Are Moles Really Necessary? Bro. Thomas McCullough, CSC St. Edward's University, Austin, TX 78704 Two universally accepted concepts in the study of chemistry are atomic weights of elements and molecular weights of wmpounds. An eauallvvalid but sometimes neglec& conrepi is the balanied edquatlon of a chemical reaction. The complete combustion of methane to yield C 0 2 and water will always be written as CH, + 20, + COZ + 2Hz0 However, the step from compounds to balanced equations often encounters a detour in the form of "moles". Truly, moles are absolutely and fundamentally important, but they should not be allowed to divert ones attention from the equally valid and equally important balanced equation. In the synthesis of ammonia from nitrogen and hydrogen we write N, + 3Hz -i2NH2 28.0
6.0
34.0
with balanced equation weights (BEW)under each constituent. There is no need to invoke the use of moles at this point, which would involve noting that one mole of Nz requires three of hydrogen and will yield two of ammonia. The equation with its three BEWs does this "bookkeeping" for us. Suppose we need to know what weight of hydrogen is needed to react with 7.00 g of nitrogen. Converting both
hydrogen and nitrogen to moles requires an extra unnecessary step, a step often misjudged by beginning students. It is far simpler to set up a simple ratio based on the balanced equation as follows:
where X is the required weight of hydrogen. The numerators of the fraction are the actual weights (in grams); the denominators are the BEWs, also in grams. The theoretical yield (Y) of ammonia can be calculated in the same manner.
A similar approach can be used in determining the limiting amount of reactant. Using the previous balanced equation and given 10.0 g of nitrogen and 3.0 g of hydrogen, we again put the actual weight over the BEW and evaluate each fraction.
Nitrogen, having the smaller value, will be the limiting reagent.
Volume 69 Number 2 February 1992
121
