Catalyst and Catalysis
A catalyst increases the rate of a particular reaction without itself being used up. A catalyst can be added to a reaction and then be recovered and reused after the reaction occurs. The process or action by which a catalyst increases the reaction rate is called catalysis. The study of reaction rates and how they change when manipulated experimentally is called kinetics.
The term catalysis was proposed in 1835 by the Swedish chemist Jöns Berzelius (1779-1848). The term comes from the Greek words kata meaning down and lyein meaning loosen. Berzelius explained that by the term catalysis he meant "the property of exerting on other bodies an action which is very different from chemical affinity. By means of this action, they produce decomposition in bodies, and form new compounds into the composition of which they do not enter."
Most chemical reactions occur as a series of steps. This series of steps is called a pathway or mechanism. Each individual step is called an elementary step. The slowest elementary step in a pathway determines the reaction rate. The reaction rate is the rate at which reactants disappear and products appear in a chemical reaction, or, more specifically, the change in concentration of reactants and products in a certain amount of time.
While going through a reaction pathway, reactants enter a transitional state where they are no longer reactants, but are not yet products. During this transitional state they form what is called an activated complex. The activated complex is short-lived and has partial bonding characteristics of both reactants and products. The energy required to reach this transitional state and form the activated complex in a reaction is called the activation energy. In order for a reaction to occur, the activation energy must be reached. A catalyst increases the rate of reaction by lowering the activation energy required for the reaction to take place. The catalyst forms an activated complex with a lower energy than the complex formed without catalysis. This provides the reactants a new pathway which requires less energy. Although the catalyst lowers the activation energy required, it does not affect reaction equilibrium or thermodynamics. The catalyst does not appear in the overall chemical equation for a pathway because the mechanism involves an elementary step in which the catalyst is consumed and another in which it is regenerated.
Catalysts exist for all types of chemical reactions. A specific catalyst can be classified into one of two main groups; homogeneous and heterogeneous. A catalyst that is in the same phase as the reactants and products involved in a reaction pathway is called a homogeneous catalyst. When a catalyst exists in a different phase than that of the reactants, it is called a heterogeneous catalyst. For example, nickel is a catalyst in the hydrogenation of vegetable oils. Nickel is a solid, while the oil is a liquid, therefore nickel is a heterogeneous catalyst. An advantage of using heterogeneous catalysts is their ease of separation from the reactants and products involved in a pathway.Metals are often used as heterogeneous catalysts because many reactants adsorb to the metal surface, increasing the concentration of the reactants and therefore the rate of the reaction. Ionic interactions between metals and other molecules can be used to orient the reactants involved so that they react better with each other, or to stabilize charged reaction transition states. Metals also can increase the rate of oxidation-reduction reactions through changes in the metal ion's oxidation state.
Another group of catalysts are called enzymes. Enzymes are catalysts that are found in biological systems. The role of catalysts in living systems was first recognized in 1833. French chemists Anselme Payen (1795-1871) and Jean François Persoz isolated a material from malt that accelerated the conversion of starch to sugar. Payen called the substance diatase. A half century later German physiologist Willy Kühne suggested the name enzyme for biological catalysts.
Enzymes are proteins and therefore have a highly folded three-dimensional configuration. This configuration makes an enzyme particularly specific for a certain reaction or type of reaction. Synthetic catalysts, on the other hand, are not nearly as specific. They will catalyze similar reactions that involve a wide variety of reactants. Enzymes, in general, will lose activity more easily than synthetic catalysts. Very slight disturbances in the protein structure of enzymes will change the three-dimensional configuration of the molecule and, as a result, its reactivity. Enzymes tend to be more active, i.e., they catalyze reactions faster, than synthetic catalysts at ambient temperatures. Catalytic activity for a reaction is expressed as the turnover number. This is simply the number of reactant molecules changed to product per catalyst site in a given unit of time. When temperature is increased, synthetic catalysts can become just as active as enzymes. With an increase in temperature, many enzymes will become inactive because of changes to the protein structure.
There are endless reactions that can undergo catalysis. One example is the decomposition of hydrogen peroxide (H2O2). Without catalysis, hydrogen peroxide decomposes slowly over time to form water and oxygen gas. A 30% solution of hydrogen peroxide at room temperature will decompose at a rate of 0.5% per year. The activation energy for this reaction is 75 kJ/mol. This activation energy can be lowered to 58 kJ/mol with the addition of iodide ions (I-). These ions form an intermediate, HIO-, which reacts with the hydrogen peroxide to regenerate the iodide ions. When the enzyme catalase is added to the hydrogen peroxide solution, the activation energy is lowered even further to 4 kJ/mol. The catalase is also regenerated in the reaction and can be separated from the solution for reuse. This example shows how a catalyst can lower the activation energy of a reaction without itself being used up in the reaction pathway.
Another example of catalysis is the catalytic converter of an automobile. Exhaust from the automobile can contain carbon monoxide and nitrogen oxides, which are poisonous gases. Before the exhaust can leave the exhaust system these toxins must be removed. The catalytic converter mixes these gases with air and then passes them over a catalyst made of rhodium and platinum metals. This catalyst accelerates the reaction of carbon monoxide with oxygen and converts it to carbon dioxide, which is not toxic. The catalyst also increases the rate of reactions for which the nitrogen oxides are broken down into their elements.
A well-known example of catalysis is the destruction of the ozone layer. Ozone (O3) in the upper atmosphere serves as a shield for the harmful ultraviolet rays from the Sun. Ozone is formed when an oxygen molecule (O2) is split into two oxygen atoms (O) by the radiation from the Sun. The free oxygen atoms then attach to oxygen molecules to form ozone. When another free oxygen atom reacts with the ozone molecule, two oxygen molecules are formed. This is the natural destruction of ozone. Under normal circumstances, the rate of destruction of ozone is the same as the rate of ozone formation, so no net ozone depletion occurs. When chlorine (Cl) atoms are present in the atmosphere, they act as catalysts for the destruction of ozone. Chlorine atoms in the atmosphere come from compounds containing chlorofluorocarbons, or CFCs. CFCs are compounds containing chlorine, fluorine, and carbon. CFCs are very stable and can drift into the upper atmosphere without first being broken down. Once in the upper atmosphere, the energy from the Sun causes the chlorine to be released. The chlorine atom reacts with ozone to form chlorine monoxide (ClO) and an oxygen molecule. The chlorine monoxide then reacts with another oxygen atom to form an oxygen molecule and the regenerated chlorine atom. With the help of the chlorine catalyst, the degeneration of ozone occurs at a faster rate than its formation, which has caused a net depletion of ozone in the atmosphere.
The previous examples illustrate some of the many practical applications of catalysis. Almost all of the chemicals produced by the chemical industry are made using catalysis. Catalytic processes used in the chemical industry decrease production costs as well as create products with higher purity and less environmental hazards. A wide variety of products are made using catalytic processes. Catalysis is used in industrial chemistry, pharmaceutical chemistry, and agricultural chemistry, as well as in the specialty chemical industry. Useful chemicals such as sulfuric acid, penicillin, and fructose are made more efficiently using catalytic processes. Research and development efforts in the chemical industry are significantly more productive with the use of catalysis in fields such as fuel refining, petrochemical manufacturing, and environmental management.
The majority of manufacturing processes in use today by the chemical industry employ catalytic reactions. These reactions are highly efficient, but research is continuing to increase the efficiency even more. The focus of this research is on separation and regeneration of the catalysts in order to decrease costs of production while increasing the purity of the product. The field of catalysis research is rapidly growing and will continue to do so as new catalysts and catalytic processes are discovered.
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