Slide presentation
Theoretical Yield, Actual Yield, and Percent Yield
General Chemistry • Chemical Reactions
Reaction yield
Theoretical Yield, Actual Yield, and Percent Yield
Theoretical yield is what stoichiometry predicts. Actual yield is what the experiment produces. Percent yield compares the real result to the prediction.
\[ \text{percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\% \]Learning target: calculate theoretical yield from stoichiometry, calculate percent yield, and interpret what the result means in the laboratory.
Theoretical yield
Maximum product amount predicted from the balanced equation.
Actual yield
Product amount collected or measured in the lab.
Percent yield
How efficiently the experiment reached the predicted amount.
Why it matters
Stoichiometry predicts; experiments test
A balanced equation can predict the maximum product mass, but real experiments often lose product or do not finish completely.
If stoichiometry predicts \(56.3\ \mathrm{g}\) H2O but the experiment collects \(44.0\ \mathrm{g}\), percent yield measures how close the experiment came to the ideal result.
Percent yield helps evaluate
- reaction efficiency,
- product loss during transfer or filtering,
- incomplete reactions,
- side reactions,
- measurement or drying errors.
Core concept
Theoretical yield is the ideal maximum
The theoretical yield comes from the limiting reactant and the balanced equation. It assumes all limiting reactant is converted into product with no loss.
The actual yield is measured after the experiment. Percent yield compares actual yield to theoretical yield.
Vocabulary
Yield terms and what they mean
| Term | Meaning | How it is found | Common unit |
|---|---|---|---|
| Theoretical yield | Maximum product predicted by stoichiometry | Calculated from limiting reactant | g or mol |
| Actual yield | Product amount obtained experimentally | Measured in the laboratory | g or mol |
| Percent yield | Actual yield compared with theoretical yield | \(\frac{\text{actual}}{\text{theoretical}} \times 100\%\) | % |
| Limiting reactant | Reactant that controls theoretical yield | Predicts the smaller product amount | mol or g |
| Experimental error | Reason actual yield differs from theoretical yield | Loss, impurity, incomplete reaction, side reaction | varies |
Main relationship
Percent yield compares laboratory result to prediction
The theoretical yield is the denominator because it is the ideal maximum predicted by stoichiometry.
Interactive simulation
Change reactant amounts and lab yield
Reaction setup
Calculated result
This percent yield is realistic: the actual yield is below the theoretical maximum.
Static fallback: if theoretical yield is \(56.3\ \mathrm{g}\) H2O and actual yield is \(44.0\ \mathrm{g}\), percent yield is \(78.2\%\).
Dynamic relationship
Actual yield should be compared to the theoretical maximum
The graph compares the predicted product mass with the mass measured in the lab.
Interpretation: the theoretical yield is the 100% reference. The actual yield usually falls below it because of experimental limitations.
Worked example
Calculate theoretical yield and percent yield
Problem: \(10.0\ \mathrm{g}\) H2 reacts with \(50.0\ \mathrm{g}\) O2. The experiment collects \(44.0\ \mathrm{g}\) H2O. What is the percent yield?
- 1. Convert reactants to moles. \[ n_{\mathrm{H_2}} = \frac{10.0\ \mathrm{g}}{2.016\ \mathrm{g/mol}} = 4.96\ \mathrm{mol\ H_2} \] \[ n_{\mathrm{O_2}} = \frac{50.0\ \mathrm{g}}{32.00\ \mathrm{g/mol}} = 1.56\ \mathrm{mol\ O_2} \]
- 2. Find the theoretical yield from the limiting reactant. O2 is limiting because it gives the smaller product amount: \[ 1.56\ \mathrm{mol\ O_2} \times \frac{2\ \mathrm{mol\ H_2O}}{1\ \mathrm{mol\ O_2}} = 3.12\ \mathrm{mol\ H_2O} \] \[ 3.12\ \mathrm{mol\ H_2O} \times 18.015\ \mathrm{g/mol} = 56.3\ \mathrm{g\ H_2O} \]
- 3. Calculate percent yield. \[ \text{percent yield} = \frac{44.0\ \mathrm{g}}{56.3\ \mathrm{g}} \times 100\% = 78.2\% \]
- Final answer: The percent yield is \(78.2\%\).
Common mistake
Do not put actual yield in the denominator
Incorrect setup
A student writes \(\frac{\text{theoretical yield}}{\text{actual yield}} \times 100\%\).
This reverses the comparison and can give a value that does not represent experimental efficiency.
Correct setup
Percent yield asks: “What fraction of the predicted product did the experiment actually produce?”
\[ \text{percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\% \]
Key idea: theoretical yield is the ideal maximum and belongs in the denominator.
Practice check
Calculate percent yield
A reaction has a theoretical yield of \(12.5\ \mathrm{g}\) CaCO3. The actual yield collected in the lab is \(10.8\ \mathrm{g}\) CaCO3. What is the percent yield?
Show answer and reasoning
Use the percent yield formula:
\[ \text{percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\% \]
Substitute the values:
\[ \text{percent yield} = \frac{10.8\ \mathrm{g}}{12.5\ \mathrm{g}} \times 100\% = 86.4\% \]
Answer: the percent yield is \(86.4\%\).
Apply the topic
How to interpret percent yield
Common and realistic
Product may be lost, reaction may not finish, or side reactions may occur.
High efficiency
The actual yield is close to the theoretical prediction.
Usually suspicious
Product may be wet, contaminated, or measured incorrectly.
Percent yield depends on theory
A good percent yield calculation requires a correct theoretical yield first.
In future problems, always identify whether the question gives theoretical yield directly or requires you to calculate it from stoichiometry.
Summary
What to remember
Theoretical yield is predicted
It comes from stoichiometry and the limiting reactant.
Actual yield is measured
It is the amount of product collected experimentally.
Percent yield compares them
Use actual yield divided by theoretical yield, then multiply by \(100\%\).
Interpret the result
Percent yield links calculations to reaction efficiency and experimental error.