Improvements upon the âColorful Cobalt Catalysis ... - ACS Publications

In the Classroom edited by

Todd P. Silverstein Willamette University Salem, OR 97301-3922

Improvements upon the “Colorful Cobalt Catalysis” Demonstration and Evidence for the Presence of an Autocatalytic Mechanism Stephen W. Wright Pfizer Global Research and Development, Eastern Point Road, Groton, Connecticut 06340 [email protected]

The cobalt-catalyzed oxidation of tartrate ion by hydrogen peroxide is one of the most popular demonstrations of catalysis presently available to instructors teaching high school and introductory college chemistry. First reported in 1974 (1), it was published in this Journal in 1978 (2) and has been summarized in collections of chemical demonstrations (3). This demonstration captivated me when I first saw it in high school. In connection with my chemical demonstration programs at local schools, this experiment has been well received by students but frequently presented technical difficulty with its preparation and presentation. This article describes some modifications of the experiment that make it simpler and more convenient to present. Furthermore, a chance observation and subsequent experimentation recounted here may lead to greater understanding of the complex mechanisms underlying this elegant demonstration.

replace the 30% hydrogen peroxide with 3% hydrogen peroxide, and I also sought to develop a heating protocol that would eliminate the need to monitor the temperature. Ideally, I wished to identify conditions under which the experiment would proceed without heating. Because students intuitively and correctly associate heating with an increase in reaction rate, the action of the catalyst upon the rate of reaction would be much more apparent to students if the reaction took place at room temperature. As described by Ruda (2), the demonstration calls for two solutions prepared as follows: 25 g (0.089 mol) of potassium sodium tartrate tetrahydrate in 300 mL of water and 20 mL (0.18 mol) of 30% hydrogen peroxide in 80 mL of water (2-fold molar excess). The two solutions are combined and heated to between 50 and 70 °C, and the cobalt(II) chloride catalyst solution is added to initiate the catalyzed reaction.

Background

Modification of the Hydrogen Peroxide Solution

The original demonstration is carried out by heating an aqueous solution of potassium sodium tartrate and hydrogen peroxide to about 60 °C. The warm solution is removed from heating and shown to the students, who see that it is colorless and more or less unremarkable. A small volume of pink cobalt(II) chloride solution is added, and the reaction mixture likewise becomes pink. After about 10 s, the mixture becomes green and gas evolution begins. Gas evolution becomes increasingly brisk while the mixture is green, to the point that the mixture appears to be in danger of foaming over. Then, the gas evolution stops abruptly, after which the mixture turns pink over the course of about a minute. It is easy for students to understand that the addition of the cobalt(II) chloride solution caused a reaction to occur, that the reaction occurred while the mixture was green, and that the cobalt(II) chloride appeared to be regenerated at the end of the reaction. As originally published, the demonstration has two minor defects that make it inconvenient to prepare and to present. First, the experiment as described requires the use of 30% hydrogen peroxide. At the time the experiment was first published, a 6% hydrogen peroxide solution was available on the consumer market that made a satisfactory substitute for the diluted 30% hydrogen peroxide,1 but this product is no longer marketed. Second, the experiment requires that the potassium sodium tartrate solution and hydrogen peroxide solutions be heated to about 65 °C before they are combined. This requires that the solutions be preheated starting prior to the class period and then carefully monitored during the class period to ensure that the solutions are not under- or overheated. Therefore, I sought to

Examination of the molar amounts of reactants used in the demonstration shows that the final concentration of hydrogen peroxide in the reaction mixture is 0.47 M, which is approximately half the concentration of 3% hydrogen peroxide (0.9 M). In as much as potassium sodium tartrate is freely soluble in water, a simple adjustment of the volume of water used to prepare the two solutions permits the use of 3% hydrogen peroxide in place of the 30% solution. Furthermore, if the solutions were not combined until use and the volumes were adjusted appropriately, then the tartrate solution could be heated to boiling and kept at that temperature until needed while the peroxide solution could be kept at room temperature. A heat capacity calculation shows that when combined, the two solutions would afford a mixture at approximately 65 °C. Solutions were therefore prepared as follows: 30 g (0.11 mol) of potassium sodium tartrate tetrahydrate in 360 mL of water and 240 mL (0.22 mol) of 3% hydrogen peroxide (2 molar excess) The solution of potassium sodium tartrate in 360 mL of water was heated to boiling. At the time of the demonstration, it was removed from the heat and diluted with 240 mL of room temperature 3% hydrogen peroxide solution. The temperature of the resulting mixture was between 65 and 70 °C, just as needed to carry out the cobalt-catalyzed demonstration.

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Concentration of Reactants Having devised a means to simplify the heating process, I wondered if the demonstration could be adapted to proceed at room temperature. This required reformulation of the

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Vol. 87 No. 10 October 2010 pubs.acs.org/jchemeduc r 2010 American Chemical Society and Division of Chemical Education, Inc. 10.1021/ed100214y Published on Web 07/30/2010

In the Classroom

demonstration. To compensate for the reduced temperature, I increased the concentration of reactants, but elected to continue to use the commercially available 3% hydrogen peroxide. I found that a solution composed of 9.88 g (0.035 mol) of potassium sodium tartrate tetrahydrate in 80 mL (0.072 mol) of 3% hydrogen peroxide did undergo the characteristic pink to green to pink reaction when treated with a 10 mL of 0.17 M cobalt(II) chloride catalyst solution. The reaction mixture became noticeably brown after 2 min, subsequently turned green and evolved gas, and by 15 min the gas evolution had ceased and the mixture was pink again.

tart2 - ðaqÞ þ CO2 ðaqÞ þ H2 OðaqÞ f Htart - ðaqÞ þ HCO3 - ðaqÞ in which the tartrate ion represented as tart2-.

Examining pH Encouraged by this result, but disappointed by the length of time required for the reaction to go to completion, I tried the reaction under various pH conditions. The addition of small quantities of various acids produced no effect, but a remarkable reduction in the time required to complete the reaction, to 5 min or less, was observed when as little as 50 mg of sodium bicarbonate was included. An experiment in which an equimolar amount of the slightly alkaline disodium phosphate was substituted for 1 g of sodium bicarbonate resulted in the slow reaction observed originally. Coupled with the observation that even very small quantities of bicarbonate resulted in an obvious increase in rate, I ruled out the possibility that these results were due to the slight pH increase of the reaction mixture. A literature search on this demonstration revealed no new information except for a letter to the editor of this Journal that stated that the gas evolved by the reaction is a mixture of oxygen and carbon dioxide, with the proportion of oxygen decreasing as the reaction proceeds (4). This led me to imagine that perhaps the rate enhancement was due in some way to the specific presence of the bicarbonate ion, which would imply that the reaction is not only catalyzed by cobalt salts, but may also be autocatalytic with respect to the carbon dioxide generated, some of which will dissolve in the mixture and form bicarbonate ion. This would be consistent with the induction period observed when the reaction is carried out at room temperature, a phenomenon that is obviated when the reaction is carried out with heating, as is customarily done. An experiment was prepared in which 1 g of dry ice was added to the reaction mixture; the solid carbon dioxide was allowed to sublime and dissolve in the solution before the catalyst solution was added. The same rate enhancement occurred that had been observed with sodium bicarbonate. The rate enhancement was also observed when the reaction was carried out in the presence of hydrogen peroxide and

• tartaric acid and sodium bicarbonate (1:2 mol ratio) • tartaric acid and sodium carbonate (1:1 mol ratio) • sodium hydrogen tartrate and sodium bicarbonate (1:1 mol ratio) • potassium sodium tartrate, sodium bicarbonate, and hydrochloric acid (1:0.09:0.09 mol ratio) • potassium sodium tartrate and carbon dioxide/bicarbonate (from club soda)

These results show that the addition of bicarbonate ion to the reaction mixture accelerates the room temperature reaction and the addition of carbon dioxide (generated in situ or introduced as dry ice or club soda) likewise accelerates the room temperature reaction. Together with the observation that

r 2010 American Chemical Society and Division of Chemical Education, Inc.

cobalt(II) catalyzes the decomposition of hydrogen peroxide and forms a green intermediate when bicarbonate ion, but not carbon dioxide, is added to the mixture (vide infra), it is clear that bicarbonate ion is a key component of the active catalytic species. Carbon dioxide may serve as a source of bicarbonate ion, dissolving in the weakly basic solution of tartrate ions, according to the equation:

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-

ð1Þ

O2CCH(OH)CH(OH)CO2- is

Mixing Order Not surprisingly, no rate enhancement was observed when the reaction was carried out using tartaric acid and sodium hydroxide (1:2 mol ratio). A set of experiments was carried out in which two of the three reagents, hydrogen peroxide, cobalt(II) chloride, and potassium sodium tartrate, were mixed and equilibrated overnight at room temperature, after which the third reagent was added. Results of these experiments showed that the order of mixing had no effect on the rate of the reaction (in the absence of added bicarbonate ion or carbon dioxide). I also considered the possibility that dissolved oxygen might somehow retard the onset of the reaction, and that pretreatment of the reaction mixture with carbon dioxide resulted in a rate enhancement due to degassing of the solution. However, no rate enhancement was observed when the solution was degassed by a nitrogen sparge prior to the experiment. Addition of Bicarbonate The quantity of bicarbonate introduced into the reaction mixture must be carefully controlled: Too little results in a demonstration that takes too long to complete and loses the students' attention and too much bicarbonate results in an immediate change in color of the reaction mixture to green upon addition of the catalyst solution. If the quantity of bicarbonate is carefully controlled, the reaction mixture remains pink for a brief period after the addition of the catalyst solution; this is followed by the appearance of the green color and induction of the tartrate oxidation and gas evolution reaction, and finally the return of the pink color at the end of the reaction. Students thus visually (and correctly) associate the pink chromophore with the inactive catalyst cobalt(II), and the green chromophore with the active catalyst cobalt(III). Reaction Temperature versus Reaction Time The temperature of the reaction was monitored as a function of time and relative reaction rates were characterized by comparing the time elapsed from the addition of the cobalt(II) chloride catalyst solution to the time at which the internal temperature reached its peak and the gas evolution abruptly slackened. A series of experiments was conducted in which the reaction mixture was treated with varying quantities of sodium bicarbonate prior to the addition of the cobalt(II) chloride catalyst solution. Reaction mixtures were adjusted to 20 °C in each case prior to the start of the experiment, to compensate for the endothermicity observed when potassium sodium tartrate is

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Figure 1. Reaction temperature versus time elapsed after addition of cobalt(II) chloride solution (0.17 M, 10 mL) to a mixture of potassium sodium tartrate (9.9 g, 0.035 mol) and 3% hydrogen peroxide (80 mL, 0.072 mol) in the presence of varying quantities of sodium bicarbonate.

dissolved in water. As is evident from the data in Figure 1, increasing quantities of sodium bicarbonate led to decreased times required to reach the maximum reaction temperature. As seen in Figure 1, the addition of 10 mL of 0.17 M cobalt(II) chloride to 9.9 g of potassium sodium tartrate tetrahydrate in 80 mL of 3% hydrogen peroxide resulted in an increase in temperature of about 35 to 37 °C. An experiment in which 80 mL of 3% hydrogen peroxide was decomposed by the addition of 10 mL of 0.02 N potassium permanganate was found to produce a 17 °C increase. Thus, the heat generated by the reaction mixture in the presence of tartrate ion cannot be due solely to the decomposition of hydrogen peroxide alone and is therefore derived from the oxidation of tartrate ion. Catalytic Species Experiments were also carried out to determine what species in the demonstration catalyze the decomposition of hydrogen peroxide and what species in the demonstration may be responsible for the transient green color observed as the reaction progresses. A mixture of 80 mL of 3% hydrogen peroxide and 10 mL of 0.17 M cobalt(II) chloride treated with 0.10 g of sodium bicarbonate changed from a pink to a green color and evolved oxygen gas over a period of about 20 min, accompanied by an increase in temperature of about 15 °C, consistent with the catalyzed decomposition of hydrogen peroxide.2 It is known that cobalt carbonate, cobalt oxide, and the Co(HCO3)þ species will catalyze the decomposition of hydrogen peroxide to water and oxygen (5). The observation by Toth (4) that the gas evolved by the demonstration is a mixture of oxygen and carbon dioxide, with the proportion of oxygen decreasing as the reaction proceeds, is therefore understandable if the catalyzed decomposition of hydrogen peroxide occurs first, and the oxidation of tartrate ion to produce carbon dioxide is autocatalytic and sets in after the passage of an induction period, after which it occurs at an accelerating rate.3 The species responsible for the green color that is observed as the demonstration proceeds has not been fully characterized. It has generally been described as an “activated complex” containing 1066


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cobalt ion(s) and tartrate ion(s) (see for examples ref 3). A green compound produced by the reaction of oxidizing agents with cobalt(II) in the presence of bicarbonate has been known for over 100 years (6). In 1942, Bobtelsky (7) noted that combination of solutions of cobalt(II) chloride and sodium citrate afforded a pink mixture that turned green upon the addition of hydrogen peroxide. It was observed that the green mixture decomposed hydrogen peroxide to liberate oxygen, and that the rate of oxygen production was significantly increased as the concentration of citrate exceeded that of cobalt(II). Bobtelksy speculated that the green substance contained cobalt “only in the divalent state”. However, earlier work by Swann and Xanthakos (8) on the oxidizing properties of cobalt(III) sulfate prepared by anodic oxidation described the compound as “greenish-blue” crystals and showed that cobalt(III) sulfate oxidized tartaric acid to afford carbon dioxide. They also reported that cobalt(III) sulfate decomposed in aqueous solutions at room temperature to evolve oxygen gas. These observations suggest that the green color is more likely due to a cobalt(III) species than to a cobalt(II) species. In this work, all solutions containing the green cobalt complex gave a positive test for peroxide4 and oxidized starch-iodide paper (moistened with dilute acetic acid) to iodine. The pink solution remaining after the completion of the reaction with tartrate ion gave a negative test for peroxide and failed to oxidize starch-iodide paper to iodine. Katz has speculated that the green “activated complex” may be a dinuclear cobalt complex with tartrate ion as a bidentate or bridging ligand.5 However, I have found that the green complex also forms when the tartrate ion is replaced by either bicarbonate ion or acetate ion.6 This shows that the green complex does not require tartrate ion for its formation and therefore does not require the dinuclear structure with tartrate ion as a bidentate or bridging ligand. On the basis of the literature descriptions and my results, I would suggest that the green color is due to a cobalt(III) species containing both bicarbonate ion and carboxylate ion ligands. The possible involvement of the peroxybicarbonate (HCO4-) ion, formed from hydrogen peroxide and carbon dioxide or bicarbonate ion, must also be considered. Reaction Stoichiometry The reaction stoichiometry merits some discussion. As originally published, and as described below for the unheated reaction, the initial reaction mixture contains hydrogen peroxide in approximately 2-fold molar excess to tartrate ion. This would suggest that the tartrate ion is the limiting reagent, and indeed Ruda appears to have come to that conclusion in 1978 (2), stating that the cobalt forms “a green complex with the tartrate” and that “as tartrate is oxidized, the activated complex is destroyed”. However, both the initial pink reaction mixture and the green mixture give a positive test for the presence of hydrogen peroxide. On the other hand, at the conclusion of the reaction, the final pink mixture gives a negative test for the presence of hydrogen peroxide. Furthermore, if the amount of tartrate ion in the mixture is reduced, while keeping the amount of hydrogen peroxide constant, the mixture will remain green after the exothermic reaction and gas evolution have ceased, and the mixture will give a positive test for the presence of peroxide. Taken together, these data clearly indicate that in fact hydrogen peroxide is the limiting reagent. This was recognized by Toth in 1980 (4), who observed that “any further reaction can be induced


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In the Classroom

only by the addition of hydrogen peroxide”. This is possible for three reasons: (a) the evolution of oxygen gas, as noted by Toth (4), indicates that hydrogen peroxide is consumed by a competing pathway in which hydrogen peroxide is catalytically disproportionated by the cobalt(II) and cobalt(III) species present; (b) cobalt(II) oxidation to cobalt(III) consumes hydrogen peroxide; and (c) the initial product(s) resulting from the oxidation of tartrate ion may themselves undergo further oxidation by hydrogen peroxide.7

Scheme 1. A Plausible Sequence of Reactions for The Cobalt Demonstration

Plausible Reaction Sequence The reaction mechanism and the product(s) produced by the oxidation of the tartrate ion, besides carbon dioxide, are not known with certainty. Oxidation of tartrate ion by cobalt(III) may occur via formation of a cyclic intermediate between the cobalt ion and the tartarte ion, as occurs in the oxidation of similar 1,2-diols by permanganate ion, or oxidation may occur by single electron transfer, or by a combination of these mechanisms, and further study of these questions is needed. A plausible sequence of reactions for this experiment is shown in Scheme 1. The active cobalt(III) catalyst is generated by hydrogen peroxide oxidation of cobalt(II) (eq 2), resulting in one electron reduction of hydrogen peroxide to water. It is known that iron(II) reacts with hydrogen peroxide in this manner (9); however, it is possible that the oxidation of cobalt(II) may occur by a process other than that shown in eq 2. The signature gas evolution that is observed in this experiment is the result of two catalytic steps involving cobalt(III). Cobalt(III) catalyzes the disproportionation of hydrogen peroxide to oxygen gas and water (eq 3). Cobalt(III) also catalyzes the oxidation of tartrate ion (eq 4), initiating a series of events that leads to the production of carbon dioxide (eqs 5 and 7). Two-electron oxidation of tartrate ion 1 affords oxaloglycolate 2 (eq 4).8 Oxaloglycolate may spontaneously evolve carbon dioxide9 to give 3-hydroxypyruvate 3a (eq 5), which may interconvert to tartronic semialdehyde 3c via the ene-diol tautomer 3b (eq 6) (10). Tartronic semialdehyde 3c may in turn undergo a second spontaneous loss of carbon dioxide to afford glycolaldehyde 4 (eq 7) (11). Further oxidation of glycolaldehyde to glyoxal, glyoxylate, and oxalate may be possible. When hydrogen peroxide is consumed, the remaining cobalt(III) is reduced by the carbon compounds remaining in the solution to return the pink color of cobalt(II) (eq 8). It should be noted that the oxidation of tartrate and subsequent decarboxylation to produce glycolaldehyde consumes only 1 mol of hydrogen peroxide per mole of tartrate. From the above discussion, it is clear that hydrogen peroxide is the limiting reactant, even though it is present in 2-fold molar excess over tartrate, because the mixture is pink and hydrogen peroxide is absent at the end of the reaction. This discrepancy can be explained by the fact that hydrogen peroxide is consumed not only by tartrate oxidation, but also by catalytic disproportionation (eq 3), catalyst generation (eq 2), and possibly by further oxidation of glycolaldeyde. Summary I have developed means by which the popular lecture demonstration oxidation of tartrate ion by hydrogen peroxide, catalyzed by cobalt(II) chloride, may be conducted with 3% hydrogen peroxide and without careful temperature control.


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I further show that the demonstration may be carried out at room temperature and that the reaction is subject to an induction period before the exothermic reaction and evolution of carbon dioxide occur. The duration of this induction period may be reduced or eliminated by the introduction of carbon dioxide or bicarbonate ion into the reaction mixture prior to the addition of the cobalt(II) chloride catalyst. It is therefore proposed that the induction period observed is due to a slow initial rate of tartrate oxidative decarboxylation, yielding carbon dioxide. As carbon dioxide is released, the concentration of bicarbonate ion in the solution increases and the oxidation of tartrate ion occurs


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more rapidly. The reaction, therefore, is catalytic in cobalt(II) chloride, but in addition, an autocatalytic mechanism involving carbon dioxide and bicarbonate is present. I provide additional evidence that hydrogen peroxide is the limiting reagent when the reaction is conducted in the customary manner. I further show that a green species may be prepared from cobalt(II) chloride and hydrogen peroxide when either bicarbonate ion or carbon dioxide and acetate ion are added, that the presence of tartrate ion is not necessary for the production of the green color, and that the chemical literature indicates that the green color is consistent with the presence of cobalt(III). Experimental Section Deionized water was used for all experiments. A 0.17 M solution of cobalt(II) chloride was prepared by dissolving 10 g of cobalt(II) chloride hexahydrate in 250 mL of water and was used as catalyst for all experiments. The starting temperature of the unheated reactions should be about 20 °C; if colder than about 17 °C, the reaction takes noticeably longer. The dissolution of potassium sodium tartrate is endothermic; if necessary, the mixture should be warmed to 20 °C. The unheated reactions are significantly exothermic, resulting in a final temperature of about 55 °C. Procedure 1 (Heated Reaction) A solution of 30 g (0.11 mol) of potassium sodium tartrate tetrahydrate in 360 mL of water in a 1 L beaker is covered with a watch glass. The liquid level is marked and the beaker is heated to a gentle boil. Any water lost during heating may be replaced by adding water from a wash bottle to restore the original volume. To perform the demonstration, the beaker is removed from the heat source and 240 mL (0.22 mol) of 3% hydrogen peroxide is added. The mixture is now at about 65 °C. The catalyst solution (10 mL of 0.17 M CoCl2) is added and the reaction is observed. The temperature typically increases to about 80 °C. A second beaker may be prepared, untreated with catalyst solution, as a control. Procedure 2 (Unheated Reaction without Added Bicarbonate Ion) Potassium sodium tartrate tetrahydrate (9.9 g, 0.035 mol) is dissolved in 80 mL (0.072 mol) of 3% hydrogen peroxide in a 250 mL beaker. To perform the demonstration, the catalyst solution (10 mL of 0.17 M CoCl2) is added and the reaction is observed. The temperature typically increases by about 35 °C over about 15 min. A second beaker may be prepared, untreated with catalyst solution, as a control. Procedure 3 (Unheated Reaction, Autocatalysis by Bicarbonate Ion) Potassium sodium tartrate tetrahydrate (9.9 g, 0.035 mol) is dissolved in 80 mL (0.072 mol) of 3% hydrogen peroxide in a 250 mL beaker. Sodium bicarbonate (0.10 g, 1 mmol) is added with stirring. To perform the demonstration, the catalyst solution (10 mL of 0.17 M CoCl2) is added and the reaction is observed. The temperature typically increases to about 55 °C. Note: It may be necessary to use lesser quantities of sodium bicarbonate to adjust the reactions induction period to a convenient length of time. Too much sodium bicarbonate will cause 1068


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the reaction to commence immediately, and students may not have time to observe the initial pink color or the mixture. Procedure 4 (Unheated Reaction, Autocatalysis by Carbon Dioxide from Supermarket Club Soda) Potassium sodium tartrate tetrahydrate (9.9 g, 0.035 mol) is dissolved in 72 mL of unflavored club soda from a freshly opened bottle in a 250 mL beaker. Some evolution of CO2 may occur during the preparation of the solution. After the potassium sodium tartrate has fully dissolved, 8 mL (0.072 mol) of 30% hydrogen peroxide is added to the solution. To perform the demonstration, the catalyst solution (10 mL of 0.17 M CoCl2) is added and the reaction is observed. The temperature typically increases to about 55 °C. Hazards Safety glasses must be worn when handling these or any other chemicals. All chemicals must be kept away from the audience at all times, particularly children. The heated experiment (procedure 1) involves liquids hot enough to cause severe burns if spilled on the body. Hydrogen peroxide is a strong oxidizer and is corrosive causing burns to the eyes, skin, and respiratory tract. Cobalt(II) chloride hexahydrate and potassium sodium tartrate tetrahydrate may cause irritation to skin, eyes, and respiratory tract and may be harmful if swallowed or inhaled. At the conclusion of the experiment, the reaction mixture may be safely discharged to the sanitary drains. Take precautions in case the beaker should overflow. Acknowledgment This article is dedicated to my high school chemistry teacher, Ronald W. Brown, with thanks for his encouragement and support. Stephanie L. Liberatore, assistant editor, The Science Teacher, is thanked for her help in obtaining a copy of the 1974 article by Deroo. I also wish to thank feature editor Todd P. Silverstein for valuable discussions and suggestions. Frederick Sauls (King's College) and Michael Roadruck (Ottawa Hills High School) are thanked for checking the demonstration. Notes 1. A 6% hydrogen peroxide solution was formerly available under the trade name “Clairoxide”. Although hydrogen peroxide is still a component of consumer hair care products, it is no longer available as a 6% solution of well-defined composition. 2. A mixture of 80 mL of 3% hydrogen peroxide and 10 mL of 0.17 M cobalt(II) chloride was found to undergo no apparent change with time. Likewise, the same mixture treated with carbon dioxide (in the form of dry ice) underwent no apparent change. 3. Editor Silverstein has pointed out that this result suggests that the rate constant for hydrogen peroxide disproportionation is substantially lower than that involved in the tartrate oxidation steps. Thus, at reduced concentrations, hydrogen peroxide may preferentially enter into the tartrate oxidation steps that evolve carbon dioxide, rather than the disproportionation step that evolves oxygen. It should also be noted, however, that the twoelectron oxidation of tartrate ion results in the production of 2 mol of carbon dioxide per 1 mol of hydrogen peroxide, whereas the disproportionation of hydrogen peroxide produces 0.5 mol of oxygen per 1 mol of hydrogen peroxide.


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4. Semiquantitative peroxide determination strips, catalog 441601, supplied by J. T. Baker, Phillipsburg, NJ 08865, were used. 5. See ref 3b. An essentially identical discussion of the green intermediate is discussed by Katz at http://www.chymist.com/ colorful%20catalysis.pdf (accessed Jul 2010). Katz asserts that the green species is a binuclear dicobalt peroxo compound, but provides no experimental data to support this assertion. The green cobalt species may or may not be a binuclear peroxo complex; as discussed above, Swann and Xanthakos (8) have shown that cobalt(III) prepared by anodic oxidation is likewise green. Further, although it is entirely possible that tartrate ion does bind as a bridging bidentate chelate to two cobalt atoms in this experiment, as suggested by Katz, in the absence of experimental data this cannot simply be assumed. This work has shown that neither a bridging bidentate ligand capable of coordinating to both cobalt atoms of a binuclear cobalt peroxo species (as shown in Katz's Scheme 5) nor the presence of tartrate ion are necessary for the production of a green color. 6. In these cases, oxygen, but not carbon dioxide, was evolved and the temperature of the mixture increased by about 15 °C until the evolution of gas slackened. This temperature increase is consistent with that which would be expected from the catalytic decomposition of the hydrogen peroxide in the reaction mixture. The resulting mixture remained green after gas evolution ceased and gave a positive test for peroxide. 7. Toth (4) also concluded that hydrogen peroxide is the limiting reagent because there is not enough hydrogen peroxide present to oxidize all of the tartrate ions to carbon dioxide. However, the assumption that the reaction mechanism is competent to oxidize tartrate ion completely to carbon dioxide has not been verified experimentally. Toth observed that if “a sufficient quantity of hydrogen peroxide is provided all of the tartaric acid becomes oxidized, but this is not manifested in any color change. The solution remains green provided hydrogen peroxide is present. With the oxidation of the tartrate salt the pH of the solution increases”. An increase in pH might be anticipated from the consumption of hydronium ions during the decarboxylation and catalyst generation steps (see eqs 5, 7, and 2 in Scheme 1). 8. This appears to be the first step in the oxidation of tartrate by hydrogen peroxide in the presence of iron(II); see: Pozhidaev,


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E. D.; Gorbachev, S. V. Zh. Fiz. Khim., 1965, 39, 1678-1684 (Chem. Abstr. 1965, 63, 71040). Equation 4 may be an oversimplification; the tartrate may be oxidized by cobalt(III), which in turn may be regenerated by hydrogen peroxide oxidation. 9. The spontaneous decarboxylation of β-keto carboxylic acids is well-known in organic chemistry.

Literature Cited 1. (a) Deroo, J. Sci. Teach. 1974, 41, 44. (b) Huxley, J. Chem 13 News 1976, 81, 15. 2. Ruda, P. T. J. Chem. Educ. 1978, 55, 652. 3. (a) Summerlin, L. R.; Ealy, J. L. Chemical Demonstrations: A Sourcebook for Teachers; American Chemical Society: Washington, DC, 1985; p 72. (b) Summerlin, L. R.; Ealy, J. L. Chemical Demonstrations: A Sourcebook for Teachers; American Chemical Society: Washington, DC, 1988; Vol. I, p 103. (c) Katz, D. A. Spectrum 1991, 29, 32–33. 4. Toth, Z. J. Chem. Educ. 1980, 57, 464. 5. Sychev, A. Y.; Isak, V. G.; Van Lap, D. Zh. Fiz. Khim. 1977, 51, 363-367 (Chem. Abstr. 1977, 86, 178055; see also Chem. Abstr. 1987, 106, 183328. 6. Durrant, R. G. J. Chem. Soc., Trans. 1905, 87, 1781–1791. 7. Bobtelsky, M.; Simchen, A. E. J. Am. Chem. Soc. 1942, 64, 2492– 2498. 8. Swann, S., Jr.; Xanthakos, T. S. J. Am. Chem. Soc. 1931, 53, 400– 404. 9. (a) Luehrs, D. C.; Roher, A. E.; Wright, S. W. J. Chem. Educ. 2007, 84, 1290–1291. (b) Prousek, J. Pure Appl. Chem. 2007, 79, 2325– 2338. 10. Sychev, A. Y.; Skutaru, Y. V.; Duka, G. G. Russ. J. Phys. Chem. ( Engl. Trans.) 1987, 61, 1188-1190 (Chem. Abstr. 1987, 108, 130860). 11. Friedemann, T. E. J. Biol. Chem. 1927, 73, 331–334.

Supporting Information Available Procedures and experimental data for the experiments reported in Figures 1. This material is available via the Internet at http://pubs.acs.org.


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