Everything about Catalytic totally explained
Catalysis is the process by which the
rate of a
chemical reaction (or
biological process) is increased by means of the addition of a species known as a
catalyst to the reaction. What makes a catalyst different from a chemical reagent is that whilst it participates in the reaction, isn't consumed in the reaction. That is, the catalyst may undergo several chemical transformations during the reaction, but at the conclusion of the reaction, the catalyst is regenerated unchanged. As a catalyst is regenerated in a reaction, often only a very small amount is needed to increase the rate of the reaction.
Overview
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 lower the activation energy, the faster the rate of the reaction.
A real example is the
disproportionation of
hydrogen peroxide to give water and
oxygen:
» 2 H
2O
2 → 2 H
2O + O
2
Whilst the above reaction is favoured in the sense that reaction products are more stable than the starting material, the reaction is slow. This can be seen by the fact that hydrogen peroxide is often available for purchase on the
high-street in bottles as a disinfectant.
However, upon the addition of a small amount of
manganese dioxide, the hydrogen peroxide undergoes a rapid reaction, which can be readily seen by the
effervescence of oxygen. The manganese dioxide may be recovered unchanged, and re-used indefinitely, and thus isn't consumed in the reaction. Accordingly, manganese dioxide catalyses this reaction.
In a more general sense, anything that increases the rate of any process is commonly called a "catalyst" (From the
Greek, meaning
to annul or
to untie or
to pick up). For example a
matchmaker might be called a catalyst, as he or she brings two people together who otherwise might not meet, with the matchmaker being unaltered by the matching process.
The opposite of a catalyst is an
inhibitor which slows the rate of a reaction without itself being consumed.
History
The phrase
catalysed processes was coined by
Jöns Jakob Berzelius in 1836 to describe
reactions which are accelerated by substances which remain unchanged after the reaction. Other early chemists involved in catalysis were
Alexander Mitscherlich who in 1831 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 1820’s.
Humphrey Davy discovered the use of platinum in catalysis. In the 1880s,
Wilhelm Ostwald at
Leipzig University started a series of systematic investigations into reactions that were catalyzed by the presence of
acids and bases, and found both 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.
Typical mechanism
Catalysts generally react with one or more reactants to form an intermediate 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, A and B are reactants, and D is the product of the reaction of A and B:
» A + C → AC (1)
B + AC → ABC (2)
» ABC → CD (3)
CD → C + D (4)
Although the catalyst (C) is consumed by reaction 1, it's subsequently produced by reaction 4, so for the overall reaction:
» A + B → D
Catalytic cycles
A catalytic cycle or catalytic mechanism is a reaction mechanism which involves a catalyst. Catalytic cycles are central to any discussion of catalysis, be it in
biochemistry,
organometallic chemistry, or solid state chemistry.
Often, a so-called
sacrificial catalyst is also part of the reaction system with the purpose of regenerating the true catalyst in each cycle. As the name implies the sacrificial catalyst isn't regenerated and is instead irreversibly consumed. This sacrificial compound is also known as a stoichiometric catalyst when added in
stoichiometric quantities compared to the main reactant. Usually the true catalyst is an expensive and complex molecule and added in quantities as small as possible. The stoichiometric catalyst on the other hand should be cheap and abundant.
Catalysts and reaction energetics
Catalysts work by providing an (alternative) mechanism involving a different transition state and lower
activation energy. The effect of this is that more molecular collisions have the energy needed to reach the transition state. Hence, catalysts can perform reactions that, albeit thermodynamically feasible, wouldn't run without the presence of a catalyst, or perform them much faster, more specific, or at lower temperatures. This can be observed on a
Boltzmann distribution and
energy profile diagram. This means that catalysts reduce the amount of energy needed to start a chemical reaction.
Catalysts
cannot make energetically unfavorable reactions possible — they've
no effect on the
chemical equilibrium of a reaction because the rate of both the forward and the reverse reaction are equally affected (see also
thermodynamics). The net free energy change of a reaction is the same whether a catalyst is used or not; the catalyst just makes it easier to activate.
The
SI derived unit for measuring the
catalytic activity of a catalyst is the
katal, which is moles per second. The degree of 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
Enzyme#Kinetics section.
Autocatalysis
In
autocatalysis, a reaction produces catalysts.
Types of catalysts
Catalysts can be either
heterogeneous or .
Biocatalysts are often seen as a separate group.
Heterogeneous catalysts are present in different
phases from the
reactants (for example, a
solid catalyst in a
liquid reaction mixture), whereas homogeneous catalysts are in the same phase (for example, a
dissolved catalyst in a liquid reaction mixture).
Heterogeneous catalysts
A simple model for
heterogeneous catalysis involves the catalyst providing a
surface on which the reactants (or
substrates) temporarily become
adsorbed.
bonds between the products and the catalyst are weaker, so the products are released. Different possible mechanisms for [[reactions on surfaces are known, depending on how the adsorption takes place (
Langmuir-Hinshelwood and
Eley-Rideal).
For example, in the
Haber process to manufacture
ammonia, finely divided
iron acts as a heterogeneous catalyst. Active sites on the metal allow partial weak bonding to the reactant
gases, which are
adsorbed onto the metal surface. As a result, the bond within the molecule of a reactant is weakened and the reactant molecules are held in close proximity to each other. In this way the particularly strong
triple bond in nitrogen is weakened and the hydrogen and nitrogen molecules are brought closer together than would be the case in the gas phase, so the rate of reaction increases.
Other heterogeneous catalysts include
vanadium(V) oxide in the
contact process,
nickel in the manufacture of
margarine,
alumina and
silica in the
cracking of
alkanes and
platinum,
rhodium and
palladium in
catalytic converters.
Mesoporous silicates have found utility in heterogeneous reaction catalysis because their large accessible surface area allows for high catalyst loading.
In car engines, incomplete
combustion of the
fuel produces
carbon monoxide, which is toxic. The electric spark and high temperatures also allow
oxygen and
nitrogen to react and form
nitrogen monoxide and
nitrogen dioxide, which are responsible for photochemical
smog and
acid rain. Catalytic converters reduce such emissions by adsorbing
CO and
NO onto catalytic surface, where the gases undergo a
redox reaction.
Carbon dioxide and nitrogen are desorbed from the surface and emitted as relatively harmless gases:
» 2CO + 2NO → 2CO
2 + N
2
Many catalysts used in refineries and in petrochemical applications are regenerated and reused multiple times to save costs and energy and to reduce environmental impact from recycling or disposal of spent catalysts.
Homogeneous catalysts
Homogeneous catalysts are in the same phase as the reactants.
In homogeneous catalysis the catalyst is a
molecule which facilitates the reaction. The catalyst initiates reaction with one or more reactants to form intermediate(s) and in some cases one or more products. Subsequent steps lead to the formation of remaining products and to the regeneration of the catalyst.
Examples of homogeneous catalysts are:
1) The ion
H+(aq) which acts as a catalyst in
esterification, as well as in the inverse reaction - hydrolysis of esters such as methyl acetate is catalysed by H
+
2)
Chlorine free radicals in the break down of
ozone. These radicals are formed by the action of
ultraviolet radiation on
chlorofluorocarbons (CFCs). They react with ozone to form oxygen molecules and regenerate the catalyst radicals. This process destroys the thin layer of
stratospheric ozone.
» Cl
· + O
3 → ClO
· + O
2
ClO
· + O
· → Cl
· + O
2
3)
Oxides of nitrogen in the
oxidation of
sulfur dioxide to
sulfur trioxide by
dioxygen in the
chamber process.
Biocatalysts
In nature
enzymes are catalysts in
metabolism. In
biochemistry catalysis is also observed with
abzymes and
ribozymes,
deoxyribozymes have also been created in the laboratory.
Biocatalysts can be thought of as a mixture of a homogenous and heterogeneous catalyst. This is because the enzyme is in solution itself, but the reaction takes place on the enzyme surface. Several factors affect the activity of enzymes. The most important are:
- Temperature
- pH
- Enzyme concentration
- Substrate concentration
Electrocatalysts
In the context of
electrochemistry, specifically in
fuel cell engineering, various metal-rich catalysts are used to promote the efficiency of a
half reaction that occurs within the fuel cell. One common type of fuel cell electrocatalyst is based upon tiny
nanoparticles of
platinum which adorn slightly larger
carbon particles. When this type of platinum electrocatalyst is in contact with one of the
electrodes in a fuel cell, it increases the rate of the
redox half reaction in which
oxygen gas is reduced to water (or
hydroxide or
hydrogen peroxide).
Significance
Catalysis is of paramount importance in the chemical industry. The production of most industrially important chemicals involves catalysis. Two notable commercial processes are the
Haber process for
ammonia synthesis and the
Fischer-Tropsch synthesis. Research into catalysis is a major field in applied science, and involves many fields of chemistry, notably in
organometallic chemistry, and physics. Catalysis is important in many aspects of
environmental science, from the
catalytic converter in automobiles to the alleged causes of the
ozone hole. Catalytic, rather than
stoichiometric reactions are preferred in environmentally friendly
green chemistry due to the reduced amount of waste generated.
Notable examples
Estimates are that 90% of all commercially produced chemical products involve catalysts at some stage in the process of their manufacture.
Manganese dioxide is used in the laboratory to prepare
oxygen by the decomposition of
hydrogen peroxide to
oxygen and
water.
Well-known applications of synthetic catalysts are:
Catalytic converters made from platinum and manganese break down some of the more harmful byproducts of automobile exhaust. The catalysts used are micro-engineered to have large surface areas.
the Haber process for the synthesis of ammonia from nitrogen and hydrogen, where iron is the catalyst.
Examples of catalysts that perform specific transformations on functional groups:
Transformations of olefinic groups:
The most common catalyst is the proton. Many transition metals and transition metal complexes are used in catalysis as well.
New directions - organocatalysis
While transition metal catalysts are well established, a new trend is toward organocatalysis which use comparatively simple organic molecules as catalysts. While typically, catalyst loading is much higher than transition metal-based catalysts, the catalysts are usually commercially available in bulk, helping to reduce costs drastically. Organocatalysts of the "new generation" are competitive to traditional metal-containing catalysts and are owing to low product inhibition applicable in substoichiometric quantities. The chemical character of organocatalysts offers new and attractive perspectives and advantages to synthesis chemists.
Catalytic processes
In 2005, Catalytic processes generated about $900 billion in products worldwide.(External Link
)
Acid-base catalysis
Catalytic converters made from platinum and rhodium break down some of the more harmful byproducts of automobile exhaust.
Fuel cells
Fischer-Tropsch synthesis.
Haber process (synthesis of ammonia from nitrogen and hydrogen, where ordinary iron is used as a catalyst)
Hydrogenation
Methanol synthesis
Nitric acid production
Petroleum refining and processing
Steam reforming of hydrocarbons to produce synthesis gas
Sulfuric acid production
Transesterification
Olefin polymerisationFurther Information
Get more info on 'Catalytic'.
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