Limiting and Excess Reactants Percentage Yield Limiting and Excess Reactants

Some Definitions Limiting Reactant – The substance that controls the quantity of product that can form in a chemical reaction Excess Reactant – The substance that is NOT used up completely in the reaction Yield – The mass of product formed Theoretic – assuming no loss or error, the mathematical calculation of the mass of product formed Actual – the measured amount of the product from a reaction

Stoichiometry We use stoichiometry to determine: Limiting Reactant Excess Reactant Theoretic Yield

Limiting vs. Excess Reactant The Limiting reactant will form the LEAST amount of product. The Excess reactant will form the MOST amount of the product. We can determine these using stoichiometry if we know how much mass (or how many moles) we have of each reactant.

Example Zn + 2HCl  ZnCl2 + H2 If you combine 0.23 mol Zn and 0.60 mol HCl, would they react completely? In other words, will they produce the same amount of product, or will one produce more ZnCl2 than the other?

Example #2 PCl3 + 3H2O  H3PO3 + 3HCl Determine the limiting and excess reactants in this equation if you have 75.0 g PCl3 and 75.0 g H2O. How about 1.o mol of PCl3 and 50.0 g of H2O?

Determining Theoretic Yield The theoretic yield is what you determined when you determined the limiting reactant. This is the most product that can be produced in ideal circumstances. There will be instances when you will have more product than you should. Why might that be? There will be instances when you will have less product. Why??

Determining Actual Yield Actual yield is determined by measurement after the reaction is complete.

Determining Percentage Yield This is a ratio of the actual yield of a reaction to its theoretic yield. This ratio determines the efficiency of the reaction. Percent Yield = Actual Yield Theoretic Yield

Examples If the theoretic yield for a reaction is 738 g, and the actual yield is 591 g, what is the percentage yield? How many grams of CH3COOC5H11 should form if 4808 g are theoretically possible, and the percentage yield is 80.5%?

One more Example N2 + 3H2  2NH3 14.0 g N2 react with 3.15 g H2 to give an actual yield of 14.5 g of NH3. Determine the limiting reactant and the percentage yield for this reaction.

#6, p. 319 A chemist reacts 8.85 g of iron with an excess of hydrogen chloride to form hydrogen gas and iron (II) chloride. Calculate the theoretical yield and the percentage yield of hydrogen if 0.27 g of H2 are collected.