By Claire Gillespie – Updated Aug 30, 2022
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The excess reactant (also called the excess reagent) refers to the quantity of a chemical that remains unreacted after a reaction has reached completion. It is determined by the fact that the other reactant has been entirely consumed and therefore cannot participate further. Knowing the excess reactant allows you to determine the final amounts of both product and remaining reactant.
Start by balancing the chemical equation to understand the exact stoichiometric ratios required. For example, consider the reaction: Mg(OH)2 + HCl → MgCl2 + H2O. The unbalanced equation shows an imbalance of hydrogen, chlorine, and oxygen atoms. Adding a coefficient of 2 before HCl and 2 before H2O balances the equation: Mg(OH)2 + 2HCl → MgCl2 + 2H2O.
Convert the given masses of reactants to moles. Use a periodic table to obtain atomic masses: Mg = 24.305, O = 16.000, H = 1.008. For Mg(OH)2, the molecular weight is 24.305 + (2 × 16.000) + (2 × 1.008) = 58.305 g/mol.
Apply the formula moles = grams ÷ molecular weight. For 65 g of Mg(OH)2: 65 ÷ 58.305 = 1.11 mol. For 57 g of HCl (H = 1.008, Cl = 35.453): 57 ÷ 36.461 = 1.56 mol.
Use the stoichiometric coefficients from the balanced equation. Two moles of HCl react with one mole of Mg(OH)2. Calculate the limiting amount: 1.56 ÷ 2 = 0.78 mol of HCl required. Since 0.78 mol < 1.11 mol of Mg(OH)2, Mg(OH)2 is in excess, and HCl is the limiting reactant.
Find the fraction of Mg(OH)2 that actually reacted: 0.78 ÷ 1.11 = 0.704 (70.4%). Multiply the original mass of Mg(OH)2 by this fraction to obtain the amount used: 65 × 0.704 = 45.78 g. Subtract this from the initial mass to find the excess: 65 – 45.78 = 19.21 g of Mg(OH)2 remain unreacted.
