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Physical Chemistry Laboratory [CHEM 335]
Lipi Akhter and Alex Stewart


Hydrogen peroxide, H2O2, decomposes very slowly because of a high activation barrier. Addition of iodide ion increases the reaction rate considerably. In this work, we investigated the decomposition of hydrogen peroxide by recording the rate of oxygen production with various starting concentrations of H2O2 and KI. The reaction exhibited first-order kinetics in both H2O2 and KI. Temperature dependence trials yielded an activation energy of –52.97 +/- 0.24 kJ and an average rate constant of 3.7 with a confidence limit of 2.24±1.4.

Corresponding Author
Contributed similarly to this work


The process of calorimetry pertains to measuring absorbed or evolved heat during a chemical reaction or simply measuring quantitatively the flow of heat in a reaction [1]. The first law of thermodynamics states that “energy can neither be created nor destroyed, but can be transformed from one state to another.” This law is made evident in bomb calorimetry. During this process, when the fuse wire and pellet are ignited in the bomb, heat is the energy seen to be given off during this combustion reaction. The bomb, used in bomb calorimetry, is a completely sealed and oxygen filled metal container. This is placed in an insulated jacket containing a pail of water and a thermometer -- all combined to form the calorimeter [3]. The bomb allows for both a constant volume of the container and for no inflow or outflow of heat (Δq) and its insulated jacket serves as an ideal environment for the bomb. All the conditions surrounding the bomb and its calorimeter, allow for an adiabatic reaction to take place. Hence, the ΔU is equal to the change in work (ΔW) done by the bomb during the reaction [2].
ΔU = ΔW (1)
All through the adiabatic process, the change in heat of the system (coming in or going out) is supposedly zero, rendering the change in enthalpy equal to the change in heat [2].
ΔH = Δq (2)
Since no reaction or the equipment used in the reaction is perfect with 100% efficiency in the process, the adiabatic conditions are also not perfect and some heat does enter and leave the bomb. The typical combustion reaction seems to resemble this form:
Organic Compound (s) + A O2 (g) --> B CO2 (g) + C H2O (g)
and when balanced,
CxHyOz (s) + (X+Y/4-Z/2) O2(g) -----> X CO2(g) + (Y/2) H2O(g) [6]

The organic compounds used in this experiment were benzoic acid (C7H6O2), naphthalene (C10H8) [5], sucrose (C12H22O11) and an unknown compound.
This experiment seeks to determine the heat capacity of the bomb using the standard, benzoic acid. This value is a constant for that particular bomb and will ultimately lead to the determination of the internal energy (ΔU) of naphthalene, sucrose and the unknown, as well as and the molar internal energy (ΔUm) and the molar enthalpy (ΔHm) of naphthalene and sucrose [3].


Fig. 1. (Left to right) Calorimeter and Bomb cross sections. [Shoemaker, D., Experiments in Physical Chemistry, McGraw-Hill, Inc., 1989]

The bomb calorimeter (Fig. 1) was assembled and cleaned. An amount just over 0.800g of benzoic acid was obtained, weighed and compacted into a pill. This pill was placed in the ignition cup and the ends of a 10cm piece of fuse wire were attached to both ends of the electrode terminals while the middle section of the wire was left touching the sample. The head of the bomb was placed in the bomb container and sealed. To remove the air from the bomb, the closed container was filled with 10 atm of oxygen through a screw on the head of the bomb. After being fully filled with oxygen, the bomb was completely closed and electrical leads were connected from the bomb insulator jacket to the electrical connection outlets on the head of the bomb. The bomb was then placed in a pail filled with 1750ml of water and this was placed in the insulated jacket. The lid of the jacket was replaced, with a Beckmann thermometer exposed from the lid used for measuring the temperature of the water in the pail. The bomb insulator was plugged into an electrical outlet, and a rubber band was connected from the motor to the stirrer, which was switched on. The temperature of the thermometer (indicating the temperature of the water in the pail) was left to become constant. The ignition controller was plugged into an electrical outlet and the button was pressed when the temperature of the water in the pail was deemed stable. The temperature values of the reaction, over a period of three minutes, were recorded every fifteen seconds. This exact procedure was performed using naphthalene, sucrose and an unknown sample as the sample. All results received from the temperature values were plotted using ExcelTM. Several trials were performed for each sample in order to obtain average and λ95 values.


The change in temperature values were obtained through this method:– 
1.Determine the minimum turning point on the curve at which the temperature 
extends upwards. 2.Determine the maximum point on the curve at which the temperature is stable
and not moving upwards. 3.Subtract the maximum from the minimum value. 4.Multiply this number by 0.60. 5.Add the number to the minimum value. 6.Extend a horizontal line from the value received in (5) to a point that
intersects with the curve. 7.Extend the lines of the top and bottom curves. 8.Rule a vertical line through the intersection point (6) to the extended maximum
and minimum lines. 9.Subtract the maximum from the minimum values obtained at the points where the top
and bottom curves (7) meet with vertical line (8). 10.The value received is the change in temperature (∆T).

For example:

Fig. 2. Determining ΔT for the combustion of 0.9033 g of benzoic acid and 0.01 g of iron wire
Finding the calorimeter heat capacity:
Mass of Iron burned: 0.01 g
U = 0.01g - 6.68 kJ/g = -0.0668 kJ
Benzoic Acid: 0.9033 g
U = 0.9033 g - 26.41 kJ/g= -23.86 kJ
ΔT = 23.85 - 21.30 0C = 2.55 0C	--> ∆U= -23.93 kJ
C = -23.93 kJ/2.55 0C = -9.38 kJ/0C
The average value of two trials gave -10.430 kJ/0C
λ95 = [t*s/(n)]1/2
= [2.91999*0.10/(2)]1/2 = 0.2065
Therefore, Cv for the calorimeter is -10.430 +/- 0.2065 kJ/0C

Finding the ΔU (kJ) and Energy (kJ/g) Values for Naphthalene, Sucrose and an Unknown Sample:
For naphthalene (Trial 1)
ΔU for the total reaction = -10.430kJ/0C * 2.300 0C = -23.990 kJ
ΔU for the fused wire = Energy (kJ/g) * Weight of fuse wire 
= -6.68 kJ/g * 0.0100 g = -0.0668 kJ
ΔU for naphthalene = ΔU for the total reaction - ΔU for the fused wire
= -23.990 kJ + 0.0668kJ = -23.936 kJ
The average of four trials gives -25.074 kJ 
λ95 = [t*s/(n)]1/2
= [2.13185*1.075/(4)]1/2 = 1.1459 
Therefore, the ΔU for naphthalene is -25.074 +/- 1.1459 kJ
Energy for naphthalene (kJ/g) = ΔU for naphthalene (kJ) / Weight of pellet used (g)
= -23.936 kJ / 0.6050 g = -39.563 kJ/g
The average of four trials gives -39.761 kJ/g
λ95 = [t * s / (n)]1/2
= [2.13185*1.322/(4)]1/2 = 1.4092 
Therefore, the energy value for naphthalene is -39.761 +/- 1.4092 kJ/g

Finding the ΔUm(kJ/mol), ∆Hm (kJ/mol) for the Benzoic Acid, Naphthalene and Sucrose Samples: 
For naphthalene (Trial 1)
ΔUm (kJ/mol) = Energy value (kJ/g)* Molecular Weight of sample (g/mol)
= -39.563 kJ/g * 128.2 g/mol = -5072.03 kJ/mol
Average ΔUm value for naphthalene = -5097.39 +/- 180.713 kJ/mol
For naphthalene (Trial 1)
ΔHm = ΔUm + RTΔn(gas)
= -5072.03 + (0.008314*298.23*-2) = -5076.99 kJ/mol
Average ΔHm value for naphthalene = -5102.35 +/- 180.714 kJ/mol

Determining the difference between ΔUm and ΔHm
R = 0.008314 kJ/(mol*K)
T = temperature at combustion found on graph (25.08 0C)
Δn(gas) = no. of moles for gas products - no. of moles for gas reactants
C10H8(s) + 12O2(g) ----> 10CO2(g) + 4H2O (l)
Δn(gas) = 10-12 = -2
Difference (ΔHm - ΔUm)
-5102.35 – (-5097.39) = -4.96 kJ/mol
Average difference is -4.961 +/- 0.0032 kJ/mol

Compound ΔT (0C) ΔU (kJ/g) ΔU (kJ) ΔU Total (kJ)Ψ ΔUm (kJ/mol) ΔHm (kj/mol) ΔHm - ΔUm (kJ/mol)
Benzoic acid 2.070 -26.410 -21.488 -21.541 -3224.66 -3225.90 -1.240
Naphthalene 2.410 -39.761 -25.074 -24.824 -5097.39 -5102.35 -4.961
Sucrose 1.100 -16.446 -11.523 -11.473 -5627.98 -5627.98 0.000
Unknown 0.947 -14.068 -9.945 -9.874 ~ ~ ~
Table 1. Summary of change in temperature, energy, internal energy and enthalpy results for the four samples used in the experiment.
ΨRefers to the combined combustion of the compound and the iron wire.


Upon calculation of the results, it has been observed that samples naphthalene and sucrose yielded different ΔU, ΔUm and ΔHm values. The unknown compound also yielded a low value ∆U compared the other samples. Although the molecular weights and make-up of the various compounds result in the varying ΔU, ΔUm and ΔHm values received, it is difficult to spot a correlation between them. Benzoic acid was used as a standard and gave rise to the heat capacity. This value is -10.430 +/- 0.2065 kJ/0C and exhibits the constant heat capacity of the bomb.
Analyzing the other three compounds revealed their ΔU values to be quite different. It seemed that the ΔU for naphthalene was closer to that of benzoic acid at -25.074 +/- 1.1459 kJ whilst those of sucrose and the unknown were -11.523 +/- 0.6184 kJ and -9.945 +/- 0.5579 kJ, respectively. These negative values indicate that energy was given off during the combustion of the organic sample into carbon dioxide and water. Unfortunately, more problems existed while performing the combustion of sucrose and the unknown as the pellet was not consumed for both reactions on two occasions.
In order to find the ΔUm and ΔHm values of naphthalene and sucrose the molecular weights of each compound were needed. The unknown sample could not be calculated, due to the lack of knowledge of its molecular weight. Calculating ΔHm requires knowledge of ΔUm. The ΔHm for naphthalene was -5102.35 +/- 180.714 kJ/mol and the literature value, -5156.95kJ/mol [4]. There existed between them a relative error percentage of 1.02%. Calculations of the ΔHm for sucrose yielded -5627.98 +/- 262.244 kJ/mol and the literature value was determined to be -5630.00 kJ/mol [4]. This exhibited a relative percentage error of 0.036% for the sucrose values.
All calculations demonstrate the highly exothermic property of these combustion reactions. They also show that the experiments performed were quite accurate in determining the given results as error between the results and the values obtained from literature was of a small percentage and does not seem significant.


[1] The American HeritageTM Dictionary of the English Language (4th Ed.), Houghton Mifflin Company, Boston, 2000
[2] Mashkevich S. V., Mashkevich V. S., “Statistical Theory of an Adiabatic Process”, Phys. Rev. E. 51: 245 – 253 (1995)
[3] MERCURY'S HELP DESK | CALORIMETRY, URL: (Retrieved April 01 2005)
[4] Heats of Combustion, URL: and URL: (Retrieved April 01 2005)
[5] Shoemaker, D., Experiments in Physical Chemistry, McGraw-Hill, Inc., 1989, p 211- 218
[6] Theory, Bomb Calorimetry, URL: (Retrieved April 01 2005)