whether the given concentrations are at equilibrium before deciding whether you're computing Q or K.
4. Treating a catalyst or an inert gas added at constant volume as a stress that shifts equilibrium. Instead, remember a catalyst changes only the rate to reach equilibrium, and an inert gas at constant volume doesn't change any partial pressure of the reacting species, so K and the position of equilibrium are unaffected.

What people get wrong

1. Treating a weak acid like a strong acid by setting [H⁺] equal to the initial concentration — instead, build an ICE table and solve for x using Ka, since weak acids only partially dissociate.
2. Assuming pH + pOH = 14 always holds — this depends on Kw = 1.0×10⁻¹⁴, which is only true at 25°C; at other temperatures Kw changes and so does the sum.
3. Assuming the half-equivalence point always gives pH = pKa — this shortcut only applies to weak acid/strong base (or weak base/strong acid) titrations, not to strong-strong titrations which have no meaningful 'pKa'.
4. Ranking acid strength purely by electronegativity — for binary acids down a group, decreasing bond strength (not increasing electronegativity) controls the trend, so HI is a stronger acid than HF.

What people get wrong

1. Flipping a sign when a half-reaction is written as oxidation instead of reduction. Always look up the reduction potential as tabulated and use E°cell = E°cathode − E°anode without reversing any table values.
2. Forgetting to convert ΔS° from J/(mol·K) to kJ/(mol·K) before combining it with ΔH° in kJ/mol in the Gibbs equation. Keep units consistent throughout, or convert ΔH° to joules instead.
3. Using the 0.0592/n shortcut in the Nernst equation at a temperature other than 25°C. That constant already has T = 298 K built in; at any other temperature you must use the full RT/nF form.
4. Multiplying E° values by the number of electrons when combining half-reactions. Cell potential is an intensive property and does not get multiplied even when you scale a half-reaction to balance electrons — only ΔG° scales with n.