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Compounds and Stoichiometry

Chapter 4: Compounds and Stoichiometry

5 min read Updated Mar 26, 2026
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🎯 Diagnostic: Test Your Starting Level 24 questions (2 per section). No prior reading required, see what you already know.

Aim to answer every question before checking. Missed questions point you to the sections you need most.

1. (4.1) The empirical formula of a compound represents:
D. Multiple molecular formulas can share the same empirical formula (e.g., CH₂O empirical for formaldehyde, acetic acid, glucose).
2. (4.1) The empirical formula of glucose (C₆H₁₂O₆) is:
A. Divide each subscript by the greatest common factor (6) to get 1:2:1 (CH₂O). Matches the name "carbohydrate" = carbon + water.
3. (4.2) Fe₂O₃ is named:
B. Roman numerals indicate oxidation state. Fe in Fe₂O₃ is +3 (since oxygen is -2 and total charge is zero).
4. (4.2) CO₂ is named:
C. Binary molecular compounds use Greek prefixes (mono is dropped from the first element if only one). CO₃²⁻ is the carbonate anion, which is different.
5. (4.3) One mole of a substance contains:
D. The mole bridges microscopic (atoms/molecules) and macroscopic (grams). Molar mass in g/mol has the same numerical value as atomic/molecular mass in amu.
6. (4.3) To convert grams to moles you:
A. Units check: g ÷ (g/mol) = mol. From moles, multiply by Avogadro's number to get particles.
7. (4.4) Percent composition by mass of an element in a compound equals:
B. Easiest in one mole of compound: grab the element's total mass (coefficient × atomic mass) over the molar mass.
8. (4.4) In water (H₂O, M = 18 g/mol), hydrogen is what percent by mass?
C. H accounts for only ~2 of the 18 g/mol; oxygen dominates mass despite being a single atom.
9. (4.5) Balanced chemical equations must respect:
D. Atoms of each element must match on both sides, and total charge must balance for ionic or redox equations.
10. (4.5) To balance H₂ + O₂ → H₂O, the smallest integer coefficients are:
A. Check: 4 H and 2 O on each side.
11. (4.6) AgNO₃ + NaCl → AgCl + NaNO₃ is classified as:
B. Cations and anions swap partners. AgCl falls out as a white solid precipitate.
12. (4.6) 2 H₂O₂ → 2 H₂O + O₂ is an example of:
C. One reactant breaks into multiple products. Catalase accelerates this in biological systems.
13. (4.7) The limiting reagent is:
D. After the limiting reagent is gone, the reaction cannot continue even if other reactants remain.
14. (4.7) To identify the limiting reagent, you should:
A. Stoichiometry works in moles, not grams. The reactant with the smallest "effective moles" is the bottleneck.
15. (4.8) Percent yield is calculated as:
B. Real yields are usually lower than theoretical because of side reactions, losses during workup, and incomplete conversion.
16. (4.8) If theoretical yield is 10 g and you obtain 7.5 g, the percent yield is:
C. 7.5 / 10 × 100% = 75%.
17. (4.9) Molarity (M) is defined as:
D. Molarity changes with temperature because volume does. Molality does not, which is why colligative properties use molality.
18. (4.9) Molality (m) is defined as:
A. Based on mass, not volume, so molality is temperature-independent. Useful in boiling point elevation and freezing point depression.
19. (4.10) The dilution equation M₁V₁ = M₂V₂ reflects:
B. Moles before = moles after: M×V is constant when only solvent is added.
20. (4.10) Diluting 50 mL of 2.0 M solution to 200 mL final volume gives:
C. Quadrupling the volume quarters the concentration.
21. (4.11) Solution stoichiometry relies on:
A. Flow: M × V → moles of one species, use the coefficients of the balanced equation, then convert moles of the other species back to whatever unit is asked.
22. (4.11) In an acid-base titration, the equivalence point is reached when:
A. Equivalence point pH depends on the strength of the acid/base (pH = 7 only for strong-strong). Indicator endpoints approximate the equivalence point.
23. (4.12) Density is defined as:
B. For water, 1 g/mL = 1 kg/L. Useful as a bridge between mass and volume in solution problems.
24. (4.12) Dimensional analysis refers to:
C. Treating units as algebraic quantities catches errors: if the units do not work out, the setup is wrong.

You follow a recipe every time you cook. Two eggs, one cup of flour, half a teaspoon of salt. If you run out of eggs after cracking one, it does not matter that you have a mountain of flour - you are stuck. That single missing egg limits the entire batch.

Chemistry works the same way. Atoms combine in fixed ratios to form compounds, and when those compounds react, the amounts matter. Stoichiometry is the math that connects what you start with to what you end up with. It answers the questions that matter on the MCAT: How much product can you make? Which reactant runs out first? What concentration is this solution?

This chapter is the quantitative backbone of general chemistry. Every topic here - from the mole concept to dilution calculations - shows up directly in MCAT passages. A passage might describe a combustion analysis experiment and ask you to determine an unknown compound’s molecular formula. Another might give you reactant masses and ask for the theoretical yield. Yet another might describe a titration and ask you to calculate the molarity of an unknown acid. The math is straightforward once you see the pattern, and the pattern is always the same: convert to moles, use the ratio, convert out.

The Kitchen is a Chemistry Lab

Think of stoichiometry as cooking with a periodic table. Your recipe (the balanced equation) tells you the exact ratio of ingredients (reactants) needed to produce your dish (products). Molar mass is your measuring cup - it converts between the grams you weigh on a balance and the moles that reactions actually “count” in. The limiting reagent is whichever ingredient runs out first, and it determines how much food you can make. Percent yield tells you how close your actual output came to the recipe’s promise - because in real kitchens and real labs, things never go perfectly.

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