Try:
Balanced: 2H2 + O2 → 2H2O
Leave a reactant blank if it's in excess (e.g. oxygen from the air).
Limiting reagent: O2
| Reactant | Moles given | ÷ Coefficient | Used | Left over |
|---|---|---|---|---|
| H2 | 4.96 | 2.48 | 4.032 g | 5.968 g |
| O2limiting | 1 | 1 | 32 g | 0 g |
| Product | Theoretical yield (mol) | Theoretical yield (g) |
|---|---|---|
| H2O | 2 | 36.03 g |
The reactant with the smallest "moles ÷ coefficient" runs out first. That value (1) times each product's coefficient gives the moles of product that can form.
How it works
The equation is balanced with the same exact-arithmetic solver as the equation balancer. Each reactant's amount is converted to moles and divided by its coefficient; the smallest of those values is the "extent" of reaction. Multiplying the extent by each coefficient gives moles of product formed and moles of each reactant used.
Frequently asked questions
- What is a limiting reagent?
- The reactant that runs out first. Once it is used up the reaction stops, so it decides the maximum amount of product (the theoretical yield). Every other reactant is left partly unused — in excess.
- How do you find the limiting reagent?
- Balance the equation, convert each reactant’s mass to moles, and divide each by its coefficient in the balanced equation. The reactant with the smallest result is limiting. Comparing raw masses or raw moles does not work, because the coefficients differ.
- Is the reactant with the smallest mass always limiting?
- No. In 2H₂ + O₂ → 2H₂O, 10 g of hydrogen (about 5 mol) is far more than 32 g of oxygen (1 mol) needs, so oxygen is limiting even though its mass is larger.
- What is percent yield?
- Actual yield ÷ theoretical yield × 100. Real reactions usually give less than 100% because of side reactions, incomplete reactions and product lost during handling.
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