Theoretical Yield from the Limiting Reactant

Two reactants go into a flask, and one of them runs out first. That one, the limiting reactant, decides how much product you get; the other is left over. The test is to divide each reactant's moles by its coefficient in the balanced equation. Each quotient counts how many times that reactant could run the equation as written, and the smaller count wins.

Worked: 2H₂ + O₂ → 2H₂O, with 3.00 mol of hydrogen and 2.00 mol of oxygen. Hydrogen supplies 3.00/2 = 1.50 batches and oxygen supplies 2.00/1 = 2.00. Hydrogen runs out first, so the water is 2 × 1.50 = 3.00 mol, and 2.00 − 1.50 = 0.50 mol of oxygen is left in the flask.

The reactant with fewer moles is not always the limiting one. In the example, oxygen had fewer moles and was still in excess, because the equation asks for twice as much hydrogen. Compare moles divided by coefficient, never raw moles, and never grams: comparing masses ignores both the molar masses and the coefficients.

The answer is a theoretical yield, a ceiling on paper. What the balance shows afterwards is the actual yield, and the ratio of the two is the percent yield. Working backwards only goes one branch at a time: given the yield, the calculator can tell you the least A that makes it, but not how much A was really there, because once B runs out any extra A changes nothing.