If a recipe calls for 2 eggs per 1 cup of flour, you'd never use 5 eggs with 1 cup of flour and expect a normal cake. Stoichiometry is exactly that — but for chemical reactions. It tells you how much of each substance you need and how much product you'll get.
Every balanced chemical equation is a recipe written in moles. Stoichiometry is the set of rules that lets you use those mole ratios to calculate masses, volumes, or concentrations of reactants and products. It's the single most practical skill in introductory chemistry — and once you see the pattern, it never really changes.
Mole Ratio
The coefficients in a balanced equation give you the exact ratio of moles that react and form.
Limiting Reactant
The reactant that runs out first controls how much product you can possibly make.
Percent Yield
Real reactions never give 100% of the theoretical product — the gap is your percent yield.
What Stoichiometry Actually Means
THE BIG PICTUREThe word comes from the Greek stoicheion (element) and metron (measure). In practice, it's the bridge between the microscopic world of atoms and the macroscopic world of grams and litres.
Here's the key idea: a balanced chemical equation tells you the ratio in which particles react. For example:
That same ratio applies to moles: 2 moles of H₂ react with 1 mole of O₂ to produce 2 moles of H₂O. And because moles convert to grams using molar mass, the ratio also tells you the mass relationship.
Before you do any stoichiometry calculation, the equation must be balanced. An unbalanced equation gives you the wrong mole ratio, which makes every calculation wrong from that point forward.
Mole Ratios: The Engine of Stoichiometry
USING COEFFICIENTSA mole ratio is simply a fraction built from the coefficients of two substances in a balanced equation. Once you have the balanced equation, you can write a mole ratio between any two substances in it.
For the reaction N₂ + 3H₂ → 2NH₃:
- Mole ratio of N₂ to H₂ is 1:3 → 1 mol N₂ / 3 mol H₂ or 3 mol H₂ / 1 mol N₂
- Mole ratio of H₂ to NH₃ is 3:2 → 3 mol H₂ / 2 mol NH₃
- Mole ratio of N₂ to NH₃ is 1:2 → 1 mol N₂ / 2 mol NH₃
You choose which ratio to use based on what you're given and what you need to find. The substance you know goes on the bottom; the substance you want goes on top.
Quick check: reading a balanced equation
Consider the combustion of methane:
If you have 3 moles of CH₄, the ratio tells you immediately that you need 6 moles of O₂, and you'll produce 3 moles of CO₂ and 6 moles of H₂O. No unit conversion needed yet — that's the beauty of working in moles.
Limiting Reactant: Who Runs Out First?
THE BOTTLENECKIn the real world, reactants aren't always present in the exact ratio the equation demands. When one reactant runs out before the others, the reaction stops — no matter how much of the other reactants remain.
The reactant that runs out first is the limiting reactant. It controls the maximum amount of product that can form. The others are in excess.
How to find the limiting reactant — every time
- Balance the chemical equation.
- Convert the amount of each reactant to moles.
- Divide each reactant's moles by its coefficient in the balanced equation.
- The smallest result is the limiting reactant.
- Use the limiting reactant's moles to calculate product amounts.
You can't just compare moles directly. A reactant with a larger coefficient needs proportionally more moles to react. Dividing by the coefficient normalises everything to the same scale, so the comparison is fair.
Worked example
Consider the reaction 2H₂ + O₂ → 2H₂O. You have 4 moles of H₂ and 3 moles of O₂. Which is limiting?
- Balanced equation: 2H₂ + O₂ → 2H₂O
- Moles of H₂ = 4; moles of O₂ = 3.
- Divide by coefficients: H₂ → 4 ÷ 2 = 2. O₂ → 3 ÷ 1 = 3.
- Smallest is 2, so H₂ is the limiting reactant.
- O₂ is in excess. Product formed depends on H₂: 4 moles H₂ × (2 mol H₂O / 2 mol H₂) = 4 moles H₂O.
| Reactant | Moles present | Coefficient | Moles ÷ coefficient | Result |
|---|---|---|---|---|
| H₂ | 4 | 2 | 2 | Limiting |
| O₂ | 3 | 1 | 3 | Excess |
Percent Yield: What You Get vs What You Expected
REALITY CHECKStoichiometry gives you the theoretical yield — the maximum product possible if everything reacted perfectly. In a real lab or factory, you almost never get that much. Side reactions, incomplete reactions, and losses during purification all reduce the amount you actually collect.
A percent yield above 100% usually means the product isn't pure — there's solvent or another impurity adding mass. A yield below 100% is normal and expected.
Worked example
Stoichiometry predicts 20 g of product, but you isolate only 15 g. What's the percent yield?
- Formula: % yield = (actual / theoretical) × 100
- Substitute: (15 g ÷ 20 g) × 100
- Solve: 75%
Interactive Mole Calculator
Use the sliders below to see how moles, mass, and molecules connect. This isn't a reaction-specific calculator — it's the fundamental conversion that underpins all stoichiometry.
Mass = moles × molar mass. Particles = moles × Avogadro's number (6.022 × 10²³).
Where Stoichiometry Shows Up in Real Life
- Medicine: calculating exact drug doses based on reaction yields and purity.
- Manufacturing: determining how much raw material to buy so nothing runs out mid-production.
- Environmental science: predicting how much CO₂ a given amount of fuel will release.
- Food science: scaling recipes in industrial food production while keeping chemical ratios correct.
- Energy: calculating fuel-to-energy ratios for combustion engines and power plants.
Convert what you're given to moles → use the mole ratio from the balanced equation → convert moles to whatever unit the question asks for. That three-step rhythm handles almost every stoichiometry problem you'll meet in a first-year chemistry course.