Heterogeneous Catalysis

HETEROGENEOUS CATALYSIS

D. C. AGUILAR AND B. N. SANCHEZ
INSTITUTE OF CHEMISTRY, COLLEGE OF SCIENCE
UNIVERSITY OF THE PHILIPPINES, DILIMAN QUEZON CITY, PHILIPPINES
RECEIVED MARCH 17, 2008

ABSTRACT

Catalysis which involves the use of a catalyst in a different phase from the reactants is known as heterogeneous catalysis. Catalysts are known to enhance rates of reaction without being consumed and they also reduce activation energies. The hydrogen peroxide decomposition reaction catalyzed in the presence of black, insoluble MnO2 solids has been investigated in this experiment. Parts 1A and 2A of the experiment are carried out at 25°C and 4°C, respectively. The same method and temperatures has been employed for the B part; only, the amount of KMnO4 added is twice the amount added in A. The calculated activation energy for B part is much lower than that of A part having a value of 24.91 kJ/mol. The calculated Arrhenius constant A, is also lower in B having a value of 34.84s-1. Results have shown that the higher the concentration of the catalyst used, the lower the activation energy, thus, the faster the reaction.

INTRODUCTION

A potential barrier called the activation energy must first be overcome by the reactants in order to become products. In the decomposition reaction of hydrogen peroxide, the barrier might be a state in which the oxygen-hydrogen bonds of one hydrogen peroxide molecule are being stretched as –OH bonds on an adjacent molecule are being formed [4]. This can be shown by the figure below.

The rate of the reaction depends on the magnitude of the activation barrier between the reactants. Here, the species with the highest energy is called the transition state. The Arrhenius equation k = Ae– Ea/RT, shows that the fraction of molecules having enough energy to overcome a potential energy barrier of magnitude Ea depends exponentially on the ratio of the activation to the thermal energy [4]. This equation also implies that the rate of the reaction exponentially increases with decreasing activation energy. In order to decrease Ea, a catalyst must be added to the system.

Catalysts have a remarkable property of facilitating a chemical reaction repeatedly without being consumed [1]. They provide a new reaction pathway, normally a lower activation barrier and enhanced rate [2]. Also, catalysts promote the production of a selected product [3]. Catalysts function by breaking the reaction into a sequence of steps. It does not appear in a balanced equation, but it does affect the rate of a reaction. Below is a reaction diagram for a catalytic reaction [4].

Generally, there are three types of catalysts- homogeneous, enzymes and heterogeneous catalysts. The latter is the one studied in this experiment.

Heterogeneous catalysis involves the interaction between a reactant and the catalytic surface of small metal particles. One or more of the reactants are adsorbed on to the surface of catalyst at active sites. This might involve an actual reaction with the surface, or some weakening of the bonds in the attached molecules. A good catalyst must adsorb strongly enough to hold and activate the reactants, but not so strongly that the products can’t break away.

DISCUSSIONS

The decomposition of hydrogen peroxide, known to be a first order reaction, was observed in this experiment. Here, the reaction was catalyzed in the presence of MnO2, a black solid that is insoluble in water. The colloidal catalyst is generated in situ by the addition of a small amount of KMnO4 to alkaline peroxide [lab]. Since the solid is insoluble, any enhancement of the rate of decomposition must occur through a heterogeneous pathway [4].

In part 1 A, a 0.5 mL of 0.02M KMnO4 was added to the reaction flask containing the peroxide, hence producing a small quantity of colloidal MnO2. This reaction was carried out at 25°C. At particular time intervals, samples were withdrawn from this reaction mixture in order to check the progress of peroxide decomposition reaction. Sulfuric acid was then added to stop the reaction and immediate titration of the solution against 0.002M KMnO4 was done so as to know the amount of remaining peroxide. The equation for the reaction is

2MnO4- + 5H202 + H+  2Mn2+ + 502 + 8 H20 .

On the assumption that the concentration of H202 was proportional to the titrant volume and that its decomposition as aforementioned, followed the first order rate law, Vt can be readily substituted for [H202] in the first-order equation. Thus, a plot of ln Vt against t, the time it takes for the titration reaction to be completed must yield an equation of a line whose negative of the slope equals the rate constant, k.

Part 1 B was very much the same as part 1 A. There was only a modification on the amount of KMnO4 added to the peroxide. Here, the original amount of KMnO4 added was doubled. A similar process was also applied in part 2 A, only, the reaction was carried out at 4°C instead of 25°C. In part 2 B, twice the amount of KMnO4 was added to initiate the reaction.

In heterogeneous catalysis, the reaction takes place at the solid or liquid interface. If an O2 gas layer forms on the solid, the H202 molecule will not be able to reach the active surface. Therefore, for a reaction to be possible, it must be ensured that molecules were continuously transported to the interfacial region [4]. For this reason, constant shaking of the reaction flask was necessary.

Plotting of ln Vt against t for each set of conditions will produce different rate constants. The rate constants for the two temperatures will then allow for the calculation of the activation energy, Ea and the pre-exponential factor, A, through manipulation of the Arrhenius equation. There is a significant difference in the values of Ea and A between the two catalyst concentrations. Part A reactions yield high values of A and Ea whereas in part B reactions, wherein the amount of KMnO4 added was doubled, a relatively lower activation energy and pre-exponential factor were obtained. Low activation energy means a fast reaction and high activation energy means a slow reaction. The rapid increase in k as T increases is due mainly to the increase in the number of collisions whose energy exceeds the activation energy [5]. Hence, from the results obtained, it can be inferred that the higher the concentration of catalyst to be used, the faster the reaction.

REFERENCES
[1]Chemistry: On Heterogeneous Catalysis. http://scienceweek.com/2004/sc041217-5.html
[2]A Colorful Catalysis Demonstration. http://chemeducator.org/sbibs/s0006004/spapers/640221tr.html
[3]Heterogeneous Catalysts. http://www.science.uwaterloo.ca/~cchieh/cact/applychem/heterocat.html
[4]Fast, Faster, Fastest: Exploring Chemical Catalysis. http://www.uiowa.edu/~c004020/exp5/exp.pdf
[5] Levine, Ira N. Physical Chemistry. 5th ed. McGraw-Hill, Inc. 2003
[6]Types of Catalysis. http://www.chemguide.co.uk/physical/catalysis/introduction.html

References: [1]Chemistry: On Heterogeneous Catalysis. http://scienceweek.com/2004/sc041217-5.html [2]A Colorful Catalysis Demonstration. http://chemeducator.org/sbibs/s0006004/spapers/640221tr.html [3]Heterogeneous Catalysts. http://www.science.uwaterloo.ca/~cchieh/cact/applychem/heterocat.html [4]Fast, Faster, Fastest: Exploring Chemical Catalysis. http://www.uiowa.edu/~c004020/exp5/exp.pdf [5] Levine, Ira N. Physical Chemistry. 5th ed. McGraw-Hill, Inc. 2003 [6]Types of Catalysis. http://www.chemguide.co.uk/physical/catalysis/introduction.html