Materials Included In Kit

Additional Materials Required

Safety Precautions

Silver oxide is slightly toxic and is a fire risk when in contact with organic material or ammonia. Handle the crucible and its lid only with tongs. Do not touch the crucible with fingers or hands. There is a significant burn hazard associated with handling a hot crucible—remember that a hot crucible looks exactly like a cold one. Always keep your face at arm’s length from the crucible. Wear chemical splash goggles and chemical-resistant gloves and apron. Remind students to wash hands thoroughly with soap and water before leaving the laboratory. Please consult current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The reaction product in the solid waste container may be disposed of according to Flinn Suggested Disposal Method #11. Method #11 also includes an inexpensive route to convert the waste silver back to silver oxide. 

Lab Hints

• The actual experimental work for this lab takes about 45 minutes. For three trials, three 50- minute class periods are required. Before beginning, a review of the safety precautions and a demonstration of the proper techniques for handling a crucible with crucible tongs and for heating a crucible in a Bunsen burner flame is recommended. The Flinn Scientific Laboratory Techniques Guide (see Materials) provides thumbnail illustrations of these and 14 other common laboratory techniques. Enough silver oxide is provided to run about 60 trials.
• The sample data are representative of the accuracy and precision that may be obtained in a typical high school classroom. Some teachers may prefer to have students calculate the empirical formula of silver oxide using classroom results for the number of moles of silver and oxygen in the product. This illustrates to students both the existence and nature of random errors in experimental measurements and also eliminates a problem area for students who have totally outlying results.

Teacher Tips

• An alternative application of this experiment is to use the results to calculate the molar mass of silver, rather than the formula of silver oxide. This can be done by assuming that the reactant must have the formula Ag₂O. (Typically, by the time students reach stoichiometry in the high school curriculum, they already know how to predict the formulas of ionic compounds.) This objective can be achieved using the same procedure as outlined in this experiment. The only difference lies in the nature of the questions and calculations that are posed to the students. See the Supplementary Information in Discussion for appropriate post-lab questions that may be assigned to determine the molar mass of silver.

Further Extensions

AP Standards

This lab fulfills the requirements for the College Board recommended Experiment #1: Determination of the Formula of a Compound.

Answers to Prelab Questions

A piece of iron weighing 85.65 g was burned in air. The mass of the iron oxide produced was 118.37 g.

1. Use the molar mass of iron to convert the mass of iron used to moles.
   
   {12963_Answers_Equation_1}
2. According to the law of conservation of mass, what is the mass of oxygen that reacted with the iron?
   Mass of oxygen = mass of product – mass of iron
   Mass of oxygen = 118.37 g – 85.65 g = 32.72 g
3. Calculate the number of moles of oxygen in the product.
   
   {12963_Answers_Equation_2}
4. Use the ratio between the number of moles of iron and number of moles of oxygen to calculate the empirical formula of iron oxide.
   Mole ratio Fe: 0 = 1.53:2.045 = 1:1.34 = 3.4
   The empirical formula or iron oxide is Fe₃O₄. This form of iron oxide is called magnetite and is actually FeO·Fe₂O_3.
   

Sample Data

Student data will vary.

Answers to Questions

1. Calculate the mass of silver oxide and the mass of the silver metal product. Use the law of conservation of mass to calculate the mass of oxygen that combined with the silver. Enter the answers in the Results Table.

Mass of Ag_xO_y = 0.522 g (Trial 1); 0.507 g (Trial 2); 0.586 g (Trial 3).
Mass of Ag = 0.488 g (Trial 1); 0.468 g (Trial 2); 0.543 g (Trial 3).
Mass of O = 0.034 g (Trial 1); 0.039 g (Trial 2); 0.043 g (Trial 3).

2. What is the percent composition of silver and oxygen in silver oxide? Enter the answers in the Results Table.

Trial 1: %Ag = 0.488 g/0.522 g x 100% = 93.5%

%O = 0.034 g/0.522 g x 100% = 6.5%

Trial 2: %Ag = 0.468 g/0.507 g x 100% = 92.3%

%O = 0.039 g/0.507 g x 100% = 7.7%

Trial 3: %Ag = 0.543 g/0.586 g x 100% = 92.7%

%O = 0.0043 g/0.586 g x 100% = 7.3%

3. Use the molar masses of silver and oxygen to calculate the number of moles of each product. Enter the answers in the Results Table.

Trial 1: Moles Ag = 0.488 g/107.87 g/mole = 0.00452 moles

Moles O = 0.034 g/16.00 g/mole = 0.0021 moles

Trial 2: Moles Ag = 0.468 g/107.87g/mole = 0.00434 moles

Moles O = 0.039 g/16.00 g/mole = 0.0024 moles

Trial 3: Moles Ag = 0.543 g/107.87 g/mole = 0.00503 moles

Moles O = 0.0043 g/16.00 g/mole = 0.0027 moles

4. Calculate the ratio between the number of moles of silver and the number of moles of oxygen in the product. What is the empirical formula of silver oxide? Enter the answers in the Results Table.

Trial 1: Mole ratio Ag:O = 2.2:1
Trial 2: Mole ratio Ag:O = 1.8:1
Trial 3: Mole ratio Ag:O = 1.9:1
The nearest whole number ratio is 2:1. The empirical formula of silver oxide is 2:1 or Ag₂O.

5. Write a balanced chemical equation for the decomposition of silver oxide to form silver metal and oxygen.

2Ag₂O(s) → 4Ag(s) + O₂(g) or Ag₂O(s) → 2Ag(s) + ½O₂(g)

6. The theoretical yield of a product in a chemical reaction is the maximum mass of product that can be obtained, assuming 100% conversion of the reactant(s). Calculate the theoretical yield of silver metal in this experiment. Hint: Calculate the molar mass of silver oxide.

Molar mass (Ag₂O) = 2(107.87) + 16.00 = 231.74 g/mole

According to the balanced chemical equation for its decomposition, one mole of silver oxide produces two moles of silver product. The maximum amount of silver metal is therefore equal to the two times the moles of silver oxide used in each trial run.

7. The percent yield reflects the actual amount of product formed versus the maximum that might have been obtained. Use the following equation to calculate the percent yield of magnesium oxide produced in this experiment.

Trial 1: Percent yield = (0.488g/0.485 g) x 100% = 101%
Trial 2: Percent yield = (0.468 g/0.472 g) x 100% = 99%
Trial 3: Percent yield = (0.543 g/0.546 g) x 100% = 99%

8. Discuss sources of error in this experiment that might account for a percent yield lower or higher than 100%. Be specific!

Two primary errors may account for the observed percent yield being less than 100%.

• • Loss of product as spattering if sample heated too rapidly.
  • Incomplete drying of the crucible prior to measuring its mass at the beginning of the experiment.

One primary error may account for the observed percent yield being greater than 100%.

• • Incomplete decomposition of the silver oxide.

Discussion

Supplementary Information

To adapt this experiment to have students use their results to calculate the molar mass of silver, assign the following Post-Lab Questions for analysis.

Create a Data and Results Table for the following calculation for each trial.

1. Find the mass of silver oxide used in this experiment.
2. Determine the mass of oxygen gas driven off when the silver oxide was heated .
3. Use the known molar mass of oxygen gas to calculate the number of moles of oxygen gas that were produced in this experiment.
4. Silver oxide is an ionic compound. Use the common ion charges for silver and oxygen to predict the formula of silver oxide.
5. Write a balanced chemical equation for the decomposition of silver oxide to produce oxygen gas and silver metal.
6. According to the balanced chemical equation, what is the mole ratio between the number of moles of silver metal and the number of moles of oxygen gas?
7. Assume 100% conversion of silver oxide to silver. Use the stoichiometric mole ratio found in Question 6 to calculate the number of moles of silver produced in this experiment.
8. Divide the mass of silver by the number of moles of silver to calculate the apparent molar mass of silver.
9. Calculate the percent error for the determination of the molar mass of silver and discuss sources of error in this experiment.

Introduction

There is an official database that keeps track of the known chemical compounds that exist in nature or have been synthesized in the lab. The database, called the chemical abstracts database, is updated daily. Currently, more than 20 million different inorganic and organic compounds have been recognized. Twenty million compounds—how is it possible to identify so many different compounds and tell them all apart?

Concepts

• Percent composition

• Empirical formula
• Law of conservation of mass
• Molecular formula
• Percent yield

Background

The composition of a chemical compound—what it is made of—can be described at least three different ways. The percent composition gives the percent by mass of each element in the compound and is the simplest way experimentally to describe the composition of a substance. According to the law of definite proportions, which was first formulated in the early 1800s by Joseph Proust (1754–1826), the elements in a given compound are always present in the same proportion by mass, regardless of the source of the compound or how it is prepared. Calcium carbonate, for example, contains calcium, carbon and oxygen. It is present in eggshells and seashells, chalk and limestone, minerals and pearls. Whether the calcium carbonate comes from a mineral supplement on a drugstore shelf or from seashells on the ocean shore, the mass percentage of the three elements is always the same: 40% calcium, 12% carbon and 48% oxygen.

The percent composition of a compound tells us what elements are present in the compound and their mass ratio. In terms of understanding how elements come together to make a new compound; however, it is more interesting and more informative to know how many atoms of each kind of element combine with one another. Since all the atoms of a given element in a compound have the same average atomic mass, the elements that are present in a fixed mass ratio in a compound must also be present in a fixed number ratio as well. The empirical formula describes the composition o a compound in terms of the simplest whole-number ratio of atoms in a molecule or formula unit of the compound. The empirical formula gives the ratio of atoms in a compound and does not necessarily represent the actual number of atoms in a molecule or formula unit. It is possible, in fact, for different compounds to share the same empirical formula.

The organic compounds acetylene and benzene, for example, have the same empirical formula, CH—one hydrogen atom for every carbon atom. These two compounds, however, have very different properties and different molecular formulas—C₂H₂ and C₆H₆ for acetylene and benzene, respectively. Notice that in both cases, the molecular formula is a simple multiple of the empirical formula. The molecular formula of a compound tells us the actual number of atoms in a single molecule of a compound. In order to find the molecular formula of a compound whose empirical formula is known, the molar mass of the compound must also be determined.

Experiment Overview

In this experiment, the percent composition and empirical formula of silver oxide will be determined. Silver oxide decomposes to silver metal and oxygen when strongly heated. Heating silver oxide causes the oxygen to be driven off, leaving only the silver metal behind. According to the law of conservation of mass, the total mass of the products of a chemical reaction must equal the mass of the reactants. In the case of the decomposition of silver oxide, the following equation must be true:

Mass of silver oxide = Mass of silver metal + Mass of oxygen

If both the initial mass of silver oxide and the final mass of the silver metal are measured, the decrease in mass must correspond to the mass of oxygen that combined with silver. The percent composition and empirical formula of silver oxide can then be calculated, based on combining the ratios of silver and oxygen in the reaction.

Materials

Prelab Questions

A piece of iron weighing 85.65 g was burned in air. The mass of the iron oxide produced was 118.37 g.

1. Use the molar mass of iron to convert the mass of iron used to moles.
2. According to the law of conservation of mass, what is the mass of oxygen that reacted with the iron?
3. Calculate the number of moles of oxygen in the product.
4. Use the ratio between the number of moles of iron and number of moles of oxygen to calculate the empirical formula of iron oxide. Note: Fractions of atoms do not exist in compounds. In the case where the ratio of atoms is a fractional number, such as ½, the ratio should be simplified by multiplying all the atoms by a constant to give whole number ratios for all the atoms (e.g., HO_½  should be H₂O).

Safety Precautions

Silver oxide is slightly toxic and is a fire risk when in contact with organic material or ammonia. Handle the crucible and its lid only with tongs. Do not touch the crucible with fingers or hands. There is a significant burn hazard associated with handling a hot crucible—remember that a hot crucible looks exactly like a cold one. Always keep your face at arm’s length from the crucible. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

1. Set up a Bunsen burner on a ring stand beneath a ring clamp holding a clay pipestem triangle. Place the crucible in the clay triangle (see Figure 1). Do NOT light the Bunsen burner.

2. Adjust the height of the ring clamp so that the bottom of a crucible sitting in the clay triangle is about 1 cm above the burner. This will ensure that the crucible will be in the hottest part of the flame when the Bunsen burner is lit.
3. Light the Bunsen burner and brush the bottom of the crucible with the burner flame for about one minute. Turn off the Bunsen burner and allow the crucible to cool.
4. Using tongs to handle the crucible, measure the mass of a clean, empty crucible and its lid to the nearest 0.001 g. Record the mass in the Data Table.
5. Using proper transfer techniques, add approximately 0.5 grams of silver oxide sample to the crucible. Measure the combined mass of the crucible, crucible lid and silver oxide to the nearest 0.001 g. Record the mass in the Data Table.
6. Place the crucible with its lid on the clay triangle as shown in Figure 2. Light the Bunsen burner again and slowly heat the crucible by brushing the bottom of the crucible with the Bunsen burner flame for 2–3 minutes.

7. Place the burner on the ring stand and gently heat the crucible for an additional 10 minutes.
8. After 10 minutes, adjust the burner to maximize the flame temperature. Heat the crucible with the most intense part of this flame for 10 minutes. Caution: Do not inhale the smoke! Do not lean over the crucible. Keep the crucible at arm’s length at all times.
9. After 10 minutes, turn off the gas source and remove the burner.
10. 10. Using tongs, remove the crucible lid and place it on a wire gauze on the bench top. With the tongs, remove the crucible from the clay triangle and place it on the wire gauze as well (see Figure 3).

11. Allow the crucible and its contents to cool completely on the bench top for at least 10 minutes.
12. Measure the combined mass of the crucible, crucible lid and silver metal product. Record the mass in the Data Table.
13. (Optional) If time permits, dump the contents of the crucible onto a watch glass. Note the appearance and consistency of the product. Is any unreacted silver oxide still present? Record all observations in the data table.
14. Dump the entire contents of the crucible into the waste container provided by the instructor. Carefully clean the crucible and crucible lid.
15. Repeat steps 4 to 14 for trials 2 and 3.

Student Worksheet PDF

12963_Student1.pdf