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Solid heterogeneous catalysts such as in automobile catalytic converters are plated on structures designed to maximize their surface area.

Catalysis is the process in which the rate of a chemical reaction is changed by a substance known as a catalyst. Unlike other reagents that participate in the chemical reaction, a catalyst is not consumed by the reaction itself. The catalyst may participate in multiple chemical transformations. Catalysts that speed the reaction are called positive catalysts. Catalysts that slow down the reaction are called negative catalysts or inhibitors. Substances that increase the activity of catalysts are called promoters and substances that deactivate catalysts are called catalytic poisons.

The general feature of catalysis is that the catalytic reaction has a lower rate-limiting free energy change to the transition state than the corresponding uncatalyzed reaction, resulting in a larger reaction rate at the same temperature. However, the mechanistic origin of catalysis is complex. Catalysts may affect the reaction environment favorably, e.g. acid catalysts for reactions of carbonyl compounds form specific intermediates that are not produced naturally, such as osmate esters in osmium tetroxide-catalyzed dihydroxylation of alkenes, or cause lysis of reagents to reactive forms, such as atomic hydrogen in catalytic hydrogenation.

Kinetically, catalytic reactions behave like typical chemical reactions, i.e. the reaction rate depends on the frequency of contact of the reactants in the rate-determining step. Usually, the catalyst participates in this slow step, and rates are limited by amount of catalyst. In heterogeneous catalysis, the diffusion of reagents to the surface and diffusion of products from the surface can be rate determining. Analogous events associated with substrate binding and product dissociation apply to homogeneous catalysts.

Although catalysts are not consumed by the reaction itself, they may be inhibited, deactivated or destroyed by secondary processes. In heterogeneous catalysis, typical secondary processes include coking where the catalyst becomes covered by polymeric side products. Additionally, heterogeneous catalysts can dissolve into the solution in a solid-liquid system or evaporate in a solid-gas system.

Contents

Background

The production of most industrially important chemicals involves catalysis. Similarly, most biochemically significant processes are catalysed. Research into catalysis is a major field in applied science and involves many areas of chemistry, notably in organometallic chemistry and materials science. Catalysis is relevant to many aspects of environmental science, e.g. the catalytic converter in automobiles and the dynamics of the ozone hole. Catalytic reactions are preferred in environmentally friendly green chemistry due to the reduced amount of waste generated,[1] as opposed to stoichiometric reactions in which all reactants are consumed and more side products are formed. The most common catalyst is the proton (H+). Many transition metals and transition metal complexes are used in catalysis as well. Catalysts called enzymes are important in biology.

A catalyst works by providing an alternative reaction pathway to the reaction product. The rate of the reaction is increased as this alternative route has a lower activation energy than the reaction route not mediated by the catalyst. The disproportionation of hydrogen peroxide to give water and oxygen is a reaction that is strongly affected by catalysts:

2 H2O2 → 2 H2O + O2

This reaction is favoured in the sense that reaction products are more stable than the starting material, however the uncatalysed reaction is slow. The decomposition of hydrogen peroxide is in fact so slow that hydrogen peroxide solutions are commercially available. Upon the addition of a small amount of manganese dioxide, the hydrogen peroxide rapidly reacts according to the above equation. This effect is readily seen by the effervescence of oxygen.[2] The manganese dioxide may be recovered unchanged, and re-used indefinitely, and thus is not consumed in the reaction. Accordingly, manganese dioxide catalyses this reaction.[3]

General principles of catalysis

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Typical mechanism

Catalysts generally react with one or more reactants to form intermediates that subsequently give the final reaction product, in the process regenerating the catalyst. The following is a typical reaction scheme, where C represents the catalyst, X and Y are reactants, and Z is the product of the reaction of X and Y:

X + C → XC (1)
Y + XC → XYC (2)
XYCCZ (3)
CZ → C + Z (4)

Although the catalyst is consumed by reaction 1, it is subsequently produced by reaction 4, so for the overall reaction:

X + Y → Z

As a catalyst is regenerated in a reaction, often only small amounts are needed to increase the rate of the reaction. In practice, however, catalysts are sometimes consumed in secondary processes.

As an example of this process, in 2008 Danish researchers first revealed the sequence of events when oxygen and hydrogen combine on the surface of titanium dioxide (TiO2, or titania) to produce water. With a time-lapse series of scanning tunneling microscopy images, they determined the molecules undergo adsorption, dissociation and diffusion before reacting. The intermediate reaction states were: HO2, H2O2, then H3O2 and the final reaction product (water molecule dimers), after which the water molecule desorbs from the catalyst surface.[4]

Catalysis and reaction energetics

Generic potential energy diagram showing the effect of a catalyst in a hypothetical exothermic chemical reaction X + Y to give Z. The presence of the catalyst opens a different reaction pathway (shown in red) with a lower activation energy. The final result and the overall thermodynamics are the same.

Catalysts work by providing an (alternative) mechanism involving a different transition state and lower activation energy. Consequently, more molecular collisions have the energy needed to reach the transition state. Hence, catalysts can enable reactions that would otherwise be blocked or slowed by a kinetic barrier. The catalyst may increase reaction rate or selectivity, or enable the reaction at lower temperatures. This effect can be illustrated with a Boltzmann distribution and energy profile diagram.

Catalysts do not change the extent of a reaction: they have no effect on the chemical equilibrium of a reaction because the rate of both the forward and the reverse reaction are both affected (see also thermodynamics). The fact that a catalyst does not change the equilibrium is a consequence of the second law of thermodynamics. Suppose there was such a catalyst that shifted an equilibrium. Introducing the catalyst to the system would result in reaction to move to the new equilibrium, producing energy. Production of energy is a necessary result since reactions are spontaneous if and only if Gibbs free energy is produced, and if there is no energy barrier, there is no need for a catalyst. Then, removing the catalyst would also result in reaction, producing energy; i.e. the addition and its reverse process, removal, would both produce energy. Thus, a catalyst that could change the equilibrium would be a perpetual motion machine, a contradiction to the laws of thermodynamics.[5]

If a catalyst does change the equilibrium, then it must be consumed as the reaction proceeds, and thus it is also a reactant. Illustrative is the base-catalysed hydrolysis of esters, where the produced carboxylic acid immediately reacts with the base catalyst and thus the reaction equilibrium is shifted towards hydrolysis.

The SI derived unit for measuring the catalytic activity of a catalyst is the katal, which is moles per second. The activity of a catalyst can also be described by the turn over number (or TON) and the catalytic efficiency by the turn over frequency (TOF). The biochemical equivalent is the enzyme unit. For more information on the efficiency of enzymatic catalysis, see the article on Enzymes.

The catalyst stabilizes the transition state more than it stabilizes the starting material. It decreases the kinetic barrier by decreasing the difference in energy between starting material and transition state.

Typical catalytic materials

The chemical nature of catalysts is as diverse as catalysis itself, although some generalizations can be made. Proton acids are probably the most widely used catalysts, especially for the many reactions involving water, including hydrolysis and its reverse. Multifunctional solids often are catalytically active, e.g. zeolites, alumina and certain forms of graphitic carbon. Transition metals are often used to catalyse redox reactions (oxidation, hydrogenation). Many catalytic processes, especially those involving hydrogen, require platinum metals.

Some so-called catalysts are really precatalysts. Precatalysts convert to catalysts in the reaction. For example, Wilkinson's catalyst RhCl(PPh3)3 loses one triphenylphosphine ligand before entering the true catalytic cycle. Precatalysts are easier to store but are easily activated in situ. Because of this preactivation step, many catalytic reactions involve an induction period.

Chemical species that improve catalytic activity are called co-catalysts (cocatalysts) or promotors in cooperative catalysis.

Types of catalysis

Catalysts can be either heterogeneous or homogeneous, depending on whether a catalyst exists in the same phase as the substrate. Biocatalysts are often seen as a separate group.

Heterogeneous catalysts

Heterogeneous catalysts are those which act in a different phases than the reactants. Most heterogeneous catalysts are solids that act on substrates in a liquid or gaseous reaction mixture. Diverse mechanisms for reactions on surfaces are known, depending on how the adsorption takes place (Langmuir-Hinshelwood, Eley-Rideal, and Mars-van Krevelen).[6] The total surface area of solid has an important effect on the reaction rate. The smaller the catalyst particle size, the larger the surface area for a given mass of particles.

For example, in the Haber process, finely divided iron serves as a catalyst for the synthesis of ammonia from nitrogen and hydrogen. The reacting gases adsorb onto "active sites" on the iron particles. Once adsorbed, the bonds within the reacting molecules are weakened, and new bonds between the resulting fragments form in part due to their close proximity. In this way the particularly strong triple bond in nitrogen is weakened and the hydrogen and nitrogen atoms combine faster than would be the case in the gas phase, so the rate of reaction increases.[citation needed]

Heterogeneous catalysts are typically “supported,” which means that the catalyst is dispersed on a second material that enhances the effectiveness or minimizes their cost. Sometimes the support is merely a surface upon which the catalyst is spread to increase the surface area. More often, the support and the catalyst interact, affecting the catalytic reaction.

Homogeneous catalysts

Homogeneous catalysts function in the same phase as the reactants, but the mechanistic principles invoked in heterogeneous catalysis are generally applicable. Typically homogeneous catalysts are dissolved in a solvent with the substrates. One example of homogeneous catalysis involves the influence of H+ on the esterification of esters, e.g. methyl acetate from acetic acid and methanol.[7] For inorganic chemists, homogeneous catalysis is often synonymous with organometallic catalysts.[8]

Electrocatalysts

In the context of electrochemistry, specifically in fuel cell engineering, various metal-containing catalysts are used to enhance the rates of the half reactions that comprise the fuel cell. One common type of fuel cell electrocatalyst is based upon nanoparticles of platinum that are supported on slightly larger carbon particles. When this platinum electrocatalyst is in contact with one of the electrodes in a fuel cell, it increases the rate of oxygen reduction to water (or hydroxide or hydrogen peroxide).

Organocatalysis

Whereas transition metals sometimes attract most of the attention in the study of catalysis, organic molecules without metals can also possess catalytic properties. Typically, organic catalysts require a higher loading (or amount of catalyst per unit amount of reactant) than transition metal-based catalysts, but these catalysts are usually commercially available in bulk, helping to reduce costs. In the early 2000s, organocatalysts were considered "new generation" and are competitive to traditional metal-containing catalysts. Enzymatic reactions operate via the principles of organic catalysis.

Significance of catalysis

Estimates are that 90% of all commercially produced chemical products involve catalysts at some stage in the process of their manufacture.[9] In 2005, catalytic processes generated about $900 billion in products worldwide.(pdf) Catalysis is so pervasive that subareas are not readily classified. Some areas of particular concentration are surveyed below.

Energy processing

Petroleum refining makes intensive use of catalysis for alkylation, catalytic cracking (breaking long-chain hydrocarbons into smaller pieces), naphtha reforming and steam reforming (conversion of hydrocarbons into synthesis gas). Even the exhaust from the burning of fossil fuels is treated via catalysis: Catalytic converters, typically composed of platinum and rhodium, break down some of the more harmful byproducts of automobile exhaust.

2 CO + 2 NO → 2 CO2 + N2

With regard to synthetic fuels, an old but still important process is the Fischer-Tropsch synthesis of hydrocarbons from synthesis gas, which itself is processed via water-gas shift reactions, catalysed by iron. Biodiesel and related biofuels require processing via both inorganic and biocatalysts.

Fuel cells rely on catalysts for both the anodic and cathodic reactions.

Bulk chemicals

Some of the largest-scale chemicals are produced via catalytic oxidation, often using oxygen. Examples include nitric acid (from ammonia), sulfuric acid (from sulfur dioxide to sulfur trioxide by the chamber process), terephthalic acid from p-xylene, and acrylonitrile from propane and ammonia.

Many other chemical products are generated by large-scale reduction, often via hydrogenation. The largest-scale example is ammonia, which is prepared via the Haber process from nitrogen. Methanol is prepared from carbon monoxide.

Bulk polymers derived from ethylene and propylene are often prepared via Ziegler-Natta catalysis. Polyesters, polyamides, and isocyanates are derived via acid-base catalysis.

Most carbonylation processes require metal catalysts, examples include the Monsanto acetic acid process and hydroformylation.

Fine chemicals

Many fine chemicals are prepared via catalysis; methods include those of heavy industry as well as more specialized processes that would be prohibitively expensive on a large scale. Examples include olefin metathesis using Grubbs' catalyst, the Heck reaction, and Friedel-Crafts reactions.

Because most bioactive compounds are chiral, many pharmaceuticals are produced by enantioselective catalysis.

Food processing

One of the most obvious applications of catalysis is the hydrogenation (reaction with hydrogen gas) of fats using nickel catalyst to produce margarine.[10] Many other foodstuffs are prepared via biocatalysis (see below).

Biology

In nature, enzymes are catalysts in metabolism and catabolism. Most biocatalysts are protein-based, i.e. enzymes, but other classes of biomolecules also exhibit catalytic properties including ribozymes, and synthetic deoxyribozymes.[11]

Biocatalysts can be thought of as intermediate between homogenous and heterogeneous catalysts, although strictly speaking soluble enzymes are homogeneous catalysts and membrane-bound enzymes are heterogeneous. Several factors affect the activity of enzymes (and other catalysts) including temperature, pH, concentration of enzyme, substrate, and products. A particularly important reagent in enzymatic reactions is water, which is the product of many bond-forming reactions and a reactant in many bond-breaking processes.

Enzymes are employed to prepare many commodity chemicals including high-fructose corn syrup and acrylamide.

In the environment

Catalysis impacts the environment by increasing the efficiency of industrial processes, but catalysis also plays a direct role in the environment. A notable example is the catalytic role of Chlorine free radicals in the breakdown of ozone. These radicals are formed by the action of ultraviolet radiation on chlorofluorocarbons (CFCs).

Cl· + O3 → ClO· + O2
ClO· + O· → Cl· + O2

History

In a general sense, anything that increases the rate of a process is a "catalyst", a term derived from Greek καταλύειν, meaning "to annul," or "to untie," or "to pick up." The phrase catalysed processes was coined by Jöns Jakob Berzelius in 1836[12] to describe reactions that are accelerated by substances that remain unchanged after the reaction. Other early chemists involved in catalysis were Alexander Mitscherlich who referred to contact processes and Johann Wolfgang Döbereiner who spoke of contact action and whose lighter based on hydrogen and a platinum sponge became a huge commercial success in the 1820s. Humphry Davy discovered the use of platinum in catalysis. In the 1880s, Wilhelm Ostwald at Leipzig University started a systematic investigation into reactions that were catalyzed by the presence of acids and bases, and found that chemical reactions occur at finite rates and that these rates can be used to determine the strengths of acids and bases. For this work, Ostwald was awarded the 1909 Nobel Prize in Chemistry.[13]

Inhibitors, poisons and promoters

Substances that reduce the action of catalysts are called catalyst inhibitors if reversible, and catalyst poisons if irreversible. Promoters are substances that increase the catalytic activity, particularly when not being catalysts unto themselves.

The inhibitor may modify selectivity in addition to rate. For instance, in the reduction of ethyne to ethene, the catalyst is palladium (Pd) partly "poisoned" with lead(II) acetate (Pb(CH3COO)2). Without the deactivation of the catalyst, the ethene produced will be further reduced to ethane.[14][15]

The inhibitor can produce this effect by e.g. selectively poisoning only certain types of active sites. Another mechanism is the modification of surface geometry. For instance, in hydrogenation operations, large planes of metal surface function as sites of hydrogenolysis catalysis while sites catalyzing hydrogenation of unsaturates are smaller. Thus, a poison that covers surface randomly will tend to reduce the number of uncontaminated large planes but leave proportionally more smaller sites free, thus changing the hydrogenation vs. hydrogenolysis selectivity. Many other mechanisms are also possible.

Promoters can cover up surface to prevent production of a mat of coke, or even actively remove such material (e.g. rhenium on platinum in platforming). They can aid the dispersion of the catalytic material or bind to reagents.

See also

References

  1. ^ "The 12 Principles of Green Chemistry". United States Environmental Protection Agency. http://www.epa.gov/greenchemistry/pubs/principles.html. Retrieved 2006-07-31. 
  2. ^ "Genie in a Bottle". University of Minnesota. 2005-03-02. http://www.chem.umn.edu/services/lecturedemo/info/genie.htm. 
  3. ^ Masel, Richard I. “Chemical Kinetics and Catalysis” Wiley-Interscience, New York, 2001. ISBN 0471241970.
  4. ^ Chemical & Engineering News, 16 February 2009, "Making Water Step by Step", p. 10
  5. ^ Robertson, A.J.B. Catalysis of Gas Reactions by Metals. Logos Press, London, 1970.
  6. ^ Helmut Knözinger, Karl Kochloefl “Heterogeneous Catalysis and Solid Catalysts” in Ullmann's Encyclopedia of Industrial Chemistry 2002, Wiley-VCH, Weinheim. doi:10.1002/14356007.a05_313. Article Online Posting Date: January 15, 2003
  7. ^ Arno Behr “Organometallic Compounds and Homogeneous Catalysis” Ullmann's Encyclopedia of Industrial Chemistry, 2002, Wiley-VCH, Weinheim. doi:10.1002/14356007.a18_215. Article Online Posting Date: June 15, 2000
  8. ^ Elschenbroich, C. ”Organometallics” (2006) Wiley-VCH: Weinheim. ISBN 978-3-29390-6
  9. ^ "Recognizing the Best in Innovation: Breakthrough Catalyst". R&D Magazine, September 2005, pg 20.
  10. ^ "Types of catalysis". Chemguide. http://www.chemguide.co.uk/physical/catalysis/introduction.html. Retrieved 2008-07-09. 
  11. ^ Nelson, D. L.; Cox, M. M. "Lehninger, Principles of Biochemistry" 3rd Ed. Worth Publishing: New York, 2000. ISBN 1-57259-153-6.
  12. ^ K.J. Laidler and J.H. Meiser, Physical Chemistry, Benjamin/Cummings (1982), p.423
  13. ^ M.W. Roberts (2000). "Birth of the catalytic concept (1800-1900)". Catalysis Letters 67 (1): 1–4. doi:10.1023/A:1016622806065. http://www.springerlink.com/content/qm3732u7x7577224/fulltext.pdf. 
  14. ^ W.P. Jencks, “Catalysis in Chemistry and Enzymology” McGraw-Hill, New York, 1969. ISBN 0070323054
  15. ^ Myron L Bender, Makoto Komiyama, Raymond J Bergeron “The Bioorganic Chemistry of Enzymatic Catalysis” Wiley-Interscience, Hoboken, U.S., 1984 ISBN 0471059919

External links


1911 encyclopedia

Up to date as of January 14, 2010

From LoveToKnow 1911

CATALYSIS (from the Gr. icara, down, and Ma y, to loosen), in chemistry, the name given to chemical actions brought about by a substance, termed the "catalyst," which is recovered unchanged after the action. The term was introduced by Berzelius, who first studied such reactions. It is convenient to divide catalytic actions into two groups: - (1) when the catalyst first combines with one of the reaction components to form a compound which immediately reacts with the other components, the catalyst being simultaneously liberated, and free to react with more of the undecomposed first component; and (2), when the catalyst apparently reacts by mere contact. The theory of catalysis is treated under Chemical Action; In This Article Mention Will Be Made Of Some Of The More Interesting Examples.

01 A Familiar Instance Of A Catalytic Action Is Witnessed When A Mixture Of Potassium Chlorate And Manganese Dioxide Is Heated To 350°, Oxygen Being Steadily Liberated, And The Manganese Dioxide Being Unchanged At The End Of The Reaction. The Action May Be Explained As Follows: Part Of The Chlorate Reacts With The Manganese Dioxide To Form Potassium Permanganate, Chlorine And Oxygen, The Chlorine Subsequently Reacting With The Permanganate To Produce Manganese Dioxide, Potassium Chloride And Oxygen, Thus 2Kc10312Mn0 2 = 2Kmn0 4 C1210 2 = 2Kc1 { 2Mn02 302.

This Explanation Is Supported By The Facts That Traces Of Chlorine Are Present In The Gas, And The Pink Permanganate Can Be Recognized When Little Dioxide Is Used. Other Oxides Bring About The Same Decomposition At Temperatures Below That At Which The Chlorate Yields Oxygen When Heated Alone; But Since Such Substances As Kaolin, Platinum Black And Some Other Finely Powdered Compounds Exercise The Same Effect, It Follows That The Explanation Given Above Is Not Quite General. Another Example Is Deacon'S Process For The Manufacture Of Chlorine By Passing Hydrochloric Acid Gas Mixed With Air Over Heated Bricks Which Had Been Previously Impregnated With A Copper Sulphate Solution. The Nitrous Gases Employed In The Ordinary Chamber Process Of Manufacturing Sulphuric Acid Also Act Catalytically. Mention May Be Made Of The Part Played By Water Vapour In Conditioning Many Chemical Reactions. Thus Sodium Will Not React With Dry Chlorine Or Dry Oxygen; Carbon, Sulphur And Phosphorus Will Not Burn In Perfectly Dry Oxygen, Neither Does Nitric Oxide Give Red Fumes Of The Peroxide. In Organic Chemistry Many Catalytic Actions Are Met With. In The Class Of Reaction Known As "Condensations," It May Be Found That The Course Of The Reaction Is Largely Dependent Upon The Nature Of Some Substance Which Acts Catalytically. One Of The Most Important Is The Friedel And Craft'S Reaction, In Which An Aromatic Compound Combines With An Alkyl Haloid In The Presence Of Aluminium, Zinc Or Ferric Chloride. It Seems In This, As In Other Cases, That Addition Compounds Are First Formed Which Subsequently React With The Re Formation Of The Catalyst. The Formation Of Benzoin From Benzaldehyde In The Presence Of Potassium Cyanide Is Another Example; This Action Has Been Investigated By G. Bredig And Stern (Zeit. Elektrochem., 1904, 10, P. 582).

The Second Class Of Catalytic Actions, Viz. Those Occasioned By The Presence Of A Metal Or Some Other Substance Which Undergoes No Change, Is Of Especial Interest, And Has Received Much Attention. The Accelerating Influence Of A Clean Platinum Plate On The Rate Of Combination Of Hydrogen And Oxygen Was Studied By Faraday. He Found That With The Pure Gases The Velocity Of Reaction Increased Until The Mixture Exploded. The Presence Of Minute Quantities Of Carbon Monoxide, Carbon Disulphide, Sulphuretted Hydrogen And Hydrochloric Acid Inhibited The Action; In The Case Of The First Two Gases, There Is No Alteration Of The Platinum Surface, Since The Plate Brings About Combination When Removed To An Atmosphere Of Pure Hydrogen And Oxygen; With The Last Two Gases, However, The Surface Is Altered, Since The Plate Will Not Occasion The Combination When Placed In The Pure Gases. M. Bodenstein (Zeit. Phys. Chem., 1904, 4 6, P. 725) Showed That Combination Occurs With Measurable Velocity At Ordinary Temperatures In The Presence Of Compact Platinum. More Energetic Combination Is Observed If The Metal Be Finely Divided, As, For Instance, By Immersing Asbestos Fibres In A Solution Of Platinum Chloride And Strongly Heating. The "Spongy" Platinum So Formed Brings About The Combination Of Ammonia And Oxygen To Form Water And Nitric Acid, Of Nitric Oxide And Hydrogen To Form Ammonia (See German Patent, 1905, 157,287), And Of Sulphur Dioxide And Oxygen To Form Sulphur Trioxide. The Last Reaction, Which Receives Commercial Application In The Contact Process Of Sulphuric Acid Manufacture, Was Studied By M. Bodenstein And W. Pohl (Zeit. Elektrochem., 1905, 1R, P. 373), Who Found That The Equilibrium Followed The Law Of Mass Action (See Also F. W. Kiister, Zeit. Anorg. Chen., 1904, 4 2, P. 453, R. Lucas, Zeit. Elektrochem., 1905, Ii, P. 457). Other Metals, Such As Nickel, Iron, &C., Can Also React As Catalysts.

The Use Of Finely Divided Nickel (Obtained By Reducing The Oxide In A Current Of Pure Hydrogen At A Temperature Of 350°) Has Been Carefully Studied By P. Sabatier And J. B. Senderens; A Summary Of Their Results Is Given In The Ann. China. Phys., 1905 (Viii.) 4, Pp. 3194 88. Of Special Interest Is The Condensation Of Acetylene. If This Gas Mixed With Hydrogen Be Passed Over The Reduced Nickel In The Cold, The Temperature May Rise To As High As 150°, The Acetylene Disappearing And Becoming Replaced By A Substance Like Petroleum. If The Nickel Be Maintained At 200°, And The Gases Circulated For Twenty Eight Hours, A Product, Condensible To A Yellow Liquid Having A Beautiful Fluorescence And Boiling At 45°, Is Obtained. This Substance Closely Resembles Ordinary Pennsylvanian Petroleum. If Acetylene Be Passed Alone Over Nickel Heated To 200° 300°, A Mixture, Boiling At 60° 70 And Having A Green Colour By Diffused And A Red By Transmitted Light, Was Obtained. This Substance Closely Resembles Caucasian Petroleum. The Decomposition Of Carbon Monoxide According To The Reaction 2Co< =2Cd Co 2 Is Purely Catalytic In The Presence Of Nickel And Cobalt, And Also In The Presence Of Iron, So Long As The Amount Of Carbon Dioxide Present Does Not Exceed A Certain Amount (R. Schenck And W. Heller, Ber., 1905, 38, Pp. 2132, 2139). It Is Of Interest That Finely Divided Aluminium And Magnesium Decompose Methane, Ethane, And Ethylene Into Carbon And Hydrogen In The Same Way As Nickel. Charcoal At 350° Also Reacts Catalytically; For Example, Senderens Found That Ethyl Alcohol Was Decomposed By Animal Charcoal Into Methane, Ethylene, Hydrogen, Carbon Monoxide And A Little Carbon Dioxide, And Propyl Alcohol Gave Propylene, Ethane, Carbon Monoxide And Hydrogen, While G. Lemoine Obtained From Ethyl Alcohol And Wood Charcoal A Mixture Of Acetaldehyde And Hydrogen.


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Wikibooks

Up to date as of January 23, 2010
(Redirected to Structural Biochemistry/Catalysis article)

From Wikibooks, the open-content textbooks collection

< Structural Biochemistry

One of the most common functions of enzymes is the ability to catalyze reactions. During a reaction, the reactants must overcome an activation energy in order for it to produce the products. The amount of activation energy needed determines how long the reaction takes to proceed. The lower the activation energy, the faster the rate of the reaction. The role of enzymes in catalyzing reactions is to stabilize the intermediate species, which is at the highest point of the activation energy, and thus dropping the activation energy. The enzyme is complementary not to the substrate but its intermediate state. If the enzyme binds to the substrate, it actually increases the activation energy. The equilibrium achieved is the same with or without the catalytic enzyme. However, what is affected is the time and rate in which it is achieved.

Generally, the higher the concentration of the substrate, the easier it is for the enzyme to bind to it. By plotting the amount of product produced as a function of time, the slope is how fast the reaction happens before the amount of substrate is saturated. This value is called the V0. Increasing the substrate concentration will increase the V0. However, there is a certain point in which the substrate concentration is too high and the reaction will not proceed any faster. This point is called the maximum veloctiy, Vmax. Every enzyme has their unique Vmax value. Another important identity of an enzyme is the Km value, defined as the substrate concentration at half of Vmax. Km is also unique to each enzyme. The turnover rate, the rate at which products are produced is called the Kcat. Dividing Kcat by Km gives the efficiency constant of the enzyme, which tells how fast the reaction is carried out and how likely the enzyme is to find the substrate. For more information, refer to Catalysis

References

  • Berg, Jeremy; Tymoczko, John; Stryer, Lubert. Biochemistry, 6th edition. W.H. Freeman and Company. 2007.

Simple English

File:Heterogeneous
Solid heterogeneous catalysts such as in automobile catalytic converters are plated on structures designed to maximize their surface area.
File:Low Temperature Oxidation
A low-temperature oxidation catalyst used to convert carbon monoxide to non-toxic carbon dioxide at room temperature. It can also remove formaldehyde from the air.

Catalysis is the change in speed (rate) of a chemical reaction due to the help of a catalyst. Unlike other chemicals which take part in the reaction, a catalyst is not consumed by the reaction itself. A catalyst may participate in many chemical reactions. Catalysts that speed the reaction are called positive catalysts. Catalysts that slow the reaction are called negative catalysts, or inhibitors. Substances that increase the activity of catalysts are called promoters, and substances that deactivate catalysts are called catalytic poisons.

A catalyst is something which changes the rate of a chemical reaction. An example is when manganese oxide (MnO2) is added to hydrogen peroxide (H2O2), and the hydrogen peroxide starts to break up into water and oxygen. Catalysts are either of natural or synthetic origin. Catalysts are useful because they leave no residue in the solution they have speeded up. A catalyst can also be used in a reaction again and again as it is not used up. There are many catalysts in our body which play an important part in many biochemical reactions. These are called enzymes. Most catalysts work by lowering the 'activation energy' of a reaction. This allows less energy to be used, thus speeding up the reaction. The opposite of a catalyst is an inhibitor. Inhibitors slow down reactions. Some of them are found in snake venom and are dangerous for our nervous system or heart.

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