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Chapter 9: Problem 4

How is percentage yield calculated?

Short Answer

Expert verified
The percentage yield can be calculated by dividing the actual yield of the experiment or process by the theoretical maximum yield, and then multiplying the result by 100.

Step by step solution

01

Identify the Actual and Theoretical Yield

In any given problem or scenario, the Actual Yield is typically an amount that would be given, it represents the quantity that is actually produced. The Theoretical Yield is the maximum amount of product that could be formed from the reactants, it is calculated based on the balanced chemical reaction.
02

Apply the Percentage Yield Formula

The formula for calculating the percentage yield is: Percentage Yield = (Actual Yield/Theoretical Yield) x 100. This formula is generally used in experiments and industrial processes to measure the efficiency of the process.
03

Calculate the Percentage Yield

Substitute the values of the Actual Yield and the Theoretical Yield into the formula and perform the division. Then multiply the result by 100 to convert it into a percentage.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Actual Yield
The actual yield is the real amount of product obtained from a chemical reaction. It is determined by performing the reaction in a lab setting or during an industrial process. Unlike theoretical yield, which is calculated, actual yield is measured.
Real-world factors like incomplete reactions, equipment efficiency, and measurement errors can affect the actual yield. Hence, it is typically less than the theoretical yield.
  • Actual yield can be expressed in grams, liters, or moles.
  • It provides insight into how successful a chemical process is.
Understanding actual yield helps in identifying the limitations of a reaction and improving the process's efficiency.
Theoretical Yield
The theoretical yield is the maximum amount of product that can be formed in a chemical reaction, assuming complete conversion of reactants to products. This value is based on the balanced chemical equation and stoichiometry principles.
To calculate the theoretical yield, follow these steps:
  • Write down the balanced chemical equation.
  • Determine the moles of limiting reactant.
  • Use stoichiometry to convert moles of reactant to moles of product.
  • Convert moles of product to mass, if necessary, using the molar mass.
Theoretical yield is usually greater than the actual yield because it assumes ideal conditions without any losses or inefficiencies. Understanding this concept will help in evaluating the efficiency of a chemical process.
Chemical Reaction
A chemical reaction is a process that leads to the transformation of one set of chemical substances to another. It involves breaking and forming chemical bonds, resulting in the production of new substances.
Key features of chemical reactions include:
  • Rearrangement of atoms.
  • Conservation of mass, following the law of conservation of mass.
  • Energy changes, as reactions absorb or release energy.
Chemical reactions are represented by chemical equations, which must be balanced to obey the law of conservation of mass. Understanding chemical reactions is fundamental in predicting the amount of products and in calculating both the theoretical and actual yields.

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Most popular questions from this chapter

Reacting 991 mol of \(\mathrm{SiO}_{2}\) with excess carbon yields 30.0 \(\mathrm{kg}\) of \(\mathrm{SiC.}\) What is the percentage yield? $$\mathrm{SiO}_{2}+3 \mathrm{C} \rightarrow \mathrm{SiC}+2 \mathrm{CO}$$

How many grams of \(\mathrm{NaNO}_{2}\) form when 256 \(\mathrm{g} \mathrm{NaNO}_{3}\) react? The yield is 91\(\% .\) $$2 \mathrm{NaNO}_{3}(s) \longrightarrow 2 \mathrm{NaNO}_{2}(s)+\mathrm{O}_{2}(g)$$

Write a balanced equation for the combus- tion of octane, \(\mathrm{C}_{8} \mathrm{H}_{18},\) with oxygen to obtain carbon dioxide and water. What is the mole ratio of oxygen to octane?

Nitrogen monoxide, NO, reacts with oxygen to form nitrogen dioxide. Then the nitrogen dioxide reacts with oxygen to form nitrogen monoxide and ozone. Write the balanced equations. What is the theoretical yield in grams of ozone from 4.55 g of nitrogen monoxide with excess \(\mathrm{O}_{2} ?\) (Hint: First calcu- late the theoretical yield for NO \(_{2}\) , then use that value to calculate the yield for ozone.)

How many liters of \(\mathrm{O}_{2},\) density \(1.43 \mathrm{g} / \mathrm{L},\) are needed for the complete combustion of \(1.00 \mathrm{L} \mathrm{C}_{8} \mathrm{H}_{18},\) density 0.700 \(\mathrm{g} / \mathrm{mL} ?\)
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