5 Calculation Mistakes That Cost You Points on Every Chemistry Exam

I have graded thousands of chemistry problems across high school, college, and graduate courses. These five mistakes show up on every single one. They're not exotic errors; they're fundamental slip-ups that cost students points on problems they otherwise know how to solve. Every one of them is preventable, and every one of them shows up in every course I've ever taught or tutored.

Mistake 1: Mixing up significant figures rules.

Multiplication and division: your answer gets the same number of significant figures as the measurement with the fewest sig figs. Addition and subtraction: your answer gets the same number of decimal places as the measurement with the fewest decimal places. These are different rules, and I cannot tell you how many exam papers I've seen where a student applies the decimal-place rule to a multiplication problem. The worst part is that this error cascades; one misapplied rule in Step 1 means every subsequent calculation carries the wrong number of digits, and suddenly you've lost points on four questions instead of one. The fix is simple: before you round, stop and ask yourself whether you're multiplying/dividing or adding/subtracting, then apply the correct rule.

Mistake 2: Forgetting to convert units before plugging in.

Volume in mL instead of L going into PV = nRT. Celsius instead of Kelvin in the ideal gas law. Kilojoules instead of joules in a Hess's Law problem. Every single one of these gives you an answer that looks plausible but is off by orders of magnitude, or, worse, off by just enough that you don't notice. Before you touch your calculator, check every value and confirm its unit matches what the equation expects. This takes ten seconds and prevents a completely wrong answer. I tell my students to circle every unit before they start computing. It sounds elementary, but the number of 400-level students I've seen make this mistake would surprise you.

Mistake 3: Comparing grams instead of moles in limiting reagent problems.

10.0 g of A and 15.0 g of B does not mean "there's less A." If A has a molar mass of 10 g/mol (that's 1.0 mol) and B has a molar mass of 150 g/mol (that's 0.10 mol), B is your limiting reagent despite having more mass sitting on the scale. The reaction doesn't care about grams. It cares about moles, because stoichiometric coefficients are mole ratios. Always convert to moles, divide by stoichiometric coefficients, then compare. The smallest number is your limiting reagent. Mass is a liar in stoichiometry; moles tell the truth.

Mistake 4: Sign and scaling errors in Hess's Law.

Reverse a reaction? Flip the sign of ΔH. Multiply a reaction by a coefficient? Multiply ΔH by the same factor. Students routinely forget one or both, and the resulting answer is often just plausible enough to seem right, which is what makes this mistake so dangerous. My advice: write out every modified reaction with its modified ΔH on a separate line, verify that all intermediates cancel algebraically, and then add the ΔH values. Treat Hess's Law like an algebraic proof, not mental math. The extra two minutes of careful bookkeeping will save you from a sign error that tanks the entire problem.

Mistake 5: Using the wrong equilibrium expression.

This one is sneaky. A problem gives you Ka for a weak acid and asks for the pH of its conjugate base. You need Kb = Kw/Ka, not Ka. Plugging Ka directly into an ICE table for the conjugate base gives you a number that looks reasonable but is completely wrong. The same thing happens with Kp vs. Kc; students use whichever one is given without checking whether the problem requires the other. Before setting up any equilibrium calculation, ask yourself two questions: Which species is actually in solution? And which K describes that species' behavior? Get those right, and the math takes care of itself.

Spot the error.

A student calculates the pH of 0.20 M sodium acetate by setting up Ka = x²/(0.20 − x) = 1.8 × 10⁻⁵. What went wrong?

Sodium acetate is the conjugate base of acetic acid. It hydrolyzes in water, producing OH⁻, not H⁺. The student needs Kb = Kw/Ka = (1.0 × 10⁻¹⁴)/(1.8 × 10⁻⁵) = 5.6 × 10⁻¹⁰, and the ICE table should solve for [OH⁻], not [H⁺]. Classic Mistake 5: right setup, wrong K, wrong answer.