In the Laboratory
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The Empirical Formula of Silver Sulfide An Experiment for Introductory Chemistry
Carlos Alexander Trujillo Departamento de Química, Universidad Nacional de Colombia, Ciudad Universitaria, Cra. 30 45-03 Bogotá, Colombia;
[email protected] Early in introductory chemistry, students learn that compounds can be given empirical formulas. The term empirical formula is in common use although “relative compositional formula” has been proposed as a substitute phrasing (1). Experimental determination of an empirical formula for a compound serves to reiterate what is probably the most important experimentally supported notion in chemistry: matter has a particulate nature. Another key idea is illustrated in the determination of an empirical formula: the question “how much” is equivalent to the question “how many”. Empirical formula calculations employ yet another basic concept of chemistry: the law of conservation of matter (2). A cursory search of chemical literature uncovers several laboratory exercises related to an experimental determination of an empirical formula (3–14). This article proposes a laboratory exercise leading to the assignment of an empirical formula to silver sulfide. Students decompose a silver sulfide sample by heating it in air. Sulfur is eliminated from the compound as sulfur dioxide, and the final product is metallic silver. The initial sample mass and the mass of silver recovered provide the experimental data needed to assign an empirical formula for silver sulfide. Students are told that the sample is a binary compound of silver and sulfur and that the mass change is the result of removing sulfur from the compound. In support of the description given, one can demonstrate the formation of silver sulfide from its elements. Small pellets of silver obtained after the completion of the exercise can be mixed with sulfur and gently warmed. The substance formed has a similar color and appearance as that of the original sample. One may also note that the tarnish that accumulates on silver objects is largely silver sulfide (15).
Procedure Students can rapidly perform the heating process to acquire data that give consistent results. The crucible is initially heated to ∼900 ⬚C, cooled, and stored in a desiccator. A high temperature burner or a muffle furnace is needed for this operation. Even the cool crucible should be handled with tongs to avoid changing its mass during manipulation. The mass of the crucible with and without the mass of the compound (between 200 mg and 500 mg) are recorded. Students gently heat the crucible and its contents for ten minutes. During this initial heating, students may observe the transformation of a dark brown compound to a white solid that has a metallic luster. During the initial heating, a mixture of metallic silver and silver sulfate is formed. The temperature required for the decomposition of the silver sulfate requires that the crucible be heated with the hottest flame of a high temperature burner or by placing the crucible in a muffle oven at 960 ⬚C for 20–30 minutes. The crucible is allowed to cool and the mass of the crucible and contents recorded. Hazards The Ag2S preparation should be accomplished in a hood. Because SO2 is toxic and some people have severe allergic reactions to it, the laboratory should be well ventilated. During the student-conducted portion of the exercise, the typical laboratory will experience SO2 levels below OSHA limits. Considerable care taken while handling hot crucibles. Silver metal is not considered toxic, but almost all silver compounds pose chemical hazards.
Sample Preparation The silver sulfide sample provided to students can be purchased or produced by the reaction of a silver nitrate solution with an ammonium sulfide solution, details are given in the Supplemental Material.W
Table 1. Reactions in the Experiment
Ag2S(s) + O2(g) Ag2S(s) + 2 O2(g) Ag2SO4(l)
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2Ag(s) + SO2(g)
(1) (2)
Ag2SO4(s)
2Ag(l) + SO2(g) + O2(g) (3)
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Figure 1. Thermogravimetric analysis of silver sulfide and silver sulfate.
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In the Laboratory
Discussion
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A summary of the reactions that transpire during the exercise is given in Table 1. The sequence of reactions can be illustrated with the results of a thermogravimetric analysis carried out in a Rheometrics STA 1500 thermal analyzer in quartz crucibles, with heating rate of 2 ⬚C兾min and an air flow rate of 40 mL/min. The results of the analysis using silver sulfide and silver sulfate samples are shown in Figure 1. The silver sulfide sample incurs a small mass loss before the temperature reaches 200 ⬚C. That change may result from the presence of ammonium nitrate remaining in the sample resulting from the method of the synthesis. Using a higher drying temperature in the sample preparation might reduce that preliminary mass change. Around 600 ⬚C the silver sulfide rapidly loses mass by oxidation of the anion sulfide to SO2 by atmospheric oxygen. A small mass increase resulting from the oxidation of the silver sulfide to silver sulfate occurs at about 620 ⬚C. The reaction is likely catalyzed by the metallic silver that was formed at lower temperatures. Figure 1 also shows that silver sulfate requires higher temperatures for its decomposition. The exercise has been used in introductory chemistry at this university for five years. During that time, the reported ratio, (mole silver)兾(mole sulfur), values have been between 1.96 and 2.05. A group of 35 students can complete experiment in less than two hours.
Instructions for the students and notes for the instructor are available in this issue of JCE Online.
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Supplemental Material
Literature Cited 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15.
Jensen, W. B. J. Chem Educ. 2004, 81, 1772–1774. DeMeo S. J. Chem Educ. 2001, 78, 1050–1052. Sanger, M. J.; Geer, K. J. Chem Educ. 2002, 79, 994–996. Johnson, S. D. J. Chem Educ. 1996, 73, 1179–1181. Wolthuis, E.; DeVries, D.; Poutsma, M. J. Chem. Educ. 1957, 34, 133–134. Masterton, W. L.; Demo, J. J., Jr. J. Chem. Educ. 1958, 35, 242–244. Carmody, W. R. J. Chem. Educ. 1967, 44, 416. Dingledy, D.; Barnard, W. M. J. Chem. Educ. 1967, 44, 693–694. Zidick, C.; Weismann, T. J. Chem. Educ. 1973, 50, 717–718. Hoffman, A. B.; Hoffman, A. J. J. Chem. Educ. 1974, 51, 418–420. Wells, N.; Boschmann, E. J. Chem. Educ. 1977, 54, 586. MacDonald, D. J. J. Chem. Educ. 1983, 60, 147. Schaeffer, R. W.; Chan, B.; Molinaro, M.; Morissey, S.; Yoder, C. H.; Yoder, C. S.; Shenk, S. J. Chem. Educ. 1997, 74, 575–577. Margolis, L. A.; Scheffer, R. W.; Yoder, C. H. J. Chem. Educ. 2001, 78, 235–236. JCE Classroom Activity #25. J. Chem. Educ. 2000, 77, 328A– 329B.
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