2.12. Percentage Yield

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1 Percentage Yield The quantities of reactants and products you calculate on the basis of the mole ratios in a balanced equation are called theoretical quantities. Theoretical quantities are the quantities that should be used or produced in a chemical reaction. They are not necessarily the quantities that are used or produced when you mix the reactants and collect the products. The quantity of product produced in a chemical reaction is called the yield. When the reaction is carried out in a laboratory, the quantity of product that is obtained and measured at the end of the reaction is called the. The quantity of product (the yield) can also be calculated using the balanced equation. The calculated quantity is called the. Theoretically, each and every entity that is used as a reactant in a chemical reaction should be accounted for in the products. Often, however, the actual yield in a chemical reaction turns out to be less than the. There are several reasons why not all the reactants end up in the products. The most common reason is related to experimental procedures. For example, loss of product may occur when transferring solutions or filtering precipitates, or from splattering during heating. This kind of loss can be reduced by improving technical skills, improving the equipment used, or reducing the number of steps in the experimental design. Low yield can also be due to impurities in the reagents used. Chemicals come in a wide variety of grades, or purities (Figure 1). Some low-purity or technical grades may only be 80% to 90% pure. If impurities are not accounted for in the amount of reactant used, the will be less than the. Impurities may also result from other processes. For example, a metal, such as magnesium, readily reacts with air to form a layer of magnesium oxide on the surface. This impurity is included in the mass of the reactant but does not proceed to form the products collected, thus causing the to differ from the. Another cause of low yield is the occurrence of side reactions that form products other than the desired products. As already mentioned, magnesium ribbon reacts with the oxygen in air to form magnesium oxide. Air contains other gases, however, such as nitrogen. A side reaction may occur, in which some of the magnesium reacts with nitrogen to form magnesium nitride. If the product that is collected is presumed to be only magnesium oxide but also contains magnesium nitride, the will be different from the. To correct this problem, the reaction may be carried out in pure oxygen instead of air. Low yield may also occur because most reactions are reversible. In reversible reactions, products may react with one another to form the reactants. Environmental conditions determine the extent to which reactions proceed in one direction or the other. When products accumulate in closed containers, they have a tendency to react with each other to regenerate the original reactant molecules. As a result, the amount of product that is 2.12 yield the quantity of product produced in a chemical reaction the quantity of product that is actually produced in a chemical reaction the quantity of product calculated from a balanced equation LEARNING TIP Different Yields? Actual yield is sometimes referred to as experimental yield. Figure 1 Chemicals come in a wide variety of grades, or purities. The purity of a chemical can significantly affect experimental results in a chemical reaction. Quantities in Chemistry 155

2 percentage yield the ratio, expressed as a percentage, of the actual or experimental quantity of product obtained () to the maximum possible quantity of product () derived from a stoichiometric calculation collected may be less than the amount predicted by simple stoichiometry calculations. To minimize these losses, the conditions for the reaction may need to be changed to allow the forward reaction to go to completion. For example, the reaction may be carried out in an open container. To compare the and the, the percentage yield is calculated. The percentage yield is obtained by dividing the by the : percentage yield 100% If, for example, the in a certain reaction is 10 kg and the is 9 kg, the percentage yield can be calculated as follows: 9 kg percentage yield 100% 10 kg percentage yield 90% The percentage yield is 90%. SAMPLE problem Calculating the Percentage Yield Iron is produced from its ore, hematite, Fe 2, by heating hematite with carbon monoxide in a blast furnace. The IUPAC name for hematite is iron(iii) oxide. If 635 kg of iron is obtained from 1150 kg of hematite, what is the percentage yield of iron? The equation for the reaction is Fe 2 3 CO (g) 2Fe (s) 3 CO 2(g) To calculate the percentage yield, you need to know the and the. In this problem, you are given the produced and the mass of hematite used. Therefore, you need to calculate the using the balanced equation provided. Use the steps you learned in section 2.9 (Stoichiometry) to calculate the mass of iron that should be produced. Hematite is the given substance. Step 1: Write Unbalanced Equation Since a balanced equation is provided, proceed to step 2. Step 2: Balance Equation, List Given Values and Molar Masses Balanced equation Fe 2 3 CO (g) 2Fe (s) 3 CO 2(g) Given mass (kg) Given mass (g) Molar mass (g/mol) Unit 2

3 Section 2.12 Step 3: Convert Mass of Given Substance to Amount of Given Substance Use the molar mass of the given substance: n Fe g Fe 2 1 mol Fe g Fe2 n Fe mol Fe 2 Step 4: Convert Amount of Given Substance to Amount of Required Substance Use a mole ratio that allows you to cancel mol Fe 2. From the balanced equation, the mole ratio of hematite and iron is 1 mol Fe 2 : 2 mol Fe. 2 mol Fe n Fe 7201 mol Fe 2 1 mol Fe2 n Fe mol Fe Step 5: Convert Amount of Required Substance to Mass of Required Substance Use the molar mass of iron to calculate the mass of iron g Fe mol Fe 1 mol Fe g Fe Convert grams to kilograms kg Fe g Fe 1000 g Fe kg Fe 804 kg Fe The of iron is 804 kg. Step 6: Calculate Percentage Yield percentage yield 100% 635 kg 100% 804 kg percentage yield 79% The percentage yield of iron is 79%. Example The most common ore of arsenic, FeSAs (s), can be heated to produce arsenic, As (s) : FeSAs (s) FeS (s) As (s) When 250 kg of this ore was processed industrially, 95.3 kg of arsenic was obtained. Calculate the percentage yield of arsenic. Quantities in Chemistry 157

4 (a) Solution Balanced equation FeSAs (s) FeS (s) As (s) Given mass (kg) Given mass (g) Molar mass (g/mol) (b) n FeSAs g FeSAs n FeSAs 1535 mol FeSAs 1 mol FeSAs g FeSAs The mole ratio from the balanced equation is 1 mol FeSAs : 1 mol As. n As 1535 mol FeSAs n As 1535 mol As 1 mol As 1 mol FeSAs (c) g As m As 1535 mol As 1 mol As g As kg As g As 1000 g As m As kg percentage yield 100% 95.3 kg 100% kg percentage yield 82.9% The percentage yield of arsenic was 82.9%. Figure 2 (a) Cochineal insects live in prickly pear cacti in the high desert plains of the Peruvian Andes. (b) (c) A vivid red dye is made of the dried and crushed bodies of female cochineal insects. Answers % 2. (b) 81.0% % Practice Understanding Concepts 1. Methyl salicylate, C 8 (l) the chemical that is responsible for wintergreen flavouring. It can be prepared by heating salicylic acid,, with methanol, CH 3 OH (l) : CH 3 OH (l) C 8 (l) H 2 O (l) If 2.00 g of salicylic acid reacts with excess methanol, and the yield of wintergreen is 1.65 g, what is the percentage yield? 2. Aluminum metal reacts with liquid bromine to produce solid aluminum bromide as the only product. (a) Write a balanced chemical equation for this reaction. (b) When 53.7 g of bromine reacts with excess aluminum, 48.4 g of aluminum bromide is produced. Calculate the percentage yield for this reaction. 3. Zinc reacts with hydrochloric acid, as shown in the following balanced chemical equation: Zn (s) 2HCl (aq) ZnCl 2(aq) H 2(g) 158 Unit 2

5 Section 2.12 Calculate the percentage yield if g of zinc chloride, ZnCl 2(aq) produced when g of hydrochloric acid reacts with excess zinc. 4. A student produces acetylsalicylic acid (Aspirin), C 9 O 4(s), and acetic acid from the reaction of salicylic acid,, with acetic anhydride, C 4 (aq). The balanced chemical equation is Answers 4. (a) g (b) 68.14% C 4 (aq) C 9 O 4(s) HC 2 H 3 O 2(aq) (a) What is the of Aspirin, if g of salicylic acid reacts with excess acetic anhydride? (b) What is the percentage yield of Aspirin, if g of Aspirin is produced? Section 2.12 Questions Understanding Concepts 1. Distinguish between the terms and. 2. Can the ever be greater than the? Explain. 3. In an experiment, 5.00 g of silver nitrate, AgN added to a solution that contains an excess of sodium bromide, NaBr (s), and 5.03 g of silver bromide, AgBr (s) produced. (a) Write a balanced equation for the reaction. (b) What is the of silver bromide? (c) What is the of silver bromide in the experiment? (d) What is the percentage yield in the experiment? 4. In an experiment, 16.1 g of iron sulfide, FeS (s) added to excess oxygen, and 14.1 g of iron(iii) oxide, Fe 2 produced. The balanced equation for the reaction is 4 FeS (s) 7O 2(g) 2Fe 2 4 SO 2(g) Calculate the percentage yield of iron(iii) oxide in the experiment. Applying Inquiry Skills 5. In an experiment to recover a precipitate that was formed in a chemical reaction, a chemistry student followed the procedure below: 1. The mass of a reactant was determined using weighing paper on an electronic balance. 2. The reactant was transferred to a large beaker. 3. A second reactant, an aqueous solution used in excess, was measured using a graduated cylinder and added to the beaker. 4. The mixture was stirred, placed in an evaporating dish, and heated to dryness on a laboratory hot plate. 5. The precipitate was transferred from the evaporating dish to the weighing paper, and the mass was determined. Suggest ways in which this procedure could be modified to improve the percentage yield. Making Connections 6. When you consume a beverage or candy that was artificially coloured with red food dye, you may be ingesting a chemical that was produced by a tiny red insect from Peru. Several synthetic red dyes have been found to be carcinogenic (cancer-causing), but a vivid red dye called carmine has been approved for use in foods, drugs, and cosmetics. Carmine is made from cochineal insects (Figure 2). Two Canadian chemists developed the process that is used to extract the red dye from the insects. The process has been improved to increase the purity and the yield of the product. Research answers to the following questions, and summarize your findings in a onepage report. (a) What is the typical percentage yield of carmine in the extraction process? How is carmine extracted? (b) What effect has the industrial production of carmine had on the people of Peru? GO Quantities in Chemistry 159