Limiting Reagent Lab

Give a reaction more of everything and it still stops. Something runs out first — the limiting reagent — and once it is gone, the reaction is over no matter how much of the other reactant is sitting there unused.

The recipe, not the pile

A balanced equation is a recipe. For 2H₂ + O₂ → 2H₂O, two moles of hydrogen need exactly one mole of oxygen. The ratio is 2:1, and it does not care which one you happen to have more of.

So equal amounts are usually wrong. With 5 mol of H₂ and 5 mol of O₂, the hydrogen runs out first — it would need 2.5 mol of oxygen but only gets to use 2.5 before the hydrogen is exhausted. Oxygen sits unused.

How to work it out

1. Work out how much of B your amount of A would need: nB_needed = nA × b / a.
2. If you have less B than that, B is limiting. Otherwise A is.
3. Product comes from the limiting reagent only: nC = (n_limiting / its coefficient) × c.
4. Convert to grams with the molar mass for the theoretical yield.

Theoretical versus actual yield

The theoretical yield is what the arithmetic promises. The actual yield is what you weigh at the end, and it is always lower: some product is lost transferring between containers, side reactions consume some starting material, and many reactions never quite finish because they reach equilibrium. The ratio, as a percentage, is the percentage yield — a key number in industrial chemistry, where a few percent can be worth millions.

Try this

• Hit “stoichiometric mix” — both reactants run out together and nothing is left over.
• Then shift one slider slightly. Watch a large pile of the other go unused.
• Compare hydrogen and oxygen: because O₂ is sixteen times heavier per mole, equal masses are wildly different mole counts.

Why it matters

Limiting reagent thinking decides how much of each feedstock a plant buys, why one reactant is deliberately made cheap and used in excess, how pharmaceutical yields are optimised, and even why a recipe fails when you scale it up incorrectly.