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{{Image|Catalysis.png|right|290px|Figure 1: The catalysis cycle and its elementary steps.}} | {{Image|Catalysis.png|right|290px|Figure 1: The catalysis cycle and its elementary steps.}} | ||
'''Catalysis''' is a process that uses a substance to accelerate the rate of a [[chemical reaction]] through an uninterrupted and repeated cycle of elementary steps until the last step regenerates the catalyst in its original form. The substance that does this is known as a '''''catalyst'''''. It is usually present in relatively small amounts and none of it is consumed in the process.<ref name=NAP>{{cite book|author=Commission on Physical Sciences, Mathematics, and Applications (CPSMA), [[National Academies]]|publisher= National Academies Press|title=Catalysis Looks to the Future|edition=|year=1992|id=ISBN 0-309-04584-3}} Available online at [http://www.nap.edu/openbook.php?record_id=1903&page=1 Executive Summary]</ref> | In [[chemistry]], '''Catalysis''' is a process that uses a substance to accelerate the rate of a [[chemical reaction]] through an uninterrupted and repeated cycle of elementary steps until the last step regenerates the catalyst in its original form. The substance that does this is known as a '''''catalyst'''''. It is usually present in relatively small amounts and none of it is consumed in the process.<ref name=NAP>{{cite book|author=Commission on Physical Sciences, Mathematics, and Applications (CPSMA), [[National Academies]]|publisher= National Academies Press|title=Catalysis Looks to the Future|edition=|year=1992|id=ISBN 0-309-04584-3}} Available online at [http://www.nap.edu/openbook.php?record_id=1903&page=1 Executive Summary]</ref> | ||
Figure 1 depicts the steps in a typical catalysis cycle. As depicted, the reactant | Figure 1 depicts the steps in a typical catalysis cycle. As depicted, the reactant [[molecule]]s A and B are reacted to yield product P. The catalysis cycle starts with the bonding of reactant molecules A and B to the catalyst. A and B then react to yield product P which is also bound to the catalyst. In the last step, the catalyst is regenerated by product P separating from the catalyst. The regenerated catalyst then begins cycle again by bonding with two more reactant molecules. <ref name=Chork>{{cite book|author=I. Chorkendorff and J. W. Niemantsverdriet|title=Concepts of Modern Catalysts and Kinetics|edition=2nd Edition|publisher=Wiley-VCH|year=2007|id=ISBN 3-527-31672-8}}</ref> | ||
Many substances can act as catalysts, including: [[metal]]s, [[chemical compounds]] (e.g., metal [[oxide]]s, [[sulfide]]s, [[nitride]]s), [[organometallic]] complexes, and [[enzyme]]s. Although a catalyst may be a [[gas]], [[liquid]] or [[solid]], most catalysts used in industrial chemical reactions are in the form of porous pellets. Since not all parts of a solid catalyst participate in the catalysis cycle, those parts that do participate are referred to as ''active sites''. A single porous pellet may have 10<sup>18</sup> active catalytic sites.<ref name=NAP/> | Many substances can act as catalysts, including: [[metal]]s, [[chemical compounds]] (e.g., metal [[oxide]]s, [[sulfide]]s, [[nitride]]s), [[organometallic]] complexes, and [[enzyme]]s. Although a catalyst may be a [[gas]], [[liquid]] or [[solid]], most catalysts used in industrial chemical reactions are in the form of porous pellets. Since not all parts of a solid catalyst participate in the catalysis cycle, those parts that do participate are referred to as ''active sites''. A single porous pellet may have 10<sup>18</sup> active catalytic sites.<ref name=NAP/> | ||
==The catalysis mechanism== | ==The catalysis mechanism== | ||
{{Image|Catalysis reaction paths.png|right| | {{Image|Catalysis reaction paths.png|right|234px|Figure 2: Diagram showing the effect of a catalyst on a hypothetical chemical reaction. The catalyst provides a different reaction path with a lower activation energy than the uncatalyzed reaction path.}} | ||
In [[chemistry]], '''''activation energy'''''<ref>A term introduced in 1889 by the Swedish scientist [[Svante Arrhenius]]</ref> is the [Energy (science)|energy]] that must be overcome in order for a chemical reaction to occur. Activation energy may also be defined as the minimum energy required to start a designated chemical reaction. It is denoted by ''E<sub>a</sub>'' in units of kilojoules per mole (kJ/mol). It may be thought of as the ''energy barrier'' that must be overcome to start a chemical reaction. | |||
For a chemical reaction to proceed at a reasonable rate, there should exist an appreciable number of reactant species (molecules, [[atom]]s, [[ion]]s, etc.) with energy equal to or greater than the activation energy of the reaction.<ref name=Clark1>[http://www.chemguide.co.uk/physical/basicrates/catalyst.html#top The Effect of Catalysts on Reaction Rates] Website provided by Jim Clarke, retired Head of Chemistry and then Head of Science at [[Truro School]] in [[Cornwall]], [[United Kingdom]].</ref> A catalyst does not lower the activation energy for a reaction, instead it provides an alternative path for the reaction that has a lower activation energy. Figure 2 depicts how a catalyzed reaction follows a lower activation energy path than the higher activation energy path followed by the same reaction when it is not catalyzed. Overall, both the catalyzed path and the uncatalyzed path have the same change in Gibbs free energy between the reactants and the reaction product. | |||
==References== | ==References== | ||
{{reflist}} | {{reflist}} |
Revision as of 21:19, 6 April 2010
In chemistry, Catalysis is a process that uses a substance to accelerate the rate of a chemical reaction through an uninterrupted and repeated cycle of elementary steps until the last step regenerates the catalyst in its original form. The substance that does this is known as a catalyst. It is usually present in relatively small amounts and none of it is consumed in the process.[1]
Figure 1 depicts the steps in a typical catalysis cycle. As depicted, the reactant molecules A and B are reacted to yield product P. The catalysis cycle starts with the bonding of reactant molecules A and B to the catalyst. A and B then react to yield product P which is also bound to the catalyst. In the last step, the catalyst is regenerated by product P separating from the catalyst. The regenerated catalyst then begins cycle again by bonding with two more reactant molecules. [2]
Many substances can act as catalysts, including: metals, chemical compounds (e.g., metal oxides, sulfides, nitrides), organometallic complexes, and enzymes. Although a catalyst may be a gas, liquid or solid, most catalysts used in industrial chemical reactions are in the form of porous pellets. Since not all parts of a solid catalyst participate in the catalysis cycle, those parts that do participate are referred to as active sites. A single porous pellet may have 1018 active catalytic sites.[1]
The catalysis mechanism
In chemistry, activation energy[3] is the [Energy (science)|energy]] that must be overcome in order for a chemical reaction to occur. Activation energy may also be defined as the minimum energy required to start a designated chemical reaction. It is denoted by Ea in units of kilojoules per mole (kJ/mol). It may be thought of as the energy barrier that must be overcome to start a chemical reaction.
For a chemical reaction to proceed at a reasonable rate, there should exist an appreciable number of reactant species (molecules, atoms, ions, etc.) with energy equal to or greater than the activation energy of the reaction.[4] A catalyst does not lower the activation energy for a reaction, instead it provides an alternative path for the reaction that has a lower activation energy. Figure 2 depicts how a catalyzed reaction follows a lower activation energy path than the higher activation energy path followed by the same reaction when it is not catalyzed. Overall, both the catalyzed path and the uncatalyzed path have the same change in Gibbs free energy between the reactants and the reaction product.
References
- ↑ 1.0 1.1 Commission on Physical Sciences, Mathematics, and Applications (CPSMA), National Academies (1992). Catalysis Looks to the Future. National Academies Press. ISBN 0-309-04584-3. Available online at Executive Summary
- ↑ I. Chorkendorff and J. W. Niemantsverdriet (2007). Concepts of Modern Catalysts and Kinetics, 2nd Edition. Wiley-VCH. ISBN 3-527-31672-8.
- ↑ A term introduced in 1889 by the Swedish scientist Svante Arrhenius
- ↑ The Effect of Catalysts on Reaction Rates Website provided by Jim Clarke, retired Head of Chemistry and then Head of Science at Truro School in Cornwall, United Kingdom.