Tamil Nadu HSE First Year Chemistry common mistakes
Tamil Nadu HSE First Year Chemistry · Understand the traps and how to avoid them.
When converting between units in stoichiometry problems (e.g., grams to moles, litres of gas at STP to moles), students plug numbers into the formula without cancelling units, then report an answer in the wrong unit or with an extra factor of 1000.
In Tamil Nadu HSE, the numerical problems in Unit 1 are the first place dimensional analysis is required, and most school practice stops at 'apply the formula'. Students memorise n = m/M and n = V/22.4 but never practise writing the unit alongside each quantity, so a mass in mg or a volume in mL silently becomes wrong by a factor of 1000.
What to do instead
For every calculation, write each given quantity with its unit and convert to base units (g, L, mol) before substituting. Use the 'unit chain' method: write the conversion as a fraction so unwanted units cancel visibly. Do 10 mixed-unit drills (mg→g, mL→L, ppm→mol) and check that the final unit matches what the question asks for.
Students write electron configurations using the n+l (Aufbau) filling order but then assign the last electrons to the wrong subshell when the configuration is for an ion, especially for transition-metal-like cases such as Fe²⁺ or Cu⁺, filling 4s before 3d is emptied.
The textbook rule '4s fills before 3d' is learned as a filling rule, and students apply it in reverse when removing electrons. They forget that once 3d is occupied, 4s electrons are removed first because 4s is higher in energy in the ion. The aufbau diagram is memorised as a picture, not as an energy argument.
What to do instead
Learn two separate rules: filling order (n+l rule, 4s before 3d) and removal order (remove from the highest n shell first, so 4s before 3d for first-row transition metals). Practise Fe→Fe²⁺→Fe³⁺ and Cu→Cu⁺→Cu²⁺ explicitly, and justify each removal by writing the (n+l) values of the orbitals involved.
When predicting periodic trends, students state that atomic radius increases across a period or that ionisation enthalpy increases down a group, because they confuse the direction of the trend with the direction of the argument ('more protons' is read as 'bigger').
The trends are taught as lists to memorise, not as consequences of effective nuclear charge and shielding. Under exam pressure, students recall the phrase 'increases across a period' but attach it to the wrong property, or they recall that 'ionisation enthalpy is high for noble gases' and generalise incorrectly down the group.
What to do instead
For each trend, write the one-sentence causal chain: across a period, Z increases while shielding stays roughly constant, so Z_eff increases, so radius decreases and ionisation enthalpy increases. Down a group, n increases, so shielding and radius increase, so ionisation enthalpy decreases. Then test yourself with the exceptions (Be>B, N>O, Mg>Al, P>S) and explain each by electron configuration.
Students draw Lewis structures and then predict molecular shape without counting lone pairs on the central atom, so H₂O is described as linear, NH₃ as trigonal planar, and SF₄ as tetrahedral, and bond angles are quoted from the wrong geometry.
VSEPR is often reduced to matching a formula to a shape from a table. Students count bond pairs but overlook lone pairs, or they treat a lone pair as having the same repulsion as a bond pair. The distinction between electron-pair geometry and molecular geometry is not internalised.
What to do instead
Use a fixed four-step routine: (1) total valence electrons, (2) draw the skeleton and place bonding pairs, (3) place remaining electrons as lone pairs on the central atom, (4) count electron domains (bond pairs + lone pairs) to get electron-pair geometry, then remove lone pairs to get molecular geometry. Always state the lone-pair repulsion order lp–lp > lp–bp > bp–bp and adjust the angle (e.g., H₂O ≈104.5°, NH₃ ≈107°).
In gas-law and kinetic-theory numericals, students substitute Celsius temperatures directly into pV = nRT or into Graham's law and into RMS-speed expressions, and they mix up the units of R (0.0821 L atm K⁻¹ mol⁻¹ vs 8.314 J K⁻¹ mol⁻¹).
School problems frequently use STP values where 273 K is memorised, so the conversion step is skipped. When a problem gives 27 °C, students write 27 in place of T. The choice of R is treated as a constant to memorise rather than a unit-dependent quantity.
What to do instead
Adopt a rule: every temperature in a gas equation is written as (t + 273) K before any substitution. Choose R by inspecting the units of p and V in the problem: atm and L → 0.0821; Pa and m³ → 8.314; and convert p and V to match. Do a set of problems at 27 °C, 127 °C and −73 °C to make the conversion automatic.
Students calculate ΔH for a reaction using Hess's law or bond enthalpies but forget to reverse the sign when a reaction is reversed, or forget to multiply ΔH by the stoichiometric coefficient when an equation is scaled, and they confuse ΔH with ΔU.
Hess's law is taught as 'add the equations', and students add the enthalpy numbers mechanically without tracking the operations applied to each equation. Bond-enthalpy problems compound the error because bonds broken are endothermic (+) and bonds formed are exothermic (−), and students apply one sign to all terms.
What to do instead
Write each manipulation next to the equation: '×2', 'reversed', 'kept'. Apply the same operation to the ΔH value and record it in a column. For bond enthalpies, always compute ΔH = Σ(bonds broken) − Σ(bonds formed) and label each term with its sign before summing. Separately, state that ΔH = ΔU + Δn_gRT and identify when Δn_g = 0.
When solving equilibrium problems, students write the equilibrium-constant expression with products over reactants but include pure solids and liquids, or they substitute initial concentrations into Kc instead of equilibrium concentrations, and they treat a large K as 'reaction goes to completion' without checking.
The expression is memorised as a pattern rather than as a ratio of activities at equilibrium. ICE tables are often skipped in school practice, so students cannot distinguish the initial row from the equilibrium row. The meaning of K as a ratio at equilibrium is not connected to the arithmetic.
What to do instead
Always build an ICE table (Initial, Change, Equilibrium) and substitute only the Equilibrium row into Kc. Cross out pure solids and pure liquids in the expression and state why (their activity is 1). Interpret K qualitatively: K >> 1 favours products, K << 1 favours reactants, and the reaction is never 'complete' unless a product leaves the system.
Students compute pH from a weak acid or weak base using [H⁺] = c·α without checking whether α is small, and they apply the pH formula to a strong base without first converting pOH to pH, or they ignore the contribution of water at very low concentrations.
The Ostwald dilution law and the pH definition are learned as separate formulas. Students do not connect Ka = cα² with the approximation α << 1, and they forget that pH + pOH = 14 applies at 25 °C. Buffer problems then fail because the Henderson–Hasselbalch form is applied without identifying the conjugate pair.
What to do instead
For weak acids, first write Ka = cα²/(1−α) and test whether c/Ka > 100; only then use α = √(Ka/c). Compute [H⁺] = cα, then pH = −log[H⁺]. For bases, compute [OH⁻], then pOH, then pH = 14 − pOH at 25 °C. For buffers, explicitly label the acid and its conjugate base before substituting into pH = pKa + log([salt]/[acid]).
In solutions and colligative-property problems, students use the molarity of the solute as the molality, or they forget the van't Hoff factor i for electrolytes, so ΔTb and ΔTf for NaCl are calculated as if it were a non-electrolyte.
Molarity and molality are both 'concentration' and differ only in the denominator (volume of solution vs mass of solvent). Students memorise the formulas ΔTb = Kb·m and ΔTf = Kf·m but do not examine what m means. Electrolytes are treated as single particles because dissociation is a separate chapter.
What to do instead
Before any colligative calculation, convert the given concentration to molality: moles of solute per kilogram of solvent, using the mass of solvent, not the mass of solution. Then multiply by i, the number of ions per formula unit (NaCl → 2, CaCl₂ → 3, glucose → 1). For non-ideal cases, note that i is measured and may be less than the theoretical value due to ion pairing.
In chemical kinetics, students write the rate law from the stoichiometric equation (e.g., for 2A + B → products, rate = k[A]²[B]) instead of from experimental data, and they confuse order with molecularity or with the stoichiometric coefficient.
The rate law is often presented after the equation, so students assume the coefficients are the orders. The distinction between elementary and complex reactions is not emphasised, and order is defined only as 'the power in the rate law' without practice in deriving it from initial-rate data.
What to do instead
Treat the rate law as an experimental result: use the initial-rate method, comparing two runs where only one concentration changes, to find each order. State explicitly that order can be zero, fractional or negative and need not match coefficients; molecularity applies only to an elementary step. Practise deriving orders from a table of initial rates before writing k with its units.
In organic chemistry, students apply Markovnikov's rule to the wrong carbon, or they write a carbocation intermediate without checking stability, and they misapply the IUPAC rules by choosing the longest chain that is not the parent or by numbering from the wrong end.
Markovnikov's rule is memorised as 'H goes to the carbon with more H' without the carbocation-stability justification, so it fails for cases with resonance or rearrangement. IUPAC naming is practised on simple alkanes only, so the priority rules for multiple bonds, substituents and principal functional groups are not automatic.
What to do instead
For addition reactions, draw the two possible carbocations and select the more stable one (3° > 2° > 1°, allylic/benzylic stabilised), then place H on the other carbon; this also predicts rearrangements. For nomenclature, follow a fixed order: identify the principal functional group, then the longest chain containing it, then number to give the lowest locants to the principal group and multiple bonds, then list substituents alphabetically.
When identifying or separating organic compounds by qualitative analysis, students describe a positive test without the specific observation (colour, precipitate, gas) and they confuse tests for functional groups, e.g., using Tollens' reagent for a ketone or bromine water for an alkane.
The tests are memorised as a list of reagent–compound pairs, without the observable change or the chemistry behind it. Under exam conditions, students recall that 'Tollens' is for aldehydes' but cannot state the silver mirror, and they generalise 'bromine water decolourises' to all hydrocarbons rather than to unsaturation.
What to do instead
Build a table with three columns: functional group, reagent, and exact observation (e.g., aldehyde + Tollens' → silver mirror; unsaturation + Br₂ in CCl₄ → colour discharged; carboxylic acid + NaHCO₃ → brisk effervescence of CO₂). Write the underlying reaction for each. Then practise distinguishing pairs such as aldehyde vs ketone, and alkene vs alkane, by naming the test and the observation.