Introduction
In the last lesson you learned how to use a balanced equation to convert between moles, masses and volumes of reactants and products. Today we go a step further: what happens when you don't have exactly the right amounts of each reactant? This is where the ideas of limiting reactant, theoretical yield and percentage yield come in — essential tools for real laboratory and industrial chemistry (Talbot, p.418).
Why equations don't always react "perfectly"
A balanced equation tells you the ratio in which substances react, not how much of each you actually have in the flask. In most real reactions, one reactant runs out before the other. The reactant that is used up first is called the limiting reactant, because it limits how much product can form. The other reactant, left over at the end, is the excess reactant.
Method for finding the limiting reactant:
- Convert all given masses/volumes into moles (n = m / M or n = cV).
- Divide each mole amount by its coefficient in the balanced equation.
- The reactant with the smallest value is the limiting reactant.
Worked Example 1
Reaction: N₂(g) + 3H₂(g) → 2NH₃(g)
Suppose you have 2.0 mol N₂ and 3.0 mol H₂.
- N₂: 2.0 / 1 = 2.0
- H₂: 3.0 / 3 = 1.0
H₂ gives the smaller number, so H₂ is the limiting reactant. All further calculations (mass of NH₃ formed, leftover N₂, etc.) must be based on H₂, not N₂.