Theoretical, actual and percentage yield

: the maximum possible mass of a product that a chemical reaction can make. It is calculated using molar ratios.

: the mass of a product that a chemical reaction makes in real life. It is usually less than the theoretical yield, for a number of reasons:

some of the product may be lost when the products are removed from the reaction mixture.
there might be side reactions – unwanted reactions that compete with the desired one.
the reactions may be reversible and may not go to completion. See also Redox, rusting and iron.

: a comparison between actual yield and theoretical yield.

\({percentage~yield} = \frac{actual~yield}{theoretical~yield} \times 100\)

The percentage yield can vary from 100% (no product lost) to 0% (no product made).

Example:

12.4 g of copper(II) carbonate are heated and decomposes. 6 g of copper(II) oxide is formed. Calculate the percentage yield.

CuCO₃ → CuO + CO₂

Calculate the theoretical yield, using the same steps as a reacting mass calculation.

\({Moles~of~CuCO_3} = \frac{mass~(g)}{M_r} = \frac{12.4}{124} = 0.1 mol\)

1 mol CuCO₃ : 1 mol CuO

∴

0.1 mol CuCO₃ forms 0.1 mol of CuO

Mass of CuO = moles x M_r = 0.1 x 80 = 8g

\({percentage~yield} = \frac{actual~yield}{theoretical~yield} \times 100 = \frac{6}{8} \times 100 = 75\%\)

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