Error Analysis Table

Common Mistake	Why It Is Wrong	Correct Understanding
Using 22.4 L/mol for gas at any condition	22.4 L/mol applies ONLY at STP (0 °C, 1 atm); at 25 °C it is ~24.5 L/mol	Always check the given temperature and pressure before using molar volume
Confusing empirical formula mass with molar mass	The empirical formula gives the simplest ratio, not the actual molecular mass	n = molar mass ÷ EF mass; multiply EF by n to get molecular formula
Assuming both reactants are limiting when ratio fits	Students identify 1:1 ratio and declare both limiting without checking stoichiometry	Check the stoichiometric coefficients — 2 $H_{2}$ : 1 $O_{2}$ means you need 2 mol $H_{2}$ per 1 mol $O_{2}$
Using molarity for colligative property calculations	Molarity is temperature-dependent, introducing error in boiling/freezing point shifts	Use molality, which is temperature-independent
Forgetting to subtract solute mass from solution mass for molality	Students divide by total solution mass instead of solvent mass	Solvent mass = solution mass − solute mass
Using molar mass directly as equivalent weight	Equivalent weight = molar mass / n-factor; n-factor depends on reaction	For $H_{2}$$SO_{4}$ : eq. wt = 98/2 = 49 g/eq in neutralisation
Rounding atomic ratios too aggressively in empirical formula	0.98 or 1.02 can be rounded to 1, but 1.48 or 1.52 cannot	Only round values within ±0.05 of a whole number; others indicate a fraction needing multiplication
Treating percentage by mass as mole fraction	% by mass is mass ratio; mole fraction is moles ratio	Convert mass % to moles using molar mass before computing mole fraction
Forgetting the factor of 1000 in molarity formula	The formula M = (d × w% × 10) / M_r comes from converting mL to L and % to fraction	Always track the 1000 mL/L conversion; use M = (1000 × d × w%) / (M_r × 100)
Confusing Normality with Molarity when n-factor ≠ 1	N = M only when n-factor = 1	Always compute n-factor before writing normality