Themes > Science > Chemistry > Miscellenous > Help file Index > Stoichiometry > Percent yield Chemists, like all other people, aren't perfect. When a chemist does a synthesis, she will end up creating less product than expected because of spills, incomplete reactions, incomplete seperations, or a dozen other reasons. The percent yield is a way of measuring how successful a reaction has been. To compute the percent yield, figure out how much product you should have made by using basic stoichiometry. (Note: this may involve a limiting reagent problem.) Then simply divide the amount of stuff you did form by the expected amount and multiply by 100%. If you get a number > 100%, you've made a serious error someplace. Obviously, you want a high percent yield: if you have a ten step synthesis where the product from one reaction ends up as the reactants for the next and each synthesis has 90% yield, you'll end up with only ~35% yield for the overall reaction. Example: You burn 10.0 grams of methane in an excess of oxygen and form 19.8 grams of water. What was your percent yield? Solution : First, you need to find out how much product you would expect to make using basic stoichiometry. The reaction of methane with oxygen is shown below CH4(g) + 2O2(g) -> CO2(g) + 2H2O(g)You start with 10.0 grams of methane, which has a molecular weight of 16.04 g/mole, so you have 10.0 g/16.04 g/mole = 0.623 moles of methane. The ratio between methane and water is 2 water for every 1 methane, so you expect to form 0.623 moles CH4 * 2 moles H2O/1 mole CH4 1.25 moles H2ONow convert back to grams: the mole weight of water is 18.02 g/mole, so you should form 1.25 moles H2O * 18.02 g/mole = 22.5 g waterif the reaction had gone perfectly. You only formed 19.8 however, so your percent yield is % yield = mass created/mass expected * 100% % yield = 19.8 g/22.5 g = 88.0% Information provided by: http://learn.chem.vt.edu