Actual yield is one of the types of yield in a chemical reaction, along with theoretical yield and percent yield. Here is the actual yield definition, how to find actual yield, and a look at why it’s always less than theoretical yield in an experiment.
Actual Yield Definition
Actual yield is the amount of product you experimentally obtain from a chemical reaction. In contrast, theoretical yield is the amount of product you obtain if all of the reactant converts to product. Actual yield is an empirical value that you measure in the lab, while theoretical yield is a calculated value.
How to Find Actual Yield
Usually, you find actual yield by weighing product using a scale:
- Weigh the container.
- Weigh the dry product in the container.
- Subtract the mass of the container from the total mass to get the mass of the product.
However, sometimes the product is measured indirectly in the unpurified reaction mixture. Measurements are taken via gas chromatography (GC), high-performance liquid chromatography (HPLC), nuclear resonance spectroscopy (NMR), or another analytical technique.
How to Calculate Actual Yield from Percent Yield
Another way of finding actual yield is from percent yield and theoretical yield.
percent yield = actual yield/theoretical yield x 100
actual yield = (percent yield x theoretical yield)/100
Isolated Yield
Many labs report isolated yield rather than actual yield. Isolated yield is the yield of product measured after it has been purified to a certain level (usually >95% spectroscopic purity). Because some product gets lost during purification, isolated yield tends to be lower than actual yield.
Reasons Why Actual Yield Is Less Than Theoretical Yield
Actual yield is lower than theoretical yield because most reactions aren’t 100% efficient and because it’s impossible to recover all of the product from a reaction. For example:
- Product remains on filter paper or passes through it.
- A tiny amount of product dissolves in a washing solvent, even if it’s insoluble in that solvent.
- Product that is a precipitate incompletely falls out of solution.
- Product evaporates.
Although less common, actual yield may be more than theoretical yield. Incompletely drying is the most common reason for this. Another reason is because an impurity is included in the product weight. Rarely, actual yield is higher than theoretical yield if another chemical reaction in the experiment also forms the same product.
Yield Inflation
In a 2010 Synlett article, Wernerova and Hudlický reported that the purification steps leading to isolated yield result in a loss of around 2% of product. Given the inherent loss, they concluded isolated yield rarely exceeds 94%. Yet, publications increasingly report higher and higher yields. This phenomenon is called yield inflation. There are multiple explanations for yield inflation.
- Improved techniques lead to higher yields.
- Small-scale reactions are more susceptible to slight measurement differences.
- Researchers artificially inflate yields to appear better in publication.
Assuming yield inflation is, in fact, a real phenomenon, the explanation is left for the reader to decide.
References
- McNaught, A. D.; Wilkinson, A., eds. (1997). Compendium of Chemical Terminology (the “Gold Book”) (2nd ed.). Oxford: Blackwell Scientific Publications. doi:10.1351/goldbook ISBN 0-9678550-9-8.
- Petrucci, R.H., Harwood, W.S.; Herring, F.G. (2002) General Chemistry (8th ed.). Prentice Hall. ISBN 0130143294.
- Wernerova, Martina; Hudlicky, Tomas (November 2010). “On the Practical Limits of Determining Isolated Product Yields and Ratios of Stereoisomers: Reflections, Analysis, and Redemption”. Synlett. 2010 (18): 2701–2707. doi:10.1055/s-0030-1259018
- Vogel, A. I.; Tatchell, A. R.; Furnis, B. S.; Hannaford, A. J.; Smith, P. W. G. (1996) Vogel’s Textbook of Practical Organic Chemistry (5th ed.). Pearson. ISBN 978-0582462366.
- Whitten, K.W., Gailey, K.D; Davis, R.E. (1992) General Chemistry (4th ed.). Saunders College Publishing. ISBN 0030723736.