# Why are yields of some chemical reactions so dependent on the stoichiometry of the reagents? [closed]

Back in chemistry class in high school I was working with a partner doing a chemical reaction. We measured the amounts of reagents well, but not perfect, and our yield suffered greatly. On the other hand my friend and his partner came up with quite a high yield. He told me it was because they measured each amount of reagent exactly.

So I wonder... Why are these reactions so dependent on perfect measuring of reagents? Why can't you just use more than enough of one reagent to completely react with the other and remove the excess reagent?

Also does this apply to all chemical reactions, or just a select few?

• It depends..... – MaxW Jun 15 '19 at 5:51
• And I'll say your friend and his partner got lucky. Or rather you did some small but crucial mistake they didn't. A lot of reasons come to mind, but exact stochiometry of the reagents is not among the first. – Karl Jun 15 '19 at 10:57

The proper stoichiometry may not be the only reason for the friend's success, but also the overall preciseness of their preparation procedure.

Chemical reactions are not generally that much critical, but there are lot of cases, that are.

There are few aspects to consider:

1. Excessive reagents may affect the effectiveness of the product isolation.
2. Excessive reagents may alter chemical environment in final mixture, affecting formation of products.
3. Excessive reagent may not remain, but may form another product, instead of the desired one

Typical case for 3. is e.g reaction of $$\ce{NaOH}$$ with $$\ce{H3PO4}$$. Phosphoric acid forms 3 salts.

$$\ce{NaH2PO4 + Na2HPO4 + Na3PO4}$$

Let suppose you want to produce $$\ce{NaH2PO4 }$$

$$\ce{NaOH + H3PO4 -> NaH2PO4 + H2O}$$

Extra $$\ce{NaOH}$$ will create from the wanted salt other salt...

$$\ce{NaOH + NaH2PO4 -> Na2HPO4 + H2O}$$

... and the amount of the final product decreases.