Question #94538

1 Answer
Nov 8, 2015

Answer:

#67.6%#

Explanation:

Percent yield are all about determining how efficient a chemical reaction is, or, in other words, how much it deviates from the ideal situation.

For any chemical reaction, you can predict how much product will result from known quantities of reactants by using stoichiometric calculations.

More specifically, you examine in what mole ratio the reactants will react and how many moles of the products you can expect. Now, stoichiometric calculations are always done assuming that the reaction has a #100%# yield.

This is what the theoretical yield actually is - the amount of product you would see if absolutely every molecule of the reactants takes part in the reaction and forms products.

However, this is not going to happen in practice. A reaction's actual yield tells you close to #100%# the reaction comes.

In your case, you did the stoichiometric calculations for a #100%# yield and determined that you should collect #"36.7 g"# of product.

However, you actually perform the reaction and only get #"24.8 g"# of the product. So it's obvious that not every molecule of reactants actually ended up producing products.

The equation for percent yield is

#color(blue)("% yield" = "actual yield"/"theoretical yield" xx 100)#

Simply put, what you get in practice divided by what you should get in ideal conditions, i.e. in theory, or "on paper".

So, plug in your values to get

#"% yield" = (24.8color(red)(cancel(color(black)("g"))))/(36.7color(red)(cancel(color(black)("g")))) xx 100 = color(green)(67.6%)#

This can be interpreted as - for every #100# molecules that should rect to give product, only #67.6# actually do so.