Percent Yield - Online Tutor, Practice Problems & Exam Prep

In chemistry, the efficiency of a chemical reaction is quantified by the percent yield, which is calculated using the formula:

(actual yield/theoretical yield)×100 The actual yield represents the pure product obtained from the laboratory experiment, while the theoretical yield is the maximum product amount predicted by stoichiometric calculations. Percent yield values are categorized as excellent (≥90%), very good (≥80%), good (≥70%), or poor (<40%). It's important to note that due to practical limitations, no chemical reaction achieves 100% efficiency, meaning the actual yield is invariably lower than the theoretical yield. Understanding and calculating percent yield is crucial for evaluating the success of a reaction and for practical applications where maximizing product yield is essential.

Percent Yield

Now, stoichiometry is a way of us determining, theoretically on paper, how much product we could make if we are given some given amount of starting material. But when we do the experiment in real life and obtain a certain amount of product, it's the percent yield that determines how successful we were. Now, percent yield determines how successful the scientist was in creating their desired product. The higher the percent yield than the higher the efficiency of a chemical reaction.

So just like you want the highest possible percentages on an exam, we want the highest possible percentages in percent yield. That'll show us being successful in creating our product. Now we're going to say in terms of percent yield values, we can have excellent, very good, good and poor. Now if you have a percent yield that is equal to or greater than 90%, you could call that an excellent yield. If you have a amount that's equal to or greater than 80%, then that would be very good. If you were good, that means you would be equal to or greater than 70% and then you'd have a poor yield if your percent yield was less than 40%.

Now with percent yield comes the percent yield formula, and the percent yield formula equals actual yield over theoretical yield times 100. And remember, we have our purple box here. That means that this is a formula you need to commit to memory because oftentimes it's not given on your formula sheet. Now we know that the percent yield measures how successful we are in terms of carrying this reaction out in real life. We know that theoretical yield is what we do as calculations on paper when they give us multiple amounts of given and we figure out how much product we're making.

Actual yield though, actual yield is the amount of pure product actually created when the experiment is done in the laboratory. Oftentimes what you do on paper, you'll see that when you do the experiment in real life, things don't exactly match U. We're going to say here that the units used in the formula are based on the units of the actual yield. So let's say that our actual yield is in grams of our product, but our theoretical yield that we calculated is in moles. You'd have to change those theoretical yield units to match the actual yield units. So you have to change moles to grams.

We're going to say here that no chemical reaction is 100% efficient, so you'll never ever get a percent yield that is 100%. There's always going to be something that happens. You're going to spill some of your compound. You're going to lose some randomly outside forces will play a part. So you'll never get to 100% yield. Because of this, that means that your actual yield is always less than your theoretical yield. So just remember, we have percent yield, we have actual yield, we have theoretical yield. Together, they give us a good insight into how efficient our chemical reaction is.

PercentYield = ActualYield TheoreticalYield × 100

Percent Yield Example 1

Here are example questions says consider the following balanced chemical reaction. So we have two moles of C₆H₆ reacting with 15 moles of oxygen gas to produce 12 moles of carbon dioxide gas plus six moles of water as a liquid. It says if a 2.6g sample of C₆H₆ reacted with excess oxygen to produce 1.25 grams of water, what is the percent yield of water?

All right, they want us to determine the percent yield. We know percent yield equals actual over theoretical times 100. If we read back on the question, they're telling me that I produced 1.25 grams of water. How can I determine if I produced something or made something made product? The only way I would know for sure is if I did the the reaction in real life and obtain that amount. That 1.25g represents our actual yield. So really what we have to do is calculate our theoretical yield.

And remember, our theoretical yield can be determined by using stoichiometry O. If we look here, it says step one, map out the portion of the stoichiometric chart you will use. Well, access, remember means we ignore. They're telling you that I have 2.6g of this reactant. That's the amount they gave to me. So that represents my grams of given. So the map that I'm going to have to workout is grams of given. Convert that into moles of given. Then I'm going to have to convert those moles of given to moles of unknown.

What's our unknown? Well, to figure out the units that I need. Remember, your theoretical yield must have the same units as your actual yield. Our actual yield is in grams per of water, so we have to go all the way to grams of water. So we're going to have here moles of unknown, which is water and then finally end here in grams of unknown. Alright. So that's the stoichiometric chart, at least the portion of it that we need to utilize to get our theoretical yield all right.

So we're going to take the given quantity and change it into moles of given. We have 2.6g of benzene, we have 6 carbons and six hydrogens. When you look on the periodic table and take their atomic masses, add up the six carbons, add up the six hydrogens, you'll have a combined mass of 78.108g per one mole of C₆H₆. Next we're going to say here, step three, do a mold to mole comparison to convert moles of given into moles of unknown. To do that, remember we look at the coefficients of the balanced equation. We're going to go for moles of C₆H₆ to moles of water and in our balance equation it is a two to six ratio.

Then it says if necessary, convert the moles of unknown into desired units into the final desired units. Since our actual yield is using grams of water, our theoretical yield needs to go to grams of water, one mole of water. Taking into account the two hydrogens in the one oxygen, it has a combined mass of 18.016g. Then if we multiply everything out into Viva what's on the bottom, we get 1.799 grams of water. This year represents our theoretical yield.

So plug that in to our answer and then step five. We're already doing it. We're plugging in the actual yield and the theoretical yield, and that'll help us define our percent yield. SO12 point 1.25 / 1.799 * * 100 gives me 69.48% as my percent yield for this particular question. So just remember, parts of this are pretty familiar. They're incorporating things we've learned about stoichiometry and limiting reagents, and now just tacking on actual yield. With that information, we're able to find our percent yield for this particular question.

For the following chemical reaction 53.1 g HBrO₄ is mixed with 25.8 g Li₂CO₃. Determine the percent yield if 7.17 g CO₂ is produced.

2 HBrO₄ (aq)  +  Li₂CO₃ (aq)  → H₂O (l)  +  CO₂ (g)  +  2 LiBrO₄ (aq)

• A. 73.1%
• B. 88.9%
• C. 90.7%

Ammonia, NH₃, can be created from the combining of H₂ and N₂ molecules. How many grams of ammonia are isolated when 10.0 g H₂ reacts with excess N₂ with a chemical reaction that has a 79.3% yield?

3 H₂  +  N₂  →  2 NH₃

• A. 34.4 g
• B. 53.0 g
• C. 44.7 g
• D. 14.5 g

The reduction of iron (III) oxide creates the following reaction: 

Fe₂O₃ (s) + 3 H₂ (g) → 2 Fe (s) + 3 H₂O (g) 

If the above reaction only went to 75% completion, how many moles of Fe₂O₃ were require to produce 0.850 moles of Fe? 

• A. 0.963 mol
• B. 2.27 mol
• C. 0.425 mol
• D. 0.567 mol
• E. 1.70 mol

Here’s what students ask on this topic:

How do you calculate percent yield?

To calculate percent yield, you'll first need to know two things: the theoretical yield and the actual yield. The theoretical yield is the maximum amount of product that could be formed from a reaction, based on stoichiometry calculations from the balanced chemical equation. The actual yield is the amount of product actually obtained from the reaction.

Here's the formula for percent yield:

$\text{Percent Yield}=\left(\frac{\text{Actual Yield}}{\text{Theoretical Yield}}\right)\times 100\%$

Follow these steps:

1. Start by determining the theoretical yield. This often involves using mole ratios from the balanced equation to find out how many moles of product you should get if everything reacts perfectly.
2. Measure the actual yield from your experiment. This is typically done by weighing the product you've obtained.
3. Plug these values into the formula above. Divide the actual yield by the theoretical yield, multiply by 100, and you'll get the percent yield.

This value tells you the efficiency of your reaction. A percent yield of 100% means you obtained all the product expected, while a lower percent yield indicates that the reaction was less efficient and some reactants did not convert to product as expected.

What is percent yield?

Percent yield is a term used in chemistry to describe the efficiency of a chemical reaction. It's a way to measure what percentage of the theoretical yield was actually obtained from a reaction. The theoretical yield is the maximum amount of product that could be formed from the given amounts of reactants, based on the stoichiometry of the balanced chemical equation. This is a calculation based on the assumption that every molecule of reactants converts to products, which doesn't always happen in practice due to side reactions, incomplete reactions, or loss of product during the process.

To calculate percent yield, you take the actual yield, which is the amount of product you actually obtained from the reaction, and divide it by the theoretical yield. Then, you multiply the result by 100 to get a percentage. The formula looks like this:

$\text{Percent Yield}=\frac{\text{Actual Yield}}{\text{Theoretical Yield}}\times 100\%$

A percent yield can be over 100% if there are impurities or measurement errors, but typically, it's less than 100% due to the practical limitations mentioned earlier. Understanding percent yield is crucial for evaluating the effectiveness of laboratory reactions and for scaling up chemical processes in an industrial setting.

What does percent yield tell you?

Percent yield is a measure used in chemistry to express the efficiency of a chemical reaction. It compares the actual yield—the amount of product actually obtained from a reaction—to the theoretical yield, which is the maximum amount of product that could be formed based on stoichiometric calculations from the balanced chemical equation. The percent yield is calculated by taking the actual yield, dividing it by the theoretical yield, and then multiplying by 100 to get a percentage.

The formula for percent yield is:

$\text{Percent Yield}=\frac{\text{Actual Yield}}{\text{Theoretical Yield}}\times 100\%$

A percent yield of 100% indicates a perfectly efficient reaction where the actual yield matches the theoretical yield. In reality, yields are often less than 100% due to factors such as incomplete reactions, side reactions, loss of product during purification, measurement errors, or limitations in reaction conditions. Percent yield gives chemists insight into the reaction's effectiveness and can help in optimizing processes, controlling costs, and assessing the feasibility of industrial-scale production.