An Estimate of the Molar Heat of Reaction for the Decomposition of Hydrogen Peroxide

An Estimate of the Molar Heat of Reaction for the Decomposition of Hydrogen Peroxide.
Frank Schmidt, Biochemistry Department, University of Missouri-Columbia, Columbia MO 65212.

Abstract:

The heat of decomposition of hydrogen peroxide was measured in a Styrofoam (TM) cup which functioned as a simple calorimetry vessel. The reaction was catalyzed by 0.05 M. ferric nitrate, and proceeded over approximately a 5 minute interval. The increase in temperature of the reaction mixture indicated that the ΔH of this exothermic reaction was determined to be - 94,000 ± 3000 joules per mole of H₂O₂.

Introduction:

Hydrogen peroxide, H₂O₂, is an oxidizing agent, commonly used for disinfection and bleaching (1, 2). The products of the reaction are water and molecular oxygen according to the overall equation

(1) H₂O₂6 H₂O + ½ O₂

H₂O₂ releases reactive oxygen species during its decomposition (2). These reactive oxygen species react quickly with biological molecules and many colored compounds, causing them to degrade. This property accounts for the widespread use of H₂O₂ as a disinfectant and bleach. The reactive oxygen species combine with each other to form molecular oxygen if they don't react with another molecule.

The fact that the decomposition of H₂O₂ occur means that reaction in equation (1) proceeds with a negative free energy change. The free energy change, ΔG , is defined as

(2) ΔG = ΔH - TΔS

where ΔH is the heat change associated with a reaction carried out at a constant temperature, T is the absolute temperature of the reaction, and ΔS is the change in entropy (molecular disorder) associated with the reaction. ΔH is negative if the reaction releases heat to the environment and positive if the reaction removes heat from the environment.

I was interested in the apparent contradiction between the ability to decompose quickly when used as a disinfectant or bleaching agent and its apparent stability (2). H₂O₂is classified as a highly reactive molecule and solutions are labeled as such as a safety precaution; however, I knew from personal experience that solutions of H₂O₂are stable for many months (e.g., on the shelf of a medicine cabinet). The solution could lie in a number of catalysts that allow H₂O₂ to decompose more rapidly than it would in their absence. These catalysts can be either biological or chemical in origin (2).

In order to protect themselves from the harmful effects of reactive oxygen, cells have evolved enzymes, called peroxidases. Peroxidases apparently protect cells by speeding up the decomposition of H₂O₂ (that is, they catalyze the reaction shown in equation (1) so that molecular oxygen is produced and other, more harmful, species aren't generated in such high amounts.

The decomposition of H₂O₂ can also be catalyzed by a number of metal salts. These catalysts are not as effective on a molecular basis as peroxidase enzymes. Whether the reaction is spontaneous or catalyzed, the products of decomposition, water and oxygen, are the same. This observation is consistent with the general rule that catalysts change the rate of a reaction but not its products.

I hypothesized that the free energy of decomposition of H₂O₂ would be negative. As part of this overall project, I measured the change in temperature of a solution of H₂O₂ when it decomposed to water and oxygen. I used ferric nitrate, Fe(NO₃)₃, as catalyst. Consistent with the hypothesis, the temperature of the solution increased during the reaction, indicating that heat was released from the decomposition of H₂O₂. The heat of reaction was determined to be - 94,000 ± 3000 joules per mole of H₂O₂.

Experimental Procedures.

Safety and environmental precautions. H₂O₂ and Fe(NO₃)₃ are both classified as reactive oxidizing agents. Safety precautions included wearing chemical safety glasses throughout the experiment and preparing dilutions by the addition of concentrated solutions to water rather by the addition of water to concentrated solutions. The plastic cups used for the reaction were clearly marked as not for food or drink. All waste solutions were placed in double-contained plastic bottles for disposal by the University of Missouri-Columbia Environmental Health and Safety Office.

Reagents. A fresh bottle of 30% H₂O₂ by weight and an unopened container of Fe(NO₃)₃ (both Fisher Reagent Grade) were obtained from the MU Chemistry Stores.

Calorimetry. A 240 mL insulated plastic cup (Styrofoam (TM)) intended for hot drinks was used as an inexpensive calorimetry vessel. The H₂O₂ solution was prepared by diluting 5 mL of 30% H₂O₂ to 50 mL with water. Fifty mL of this solution were added to the cup. The cup was placed on a stir plate so that the bar turned about 3 times per sec. A thermometer was clamped above the cup, positioned so that the bulb was covered as much as possible by the solution without the stir bar hitting it. The temperature of the H₂O₂ solution was recorded at 1 min. intervals for 5 min. Then 50 mL of 0.1 M Fe(NO₃)₃ were added all at once. The temperature was recorded at 1 min. intervals for a further 15 min.

Data treatment. The solution temperature was plotted as a function of time. Heat loss from the calorimeter during and after the reaction was accounted for by extrapolating the line of temperature vs. time back to the time of mixing. The temperature value at this point was taken as the maximum temperature (T_max) for the reaction. The difference between T_max and the initial temperature (T₀), ΔT, was used to calculate the heat released by the reaction.

Results.

Reaction course. Immediately after addition of the light-brown solution of Fe(NO₃)₃ to the H₂O₂ solution, the reaction mixture turned dark brown and oxygen gas bubbled out of solution. The temperature of the solution increased, reaching a maximum 3 min. after mixing. The solution color changed throughout the reaction so that at 5 min. after mixing, very few bubbles evolved and the Fe(NO₃)₃ had returned to its original color.

Temperature and changes during reaction. I did three separate trials of the experiment. The plot of temperature vs. time for a single experiment is shown in Fig. 1. The temperature of the reaction mixture was constant before the reaction began and rapidly increased as the catalyst was added. The maximum temperature observed was 33.2E.

I calculated ΔT by extrapolating the line of temperature vs. time back to the time of mixing and using this temperature as T_max for calculation of ΔT by Equation 3:

(3) ΔT = T_max - T₀

where T₀ is the temperature of the H₂O₂ solution before addition of the catalyst. The mean value of ΔT was 11.3 ± 0.3 degrees.

Using the value of ΔT above, I calculated the heat released during the reaction. The heat capacity of water is 4.18 J/(degree mL). I ignored any contribution to the heat capacity from the Fe(NO₃)₃ since it was in aqueous solution. The solution contained 100 mL. so the heat released by the reaction was

(4) ΔT × volume × heat capacity = 11.3 ×100 × 4.18 = 4720 J.

Since heat was given off by the reactants to the water in the solution ΔH = - 4720 J.

Heat given off per mole of reactant. In order to calculate the heat of reaction on a molar basis it was first necessary to estimate the amount of H₂O₂ in the solution. The original concentration of H₂O₂ in the reagent bottle was 30% by weight (i.e., 30 g H₂O₂ per 100 g of solution). The density of 30% H₂O₂ solution is 1.11 g/mL (3) so that the weight of H₂O₂ in 1 mL of 30% solution is given by Equation 5.

(5) (30 g H₂O₂ /100 g solution ) × 1.11 g solution = 0.333 g H₂O₂

Since the molecular weight of H₂O₂ is 34, and since I used 5 mL of 30% H₂O₂ per reaction, the reaction mixture contained 0.050 moles of H₂O₂, according to equations 6 and 7:

(6) 0.333 (g/ mL of 30% H₂O₂) × 5 (mL of 30% H₂O₂ per reaction) = 1.7 g H₂O₂ per reaction

(7) 1.7 g H₂O₂ per reaction / 34 (g/mole H₂O₂) = 0.050 moles of H₂O₂ per reaction.

This determination allowed the molar heat of decomposition to be calculated as

(8) Heat/mole = - 4720 J / 0.050 moles H₂O₂ = - 94, 500 J/mole.

Experimental uncertainty. The mean value of ΔT was 11.3 ± 0.3 degrees. The percentage error was

(9) 0.3 /11.3 = 0.03 = 3 %

Since standard errors are only given to one significant figure, and since the final determination can't be more precise than the measurement of ΔT , my estimate of the ΔH per mole of reaction could only be precise to ± 3%. Three per cent of -94,500 is - 3,000, so the final value of the specific heat of decomposition of H₂O₂ is - 94,000 ± 3000 joules per mole.

Discussion:

The purposes of this study were: first, to determine the molar heat of decomposition of H₂O₂, and secondly, to evaluate the utility of insulated plastic cups as inexpensive calorimetry vessels.

The molar heat of reaction, ΔH, was determined to be -94,000 ± 3,000 joules per mole. Thus, energy in the form of heat was given off in the experiment. The chemical literature gives a ΔH for this reaction of -94,500 J/mole (4). This value is well within the experimental error of my determination. The chief source of error in my determination was likely to be the estimation of the temperature of the reaction. I used a standard laboratory thermometer for this experiment. This thermometer has gradations of one degree and requires that tenths of a degree be estimated visually. A more accurate thermometer with finer calibration, as is used in calorimetry research, would have reduced this source of error. Other potential sources of nonsystematic error, such as from volume measurement, are included in the standard error of the experimental result.

I concluded that these cups were adequate for calorimetry of this reaction in aqueous solution. I knew from previous experience that the cups are imperfect insulators, since hot drinks kept in them do not stay hot forever but rather lose heat to the environment. This meant that I had to correct for this heat loss in determining ΔT. The rate of heat loss was approximately 0.2 degrees per minute (Figure 1). This rate of heat loss was corrected for by extrapolating the temperature to the instant of mixing; this value was used as T_max. When this correction was made, T_max and therefore ΔT increased by 0.3E. A better-insulated reaction vessel would have made this correction smaller; however, it still would likely have been necessary.

Since H₂O₂ decomposes on its own, it is also possible that the concentration of H₂O₂ in the stock solution was less than 30% by weight, even though I used a fresh bottle and kept it refrigerated until the experiment. This inaccuracy would have made the determination of ΔH per mole too low, since less heat would have evolved but I still would have assumed that 0.05 moles of H₂O₂ was present in the reaction mixture.

The reaction proceeded in the presence of a catalyst, in this case, Fe(NO₃)₃. Catalysts are not themselves changed by participating in a chemical reaction. My observation that the color of the reaction returned to that of the Fe(NO₃)₃ solution at the end of the reaction was consistent with this definition. If the Fe(NO₃)₃ catalyst had been altered during the reaction, I would have expected it to have a different color at the end of the reaction.

Heat was a product of the decomposition of H₂O₂, in addition to water and molecular oxygen. This observation has several implications. Because ΔH is a thermodynamic function of the reaction itself, and not of the rate at which it proceeds, one could measure the amount of H₂O₂ in an unknown solution by measuring the heat given off during its decomposition. More interestingly, since catalysts alter the rate but not the thermodynamic or equilibrium properties of a chemical reaction, I would predict that different catalysts or amounts of a single catalyst shouldn't affect the ΔH of the reaction, irrespective of the differences in the rate of reaction under varying conditions. It would be interesting to test these predictions by determining the molar ΔH for decomposition of H₂O₂in the presence of a different catalyst, for example a peroxidase enzyme, or in the presence of differing concentrations of Fe(NO₃)₃.