# Actual Yield Calculator

Created by Juhi Raj, PhD candidate
Reviewed by Hanna Pamuła, PhD candidate and Jack Bowater
Last updated: Sep 27, 2022

This actual yield calculator will answer your questions about how to calculate the actual yield and help you find the true yield of a chemical reaction.
Because reality is much different than theory, no chemical reaction is ever carried out flawlessly. Because of this, the actual yield definition states that it should always be lower than the theoretical yield.
Before commencing any laboratory work, figure out what the theoretical yield is so you know how much of your product to expect from a given amount of starting material. The efficiency of your reaction will help you compute the actual yield of a chemical process using the actual yield formula - or you can simply use our amazing tool!

Remember, that we also have a theoretical yield calculator and a percent yield calculator available for your chemistry calculations!

## What is actual yield? Actual yield definition

The quantity of a product received from a chemical reaction is known as the actual yield.
The theoretical yield, on the other hand, is the quantity of product that could be obtained if all of the reactants were converted to product perfectly, i.e., there were no by-products.
The limiting reactant is used to calculate theoretical yield.
By determining the percent yield, also known as the efficiency of the chemical reaction, you can calculate the actual yield using our actual yield calculator.

## How to calculate the actual yield? Actual yield formula

Now that we know the actual yield definition, let's proceed to the actual yield formula for a chemical process:

$Y_{\text{a}} = \frac{Y_{\text{p}}}{100}\cdot Y_{\text{t}}$

where:

• $Y_{\text{a}}$ — Actual yield;
• $Y_{\text{p}}$ — Percent yield ($\%$); and
• $Y_{\text{t}}$ — Theoretical yield.

## Difference between theoretical yield and actual yield

The actual yield is usually lower than the theoretical yield because few reactions proceed to absolute completion (i.e., they aren't 100% efficient) or because not all of the product in a reaction is recovered.
For example,

• If you're recovering a precipitate, you can lose some of it if it doesn't completely crash out of the solution;
• Some products may remain on the filter paper or make their way through the mesh and wash away if you filter the solution through filter paper;
• Even if the substance is insoluble in a particular solvent, if you rinse it, a little bit of it may be lost due to dissolving in the solvent; and
• Sometimes your product may undergo unwanted reactions after you have formed it due to the presence of high temperatures or other chemicals.

But, it is also possible that the actual yield exceeds the theoretical yield:

• This happens most frequently when the solvent is still present in the product (incomplete drying), when the product is weighed incorrectly, or when an unaccounted ingredient in the reaction acts as a catalyst or causes the creation of by-products; and
• Another reason for the increased yield is that the result is impure because another ingredient than the solvent is present.

## Examples of how to find actual yield

Let's move on to real-world applications now that we've learned about the distinctions between actual and theoretical yields.

Example 1: If the percent yield of a reaction was $45\%$ with a theoretical yield of $Y_{\text{t}} = 4\ \text{g}$, what is the actual yield?

You can solve this problem by using our actual yield calculator:

1. Input the percent yield: $Y_{\text{p}} = 45\%$.

2. Input the theoretical yield: $Y_{\text{t}} = 4\ \text{g}$.

3. The actual yield is $Y_{\text{a}}=1.8\ \text{g}$.

Example 2: When making a new substance from other substances, chemists say that they have synthesized a new compound.
Reactants are converted to products, and the process is symbolized by a chemical equation. For example, the decomposition of magnesium carbonate to form magnesium oxide has an experimental percent yield of $79\%$. The theoretical yield is known to be $Y_{\text{t}}=19\ \text{g}$. What is the actual yield of magnesium oxide?

$\text{MgCO}_3\rightarrow\text{MgO}+\text{CO}_2$

So, how to find the actual yield? The calculation is simple if you know the percent and theoretical yields.
All you need to do is plug the values into the calculator again:

1. Input the percent yield: $Y_{\text{p}} = 79\%$.

2. Enter the theoretical yield: $Y_{\text{t}} = 19\ \text{g}$.

3. The actual yield is $Y_{\text{a}}=15\ \text{g}$.

Usually, you have to calculate the theoretical yield based on the balanced equation. In this equation, the reactant and the product have a $1:1$ mole ratio, so, if you know the amount of reactant, you know the theoretical yield is the same value in moles (not grams!). You take the number of grams of reactant you have, convert it to moles (maybe with our grams to moles calculator), and then use this number of moles to find out how many grams of product to expect.

You can find the yield of the reaction at every moment using the concentrations you can calculate with our reaction quotient calculator!

🙋 Make sure the theoretical yield and actual yield are expressed in the same unit: if you need to know the conversion, visit our weight converter.

Juhi Raj, PhD candidate
Theoretical yield
g
Percent yield
%
Actual yield
g
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