1. Kinetics of biological reaction. Influence of different factors on the chemic reaction rate.

3. Catalysis and catalysts. Influence of inorganic catalysts and enzymes on chemical reaction rate.

Chemical kinetics

Chemical kinetics, also known as reaction kinetics, is the study of rates of chemical processes. Chemical kinetics includes investigations of how different experimental conditions can influence the speed of a chemical reaction and yield information about the reaction's mechanism and transition states, as well as the construction of mathematical models that can describe the characteristics of a chemical reaction. In 1864, Peter Waage and Cato Guldberg pioneered the development of chemical kinetics by formulating the law of mass action, which states that the speed of a chemical reaction is proportional to the quantity of the reacting substances.

Chemical kinetics deals with the experimental determination of reaction rates from which rate laws and rate constants are derived. Relatively simple rate laws exist for zero-order reactions (for which reaction rates are independent of concentration), first-order reactions, and second-order reactions, and can be derived for others. In consecutive reactions, the rate-determining step often determines the kinetics. In consecutive first-order reactions, a steady state approximation can simplify the rate law. The activation energy for a reaction is experimentally determined through the Arrhenius equation and the Eyring equation. The main factors that influence the reaction rate include: the physical state of the reactants, the concentrations of the reactants, the temperature at which the reaction occurs, and whether or not any catalysts are present in the reaction.

Factors affecting reaction rate

Nature of the reactants

Depending upon what substances are reacting, the reaction rate varies. Acid/base reactions, the formation of salts, and ion exchange are fast reactions. When covalent bond formation takes place between the molecules and when large molecules are formed, the reactions tend to be very slow. Nature and strength of bonds in reactant molecules greatly influence the rate of its transformation into products.

Physical state

The physical state (solid, liquid, or gas) of a reactant is also an important factor of the rate of change. When reactants are in the same phase, as in aqueous solution, thermal motion brings them into contact. However, when they are in different phases, the reaction is limited to the interface between the reactants. Reaction can occur only at their area of contact; in the case of a liquid and a gas, at the surface of the liquid. Vigorous shaking and stirring may be needed to bring the reaction to completion. This means that the more finely divided a solid or liquid reactant the greater its surface area per unit volume and the more contact it makes with the other reactant, thus the faster the reaction. To make an analogy, for example, when one starts a fire, one uses wood chips and small branches — one does not start with large logs right away. In organic chemistry, on water reactions are the exception to the rule that homogeneous reactions take place faster than heterogeneous reactions.

Concentration

The reactions are due to collisions of reactant species. The frequency with which the molecules or ions collide depends upon their concentrations. The more crowded the molecules are, the more likely they are to collide and react with one another. Thus, an increase in the concentrations of the reactants will result in the corresponding increase in the reaction rate, while a decrease in the concentrations will have a reverse effect. For example, combustion that occurs in air (21% oxygen) will occur more rapidly in pure oxygen.

Temperature

Temperature usually has a major effect on the rate of a chemical reaction. Molecules at a higher temperature have more thermal energy. Although collision frequency is greater at higher temperatures, this alone contributes only a very small proportion to the increase in rate of reaction. Much more important is the fact that the proportion of reactant molecules with sufficient energy to react (energy greater than activation energy: E > Ea) is significantly higher and is explained in detail by the Maxwell–Boltzmann distribution of molecular energies.

The 'rule of thumb' that the rate of chemical reactions doubles for every 10 °C temperature rise is a common misconception. This may have been generalized from the special case of biological systems, where the α (temperature coefficient) is often between 1.5 and 2.5.

A reaction's kinetics can also be studied with a temperature jump approach. This involves using a sharp rise in temperature and observing the relaxation time of the return to equilibrium. A particularly useful form of temperature jump apparatus is a shock tube, which can rapidly jump a gases temperature by more than 1000 degrees.

Catalysts

A catalyst is a substance that accelerates the rate of a chemical reaction but remains chemically unchanged afterwards. The catalyst increases rate reaction by providing a different reaction mechanism to occur with a lower activation energy. In autocatalysis a reaction product is itself a catalyst for that reaction leading to positive feedback. Proteins that act as catalysts in biochemical reactions are called enzymes. Michaelis–Menten kinetics describe the rate of enzyme mediated reactions. A catalyst does not affect the position of the equilibria, as the catalyst speeds up the backward and forward reactions equally.

In certain organic molecules, specific substituents can have an influence on reaction rate in neighbouring group participation.

Agitating or mixing a solution will also accelerate the rate of a chemical reaction, as this gives the particles greater kinetic energy, increasing the number of collisions between reactants and, therefore, the possibility of successful collisions.

Pressure

Increasing the pressure in a gaseous reaction will increase the number of collisions between reactants, increasing the rate of reaction. This is because the activity of a gas is directly proportional to the partial pressure of the gas. This is similar to the effect of increasing the concentration of a solution.

In addition to this straightforward mass-action effect, the rate coefficients themselves can change due to pressure. The rate coefficients and products of many high-temperature gas-phase reactions change if an inert gas is added to the mixture; variations on this effect are called fall-off and chemical activation. These phenomena are due to exothermic or endothermic reactions occurring faster than heat transfer, causing the reacting molecules to have non-thermal (non-Boltzmann) energy distributions. Increasing the pressure increases the heat transfer rate between the reacting molecules and the rest of the system, reducing this effect.

Condensed-phase rate coefficients can also be affected by (very high) pressure; this is a completely different effect than fall-off or chemical-activation. It is often studied using diamond anvils.

A reaction's kinetics can also be studied with a pressure jump approach. This involves making fast changes in pressure and observing the relaxation time of the return to equilibrium.

Equilibrium

While chemical kinetics is concerned with the rate of a chemical reaction, thermodynamics determines the extent to which reactions occur. In a reversible reaction, chemical equilibrium is reached when the rates of the forward and reverse reactions are equal and the concentrations of the reactants and products no longer change. This is demonstrated by, for example, the Haber–Bosch process for combining nitrogen and hydrogen to produce ammonia. Chemical clock reactions such as the Belousov–Zhabotinsky reaction demonstrate that component concentrations can oscillate for a long time before finally attaining the equilibrium.

Free energy

In general terms, the free energy change (ΔG) of a reaction determines whether a chemical change will take place, but kinetics describes how fast the reaction is. A reaction can be very exothermic and have a very positive entropy change but will not happen in practice if the reaction is too slow. If a reactant can produce two different products, the thermodynamically most stable one will in general form, except in special circumstances when the reaction is said to be under kinetic reaction control. The Curtin–Hammett principle applies when determining the product ratio for two reactants interconverting rapidly, each going to a different product. It is possible to make predictions about reaction rate constants for a reaction from free-energy relationships.

The kinetic isotope effect is the difference in the rate of a chemical reaction when an atom in one of the reactants is replaced by one of its isotopes.

Chemical kinetics provides information on residence time and heat transfer in a chemical reactor in chemical engineering and the molar mass distribution in polymer chemistry.

Applications

The mathematical models that describe chemical reaction kinetics provide chemists and chemical engineers with tools to better understand and describe chemical processes such as food decomposition, microorganism growth, stratospheric ozone decomposition, and the complex chemistry of biological systems. These models can also be used in the design or modification of chemical reactors to optimize product yield, more efficiently separate products, and eliminate environmentally harmful by-products. When performing catalytic cracking of heavy hydrocarbons into gasoline and light gas, for example, kinetic models can be used to find the temperature and pressure at which the highest yield of heavy hydrocarbons into gasoline will occur. Kinetics is also a basic aspect of chemistry.

Concentration Changes

Chemical kinetics is the study of the speed with which a chemical reaction occurs and the factors that affect this speed. This information is especially useful for determining how a reaction occurs.

What is meant by the speed of a reaction? The speed of a reaction is the rate at which the concentrations of reactants and products change.

Consider the following hypothetical example. The letters A, B, and C represent chemical species (in this context, the letters do not represent elements). Suppose the following imaginary reaction occurs:

A + 2 B → 3 C

The simulation below illustrates how this reaction can be studied. The apparatus at the left is called a stopped-flow apparatus. Each syringe contains a solution filled with a different reactant (A or B). When the two solutions are forced out of the syringes, they are quickly mixed in a mixing block and the reaction starts. The reacting solution passes through the tube at the bottom. An analytical technique such as spectrophotometry is used to measure the concentrations of the species in the reaction mixture (which is in the tube at the bottom) and how those concentrations change with time.

In this example, the syringe at the left contains a solution of species A, which has a yellow color. The syringe at the right contains a solution of species B, which has a light blue color. The product C has a red color.

Chemical Kinetics

Chemical kinetics is the study and discussion of chemical reactions with respect to reaction rates, effect of various variables, re-arrangement of atoms, formation of intermediates etc. There are many topics to be discussed, and each of these topics is a tool for the study of chemical reactions. By the way, the study of motion is called kinetics, from Greek kinesis, meaning movement.

At the macroscopic level, we are interested in amounts reacted, formed, and the rates of their formation. At the molecular or microscopic level, the following considerations must also be made in the discusion of chemical reaction mechanism.

Molecules or atoms of reactants must collide with each other in chemical reactions.

The molecules must have sufficient energy (discussed in terms of activation energy) to initiate the reaction.

In some cases, the orientation of the molecules during the collision must also be considered.

Reaction Rates

Chemical reaction rates are the rates of change in concentrations or amounts of either reactants or products. For changes in amounts, the units can be one of mol/s, g/s, lb/s, kg/day etc. For changes in concentrations, the units can be one of mol/(L s), g/(L s), %/s etc.

With respect to reaction rates, we may deal with average rates, instantaneous rates, or initial rates depending on the experimental conditions.

Thermodynamics and kinetics are two factors that affect reaction rates. The study of energy gained or released in chemical reactions is called thermodynamics, and such energy data are called thermodynamic data. However, thermodynamic data have no direct correlation with reaction rates, for which the kinetic factor is perhaps more important. For example, at room temperature (a wide range of temperatures), thermodynamic data indicates that diamond shall convert to graphite, but in reality, the conversion rate is so slow that most people think that diamond is forever.

Factors Influence Reaction Rates

Many factors influence rates of chemical reactions, and these are summarized below. Much more extensive discussion will be given in other pages.

Nature of Reactants

Acid-base reactions, formation of salts, and exchange of ions are fast reactions. Reactions in which large molecules are formed or break apart are usually slow. Reactions breaking strong covalent bonds are also slow.

Temperature

Usually, the higher the temperature, the faster the reaction. The temperature effect is discussed in terms of activation energy.

Concentration Effect

The dependences of reaction rates on concentrations are called rate laws. Rate laws are expressions of rates in terms of concentrations of reactants. Keep in mind that rate laws can be in differential forms or integrated forms. They are called differential rate laws and integrated rate laws. The following is a brief summary of topics regarding rate laws.

rate laws: differential and integrated rate laws.

Integrated rate laws: First Order Reactions

Second Order Reactions

Rate laws apply to homogeneous reactions in which all reactants and products are in one phase (solution).

Heterogeneous reactions: reactants are present in more than one phase

For heterogeneous reactions, the rates are affected by surface areas.

Catalysts: substances used to facilitate reactions

By the nature of the term, catalysts play important roles in chemical reactions.

Reaction Mechanisms

The detailed explanation at the molecular level how a reaction proceeds is called reaction mechanism. The explanation is given in some elementary steps. Devising reaction mechanisms requires a broad understanding of properties of reactants and products, and this is a skill for matured chemists. However, first year chemistry students are often given a mechanism, and be asked to derive the rate law from the proposed mechanism. The steady-state approximations is a technique for deriving a rate law from the proposed mechanism.

Rate of Reaction

The rate of a reaction is the speed at which a reaction happens. If a reaction has a low rate, that means the molecules combine at a slower speed than a reaction with a high rate. Some reactions take hundreds, maybe even thousands, of years while others can happen in less than one second. The rate of reaction depends on the type of molecules that are combining. If you want to think of a very slow reaction, think about how long it took dinosaur bones to become fossils through breakdown. You can thank chemical processes in bacteria for most of those dinosaur bones in the museum.

There is another big idea for rates of reaction called collision theory. The collision theory says that as more collisions in a system occur, there will be more combinations of molecules bouncing into each other. If there are a higher number of collisions in a system, more combinations of molecules can occur. The reaction will go faster and the rate of that reaction will be higher. Even though they are both liquids, think about how slowly molecules move in honey when compared to your soda. There are a lower number of collisions in the honey.

Reactions happen - no matter what. Chemicals are always combining or breaking down. The reactions happen over and over, but not always at the same speed. A few things affect the overall speed of the reaction and the number of collisions that can occur.

Concentration: If there is more of a substance in a system, there is a greater chance that molecules will collide and speed up the rate of the reaction. If there is less of something, there will be fewer collisions and the reaction will probably happen at a slower speed. Sometimes when you are in a chemistry lab, you will add one solution to another. When you want the rate of reaction to be slower, you will add only a few drops at a time instead of the entire beaker.

Temperature: When you raise the temperature of a system, the molecules bounce around a lot more because they have more energy. When they bounce around more, they are more likely to collide. That fact means they are also more likely to combine. When you lower the temperature, the molecules are slower and collide less. That temperature drop lowers the rate of the reaction. Back to the chemistry lab! Sometimes you will mix solutions in ice so that the temperature of the system stays cold and the rate of reaction is slower.

Pressure: Pressure affects the rate of reaction, especially when you look at gases. When you increase the pressure, the molecules have less space in which they can move. That greater density of molecules increases the number of collisions. When you decrease the pressure, molecules don't hit each other as often. The lower pressure decreases the rate of reaction.

Rates of Reaction

What factors influence the rate of a chemical reaction?

1. Temperature

2. Catalysts

3. Concentrations of reactants

3. Surface area of a solid reactant

4. Pressure of gaseous reactants or products

If you are planning an investigation, I suggest that you investigate the effects of temperature or the effects of the concentration of the reactants because these will allow you to choose a suitable range of values for the controlled or independent variable. The dependent variable will be the rate of the reaction. Keep all the other variables fixed.

To make a prediction for your investigation you will have to ask yourself the question: What will happen to the rate of the reaction when I increase the temperature? or What will happen to the rate of the reaction if I increase the concentration of one of the reactants? The answer to that question is your prediction. The next thing to do is to explain your prediction. You will have to answer the question: Why will the reaction go faster if I increase the temperature? or Why will the reaction go faster if I increase the concentration of one of the reactions? The answer to this question is your explanation, and to get the highest possible marks, you will have to provide a full scientific explanation.

Once you have written your hypothesis (prediction with explanation) you will decide how to do the experiments, i.e. write the proposed method.

How does temperature affect the rate of a chemical reaction?

When two chemicals react, their molecules have to collide with each other with sufficient energy for the reaction to take place. This is collision theory. The two molecules will only react if they have enough energy. By heating the mixture, you will raise the energy levels of the molecules involved in the reaction. Increasing temperature means the molecules move faster. This is kinetic theory. If your reaction is between atoms rather than molecules you just substitute “atom” for “molecule” in your explanation.

How do catalysts affect the rate of a reaction?

Catalysts speed up chemical reactions. Only very minute quantities of the catalyst are required to produce a dramatic change in the rate of the reaction. This is really because the reaction proceeds by a different pathway when the catalyst is present. Adding extra catalyst will make absolutely no difference. There is a whole page on this site devoted to catalysts.

How does concentration affect the rate of a reaction?

Increasing the concentration of the reactants will increase the frequency of collisions between the two reactants. So this is collision theory again. You also need to discuss kinetic theory in an experiment where you vary the concentration. Although you keep the temperature constant, kinetic theory is relevant. This is because the molecules in the reaction mixture have a range of energy levels. When collisions occur, they do not always result in a reaction. If the two colliding molecules have sufficient energy they will react.

If reaction is between a substance in solution and a solid, you just vary the concentration of the solution. The experiment is straightforward. If the reaction is between two solutions, you have a slight problem. Do you vary the concentration of one of the reactants or vary the concentration of both? You might find that the rate of reaction is limited by the concentration of the weaker solution, and increasing the concentration of the other makes no difference. What you need to do is fix the concentration of one of the reactants to excess. Now you can increase the concentration of the other solution to produce an increase in the rate of the reaction.

How does surface area affect a chemical reaction?

If one of the reactants is a solid, the surface area of the solid will affect how fast the reaction goes. This is because the two types of molecule can only bump into each other at the liquid solid interface, i.e. on the surface of the solid. So the larger the surface area of the solid, the faster the reaction will be.

Smaller particles have a bigger surface area than larger particle for the same mass of solid. There is a simple way to visualize this. Take a loaf of bread and cut it into slices. Each time you cut a new slice, you get an extra surface onto which you can spread butter and jam. The thinner you cut the slices, the more slices you get and so the more butter and jam you can put on them. This is “Bread and Butter Theory”. You should have come across the idea in your biology lessons. By chewing your food you increase the surface area so that digestion can go faster.

What affect does pressure have on the reaction between two gasses?

You should already know that the atoms or molecules in a gas are very spread out. For the two chemicals to react, there must be collisions between their molecules. By increasing the pressure, you squeeze the molecules together so you will increase the frequency of collisions between them. This is collision theory again.

In a diesel engine, compressing the gaseous mixture of air and diesel also increases the temperature enough to produce combustion. Increasing pressure also results in raising the temperature. It is not enough in a petrol engine to produce combustion, so petrol engines need a spark plug. When the petrol air mixture has been compressed, a spark from the plug ignites the mixture. In both cases the reaction (combustion) is very fast. This is because once the reaction has started, heat is produced and this will make it go even faster.

The Rate Law

The rate law for a chemical reaction links the reaction rate with concentrations or pressures of reactants and constant parameters.

Rate Laws for Various Reactions

A variety of reaction orders are observed, and they cannot be easily correlated with the stoichiometry of the reaction.

Rate Law

Rate = K[A]x[B]y

Rate = Reaction Rate Reaction Order

The sum of x and y.

The reaction of bromine and formic acid is first order in bromine, zeroth order in formic acid, and first order overall.

K = rate constant

x/y = determined numbers

A/B = concentrations

How fast is Fast? The Mathematics of Change

Consider a reaction of Red molecules (A) to make Blue molecules (B), i.e. A -> B. If we were able to see the reaction on a molecular scale, the reaction of each individual molecule of occurs very rapidly, but the overall color of the vessel changes more slowly. Snapshots of the reaction in progress might look like this:

The number of reactants and products in the reaction vessel changes with time, with the relative number of reactant molecules destroyed and number of products formed per reaction event determined by the reaction stoichiometry. Each reactant molecule is identical to every other one, but they all don't react at the same instant. At each point in time, the probability of reaction per unit time is the same for each molecule in the sample, and that probability influences the overall reaction rate. But that isn't the only thing that determines the overall reaction rate. The total number of reactions at any instant is the probability of reaction per unit time multiplied by the number of reactants remaining in the vessel. Thus the reaction proceeds quickly at first, when there are lots of reactants around, and appears to slow as the reactants are consumed. A simulation of a similar reaction involving reactive collisions between molecules can be run on you browser. A plot of the time dependence of the number of molecules of each type looks 'smooth' when there are lots and lots of molecules in the sample, so individual reaction events get 'averaged' out. The concentrations of the reactants and products change in time like this:

The Rate Law

The equation that describes the dependence of the reaction rate on the concentrations of the species in the reactor is called a rate law. The rate law for a given reaction is determined from the reaction mechanism. Several important kinds of simple rate laws are worth noting

First Order Rate Law

The simplist reaction mechanism is that of unimolecular decomposition (or isomerisation). In such a process, a single reactant undergoes a transformation at a constant probability per unit time. Such a mechanism leasds to a first-order reaction rate law. Examples of reactions such as these are radioactive decay, bacterial growth, and compound interest. Let's assume the reaction has a simple stoichimetry:

A → B

A First-Order Rate Law is called such because the rate of product formation ( or reactant depletion ) is proportional to the first power of the number of available reactants (or reactant concentration):

rate =k[A]

where [A] represents the concentration (number density) of species A in the sample.

Second Order Rate Law

If two molecules undergo a bimolecular reaction such as a reaction that involves a collisional encounter to produce products, and has a stoichiometry like this:

A + A → B + C

Zeroth Order Rate Law

If a reaction is catalysed by a surface and has enough (excess) reactant, the rate of the reaction depends on the area of the catalyst, not on how much reactant is present. This is an unusual circumstance outside of the realm of catalyzed reactions and is described by a Zeroth Order rate law:

rate =k

THE EFFECT OF CONCENTRATION ON REACTION RATES

This page describes and explains the way that changing the concentration of a solution affects the rate of a reaction. Be aware that this is an introductory page only. If you are interested in orders of reaction, you will find separate pages dealing with these. You can access these via the rates of reaction menu (link at the bottom of the page).

For many reactions involving liquids or gases, increasing the concentration of the reactants increases the rate of reaction. In a few cases, increasing the concentration of one of the reactants may have little noticeable effect of the rate. These cases are discussed and explained further down this page.

Don't assume that if you double the concentration of one of the reactants that you will double the rate of the reaction. It may happen like that, but the relationship may well be more complicated.

The Temperature Dependence of Reaction Rates

Chemical Activation

Consider the reaction

H₂ + Cl₂ -> 2HCl

On a molecular level, bonds must be broken (H-H and Cl-Cl) before the reaction can proceed too far into products. This means that as the reactant molecules come together, the collision must have enough energy to initiate the bond breakage for the reaction to occur. Not all collisions will have this amount of energy. The collisions that do not have sufficient energy to react end up as elastic scattering events.

Only collisions with enough energy react to form products. The energy of the system changes as the reactants approach each other. The critical amount of energy to make the reaction proceed is called the Activation Energy.

The Reaction Coordinate is the 'distance' along the path of the reaction, and is plotted along the horizontal axis. The energy of interaction of the reactive system is plotted vertically, and is called the Chemical potential, or just potential energy. You fight gravitational potential energy when you try to roll a boulder over a mountain.

A chemical potential of interaction usually looks like something like the graph above, which is similar to the 'pushing a boulder over a hill' graph above. The Graph above is drawn for the isomerization of an isonitrile that we discussed before. The barrier to the isomerization keeps the unstable CH3NC from reacting away quickly at low temperature, even though energy is released upon the net reaction.

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.

Factors affection the reaction rate.

The fate of any particular reaction depends upon the following factors:

1. Nature of the reactants. Consider the following two reactions

These reactions appear to be similar but the first is fast while the second is slow. This is because different amounts of energies are required for breaking of different bonds and different amounts of energies are released in the formation of different bonds.

2. Concentration of the reactants. Greater are the concentrations of the reactants, faster is the reaction, as the concentrations of the reactants decrease, the rate of reaction also decreases.

3. Temperature. The rate of reaction increases with increase of temperature. In most of the cases, the rate of reaction becomes nearly double for 10К rise of temperature. In коте смея, reactions do not take place at room temperature but take place at higher temperature.

4. Presence of Catalyst. А catalyst generally increases the speed of а reaction without itself being consumed in the reaction. In case of reversible reactions, а catalyst helps to attain the equilibrium quickly without disturbing the state of equilibrium.

5. Surface area of the reactants. For а reaction involving а solid reactant or catalyst, the smaller is the particle size i.е., greater is the surface area, the fast r is the reaction.

6. Presence of light. Some reactions do not take place in the dark but take place in the presence of light e.g.,

Н₂ + С1₂ = 2НС1. Such reactions are called “photochemical reactions”

Rate laws and rate constants. It is often found that the rate of reaction is proportional to the concentrations of the reactants raised to а power. For example, 1с may be found that the rate is proportional to the concentrations of two reactants А and В, and that:

where each concentration is raised to the first power. The coefficient k is called the rate constant for the reaction or velocity constant. The rate constant is independent of the concentrations but depends on the temperature. An experimentally determined equation of this kind is called the rate law of the reaction. More formally, а rate law is an equation that expresses the rate of reaction as а function of the concentrations of all the species present in the overall chemical equation for the reaction.

If all concentrations are take as unity, [A] = [B] = 1 mole/liter, then rate = k.

Hence rate constant may be defined as the rate of the reaction when the concentration of each reactants is take as unity. That is why the rate constant is also called specific reaction rate.

Characteristics of rate constant. Some important characteristics of the rate constant are as follows:

1. Rate constant is a measure of the rate of reaction. Greater is the value of the rate constant, factors is the reaction.

2. Each reaction has a definite value of the rate constant at particular temperature.

3. The value of the rate constant for the same reaction changes with temperature.

4. The value of the rate constant of a reaction does not depend upon the concentration of the reactants.

5. The units of the rate constant depend upon the order of reaction.

А practical application of а rate law is that, once we know it and the value of the rate constant, we can predict the rate of reaction from the composition of the mixture. Moreover, as we shall see later, by knowing the rate law we can go on to predict the composition of the reaction mixture at а later stage of the reaction. The theoretical usefulness of a rate law is that it is а guide to the mechanism of the reaction, for any proposed mechanism must be consistent with the observed rate law.

Order of reaction

The sum of the concentration terms on which the rate of а reaction actually depends as observed experimentally is called the order of the reaction. For example, in the above case, order of reaction = а + p. Thus the orders a reaction may also be defined as the sum of the exponents (powers) to which the concentration terms in the rate law equation are raised to express the observed rate of the reaction.

The power to which the concentration of а species is raised in а rate law is the order of the reaction with respect to that species. А reaction with the rate law is first-order in А and first-order in В. The overall order of а reaction is the sum of the orders of all the components. The rate law is therefore second-order overall.

Some reactions obey а zero-order rate law, and therefore have а rate that is independent of the concentration of the reactant (so long as some is present). Thus, the catalytic decomposition of phosphine (РН₃) on hot tungsten at high pressures has the rate law: u = k

The PH₃ decomposes at а constant rate until it has almost entirely disappeared. Only heterogeneous reactions can have rate laws that are zero-order overall.

These remarks point to three problems. First, we must see how to identify the rate law and obtain the rate constant from the experimental data. We shall concentrate on this aspect in this chapter. Second, we must see how to construct reaction mechanisms that are consistent with the rate law. k₂ + k₃[В]₀

It is most important to distinguish molecularity from order: Reaction order is an empirical quantity, and obtained from thy experimental rate law. The molecularity refers to an elementary reaction proposed as an individual step in а mechanism.

In contrast to reactions in general, the rate law of an elementary reaction can be written down from its chemical equation. Thus, the rate law of а unimolecular elementary reaction is first-order in the reactant;

А ® Products : d[А]/dt = - k [А]

А unimolecular reaction is first-order because the number of А molecules that decay in а short interval is proportional to the number available to decay. (Ten times as many decay in the some interval when there are initially 1000 А molecules than when there are only 100 present). Therefore, the rate of decomposition of А is proportional to its molar concentration.

An elementary bimolecular reaction has а second-order rate law:

А bimolecular reaction is second-order because its rate is proportional to the rate at which the reactant species meet, which is proportional their concentrations. Therefore, if we believe (or simply postulate) that a reaction is а single-step, bimolecular process, then we can write down the rate law (and then go on to test it). Bimolecular elementary reactions are believed to account for many homogeneous reactions, such as the dimerizations of alkenes and dienes and reactions such as:

CH₃I(alc) + СН₃СН₂О^- (alc) ®СН₃ОСН₃СН₂(аlс) + I^-(а1с);

(where 'alc' signifies alcohol solution). The mechanism of the last reaction is believed to be the single elementary step: CH₃I + СН₃СН₂О^- ®СН₃ОСН₃СН₂ + I^-

u = k[CH₃I] [CH₃CH₂O^-]

The interpretation of а rate law is full of pitfalls, partly because а second-order rate law, for instance, can also result from а complex reaction scheme. We shall see below how to string simple steps together into а mechanism and how to arrive at the corresponding rate law. For the present we emphasize that if the reaction is an elementary bimolecular process, then it has second-order kinetics, but if the kinetics are second-order, then the reaction might be complex. The postulated mechanism can be explored only by detailed detective work on the system, and by investigating whether side products or intermediates appear during the course of the reaction. Detailed analysis of this kind was one of the ways, for example, in which the reaction H₂(g) + I₂(g) ® 2HI(g) was shown to proceed by а complex reaction after many years during which it had been accepted on good, but insufficiently meticulous evidence, that it was а fine example of а simple bimolecular reaction in which atoms exchanged partners during а collision.

It is found that the rates of most reactions increase as the temperature is raised. Many reactions fall somewhere in the range spanned by the hydrolysis of methyl ethanoate (where the rate constant at 35⁰С is 1.82 times that at 25⁰С) and hydrolysis of sucrose (where the factor is 4.13).

The Arrhenius parameters. An empirical observation is that many reactions have rate constants that follow the Arrhenius equation.

That is, for many reactions it is found that а plot of ln k against 1/Т gives а straight line.

The factor А is called the pre-exponential factor or the frequency factor; Е_a is called the activation energy. Collectively, the two quantities are called the Arrhenius parameters of the reaction. This equation is sometimes written in an alternative form that combines the two parameters:

The quantity D⁺G is called the activation Gibbs energy. In this form, the expression for the rate constant strongly resembles the formula for the equilibrium constant in terms of the standard reaction Gibbs energy.

For the present chapter we shall regard the Arrhenius parameters as purely empirical quantities that enable us to discuss the variation of rate constants with temperature. There we shall see that the activation energy is the minimum energy that reactants must have in order to from products. For example, in а gas-phase reaction there are numerous collisions each second, but only а tiny proportion of them are sufficiently energetic to lead to reaction. The fraction of collisions with a kinetic energy in excess of an energy Е_a is given by the Boltzmann distribution as е^-Ea/RT. Hence, the exponential factor can be interpreted as the fraction of collisions that have enough energy to lead to reaction.

The analogous interpretation of the pre-exponential factor is that it is a measure of the rate at which collisions occur irrespective of their energy. Hence the product of А and the exponential factor gives the rate of successful collisions.

The temperature dependence of some reactions is not Arrhenius-like.

This definition reduces to the earlier one (as the slope of an Arrhenius plot) for а temperature-independent activation energy. Thus, by using d(l/Т) = - dT/Т² we can rearrange equation:

CATALYSIS.

It is found that the rates of many reactions are increased by the presence of а catalyst, а substance that increases the rate of а reaction without being consumed by it. Although at first thought this may seem impossible, it can indeed occur, because а catalyst is а substance that is used in one step in the mechanism for а reaction and is regenerated in а subsequent step. А catalyst acts by making available а new reaction mechanism with а lower activation energy.

Figure.1.shows the uncatalyzed path of а reaction contrasted with its catalyzed path. (Each potential-energy maximum corresponds to the formation of an activated complex.) Note that ЬН for the reaction is independent of the reaction mechanism, and depends only upon the identity of the reactants and products. However, the activation energy for the catalyzed path is less than that for the uncatalyzed path. Thus, at any given temperature more reactant molecules possess the activation energy for the catalyzed reaction than for the uncatalyzed one. The catalyzed mechanism thus predominates. А catalyst does not eliminate а reaction mechanism; rather, it offers а new, faster one. Mоre molecules, often almost all of them, will follow the new (catalyzed) pathway the products, instead of the old.

If the activation energy of а reaction is high, at normal temperatures only а small proportion of molecular encounters result in reaction. А catalyst lowers the activation energy of the reaction by providing an alternative path that avoids the slow, rate-determining step of the uncatalysed reaction, and results in а higher reaction rate at the same temperature. Catalysts can be very effective; for instance, the activation energy for the decomposition of hydrogen peroxide in solution is 76 kJ/mol, and the reaction is slow at room temperature. When а little iodide is added, the activation energy falls to 57 kJ/mol, and the rate increases by а factor of 2000. Enzymes, which are biological catalysts, are very specific and can have а dramatic effect on the reactions they control. The activation energy for the acid hydrolysis of sucrose is 107kJ/mol, but the enzyme saccharase reduces it to 36 kJ/mol, corresponding to an acceleration of the reaction by а factor of 10⁰ at blood temperature (310 К).

А homogeneous catalyst is а catalyst that is in the same phase as the reaction mixture (е.g. an acid added to an aqueous solution).

А heterogeneous catalyst is in а di6erent phase (е.g. а solid catalyst for а gas-phase reaction).

Homogeneous catalysis.

In homogeneous catalysis, the catalyst and the reactants are present in the same phase. Consider the elementary process

А + В ® products (slow)

Assume that this process has а high activation energy. If we now add catalyst C the reaction mixture, а new, two-step mechanism is possible, in which rate-determining step (step 1, below) has а lower activation energy:

Step 1: А + С ® АС (fast)

Step 2: АС + В ® products + С (faster)

Here, both activation energies are low, and each reaction is faster than original, uncatalyzed reaction. Notice that the overall net equation is changed, and that while catalyst С is used up in step 1, it is regenerated step 2. The rate law for the uncatalyzed reaction is: rate = k[A][B]

and for the catalyzed reaction, rate = k'[А][C]

An example of homogeneous catalysis is found in the oxidation of sulfur dioxide to sulfur trioxide by oxygen, using nitrogen oxide, NO, as а catalyst.

The net equation for the reaction is

2SO₂ (g) + O₂ (g) ®2 SO₃ (g)

The uncatalyzed reaction is very slow, either because it is termolecular (unlikely) or because one step in its reaction mechanism has а very high activation energy. Addition of nitrogen oxide, NO, to the mixture greatly speeds the reaction by making the following mechanism available:

Step 1: O₂ (g) + 2NO(g) ® 2NO₂ (g)

Step 2: [NO ₂ (g) + SO₂ (g) ® NO (g) + SO₃(g)] х 2

The sum of these gives the original net equation, and because the activation energy for each step is fairly low, the reaction proceeds more rapidly than via the uncatalyzed path.

Sоmе idea of the mode of action of homogeneous catalysts can be obtained by examining the kinetics of the bromide-catalysed decomposition of hydrogen peroxide:

2Н₂О₂(aq) ®2Н₂О₂(aq) + О₂(g)

Heterogeneous catalysis.

А heterogeneous catalyst is one which provides а surface on which molecules can readily combine. The process of heterogeneous catalysis begins with the adsorption of а molecule on the surface of the catalyst. There are two general types of adsorption: the relatively weak physical, or van-der-Waals, adsorption and the stronger chemisorption. Evidence that а chemisorbed molecule is relatively strongly bonded at the surface comes from the fact that much more heat is usually evolved during chemisorption than during physical adsorption.

Chemisorption is common in surface catalysis; it apparently takes place preferentially at certain sites on the surface, called active sites or active centers.

These are believed to be related to surface defects or emergences of dislocations.

The chemisorbed molecule is structurally changed at the active site so that it can more readily react with another molecule. There is evidence that some molecules become dissociated into highly reactive fragments. On certain metal surfaces hydrogen, for example, is dissociated into atoms which can react more rapidly than H~ molecules. The reaction of ethylene, С₂Н₂, with hydrogen,

H₂ (g) + C₂H₄ (g) ® C₂H₆ (g)

is thought To be surface-catalyzed by nickel metal in this way.

Catalytic activity at surfaces.

А catalyst acts by providing an alternative reaction path with а lower activation energy. It does not disturb the final equilibrium composition of the system, only the rate at which that equilibrium is approached. In this section we shall consider heterogeneous catalysis, in which the catalyst and the reagents are in different phases. For simplicity, we shall consider only gas/solid systems and the solids we consider will be primarily metals. In practice, industry makes use of а wide range of complex solid catalysts, including oxides and zeolites.

Adsorption and catalysis. Heterogeneous catalysis normally depends on at least one reactant being adsorbed (usually chemisorbed) and modified to а form in which it readily undergoes reaction. Often this modification takes the form of а fragmentation of the reactant molecules.

The Eley-Rideal mechanism. In the Еlеу-Rideal mechanism of а surface-catalysed reaction, а gas-phase molecule collides with another molecule adsorbed on the surface. The rate of formation of product is expected to be proportional to the partial pressure p_b of the non-adsorbed gas В and the extent of surface coverage О„of the adsorbed gas А. It follows that the rate law should be

А + В ® Р; u = kp_Bq

The rate constant k might be much larger than for the uncatalysed gas-phase reaction because the reaction on the surface has а low activation energy and the adsorption itself is often not activated.

Molecular beam studies are able to give detailed information about catalysed reactions. It has become possible to investigate How the catalytic activity of а surface depends on its structure as well as its composition. For instance, the cleavage of С-Н and Н -Н bonds appears to depend on the presence of steps and kinks, and а terrace often has only minimal catalytic activity. The reaction

Н₂ +D₂ ®2HD

has been studied in detail, and it is found that terrace sites are inactive but one molecule in ten reacts when it strikes а step. Although the step itself might be the important feature, it may be that the presence of the step merely exposes а more reactive crystal face (the step face itself). Likewise, the dehydrogenation of hexane to hexene depends strongly on the kink density, and it appears that kinks are needed to cleave С – С bonds. These observations suggest а reason why even small amounts of impurities may poison а catalyst: they are likely to attach to step and kink sites, and so impair the activity of the catalyst entirely. А constructive outcome is that the extent of dehydrogenation may be controlled relative to other types of reactions by seeking impurities that adsorb at kinks and act as specific poisons.

Examples of catalysis. Almost the whole of modern chemical industry depends on the development, selection, and application of catalysts. All we can hope to do is this section is to give а brief indication of some of the problems involved. Other than the ones we consider, these include the danger of the catalyst being poisoned by by-products or impurities and economic considerations relating to cost and lifetime.

In order to be active, the catalyst should be extensively covered by adsorbate, which is the case if chemisorption is strong. On the other hand, if the strength of the substrate-adsorbate bond becomes too great, the activity declines either because the other reactant molecules cannot react with the adsorbate or because the adsorbate molecules are immobilized on the surface. This suggests that the activity of а catalyst should initially increase with strength of adsorption (as measured, for instance, by the enthalpy of adsorption) and then decline, and that the most active catalysts should be those lying near the summit of the volcano. The most active metals are those lying close to the middle of the d block..

Manу metals are suitable for adsorbing gases, and the general order of adsorption strengths decreases along the series O₂, С₂Н₂, С₂Н₄, CO, Н₂, CO₂, N₂. Some of these molecules adsorb dissociatively (е.g. Н,). Elements from the d block, such as iron, vanadium, and chromium, show а strong activity towards all these gases, but manganese and copper are unable to adsorb N₂ and CO₂. Metals towards the left of the periodic table (е.g. magnesium and lithium) can adsorb (and, in fact, react with) only the most active gas (О₂).

Hydrogenation. An example of catalytic action is found in the hydrogenation of alkenes. The alkene (5) adsorbs by forming two bonds with the surface (6), and on the вате surface there may be adsorbed Н atoms. When an encounter occurs, one of the alkene - surface bonds is broken (6 ®7 or 8) and later an encounter with а second Н atom releases the fully hydrogenated hydrocarbon, which is the thermodynamically more stable species.

The evidence for а two-stage reaction is the appearance of different isomeric alkenes in the mixture. The formation of isomers comes about because while the hydrocarbon chain is waving about over the surface of the metal, it might chemisorb again (8 ® 9) and desorb to 10, an isomer of the original 5. The new alkene would not be formed if the two hydrogen atoms attached simultaneously.

А major industrial application of catalytic hydrogenation is to the formation of edible fats from vegetable and animal oils. Raw oils obtained from sources such as the soya bean have the structure CH₂(O₂CR)CH-(O₂CR')CH₂(О₂CR’’), where R, R', and R’’ are long-chain hydrocarbons with several double bonds. One disadvantage of the presence of many double bonds is that the oils are susceptible to atmospheric oxidation, and therefore are liable to become rancid. The geometrical configuration of the chains is responsible for the liquid nature of the oil, and in many applications а solid fat is at least much better and often necessary. Controlled partial hydrogenation of an oil with а catalyst carefully selected so that hydrogenation is incomplete and so that the chains do not isomerize (nickel, in fact), is used on а wide scale to produce edible fats. The process, and the industry, is not made any easier by the seasonal variation of the number of double bonds in the oils.

Oxidation: Catalytic oxidation is also widely used in industry and in pollution control. Although in после cases it is desirable to achieve complete oxidation (as in the production of nitric acid from ammonia); in others partial oxidation is the aim. For example, the complete oxidation of propene to carbon dioxide and water is wasteful, but its partial oxidation to propenal (acrolein, СН₂=СНСНО) is the start of important industrial processes. Likewise, the controlled oxidations of ethene to ethanol, acetaldehyde, and (in the presence of acetic acid or chlorine) to vinyl acetate or vinyl chloride are the initial stages of very important chemical industries.

Some of these reactions are catalysed by d-metal oxides of various kinds. The physical chemistry of oxide surfaces is very complex, as can be appreciated by considering what happens during the oxidation of propene to acrolein on bismuth molybdate. The first stage is the adsorption of the propene molecule with loss of а hydrogen to form the allyl radical, СН₂=СНСН₃. An O atom in the surface can now transfer to this radical, leading to the formation of acrolein and its desorption from the surface. The Н atom also escapes with а surface O atom, and goes on to form Н₂О, which leaves the surface. The surface is left with vacancies and metal ions in lower oxidation states. These vacancies are attacked by О, molecules in the overlying gas, which then chemisorb as О^2- ions, so reforming the catalyst. This sequence of events involves great upheavals of the surface, and some materials break up under the stress.

Cracking and reforming. Many of the small organic molecules used in the preparation of all kinds of chemical products toте from oil. These small building blocks of polymers, perfumes, and petrochemicals in general, are usually cut from the long-chain hydrocarbons drawn from the Earth as petroleum. The catalytically induced fragmentation of the long-chain hydrocarbons is called cracking, and is often brought about on silica - alumina catalysts. These catalysts act by forming unstable carbocations, which dissociate and rearrange to more highly branched isomers. These branched isomers burn more smoothly and ef5ciently in internal combustion engines, and are used to produce higher octane fuels.

Catalytic reforming uses а dual-function catalyst, such as а dispersion of platinum and acidic alumina. The platinum provides the metal function, and brings about dehydrogenation and hydrogenation. The alumina provides the acidic function, being able to form carbocations from alkenes. The sequence of events in catalytic reforming shows up very clearly the complications that must be unravelled if а reaction as important as this is to be understood and improved. The first step is the attachment of the long-chain hydrocarbon by chemisorption to the platinum. In this process first one and then а second Н atom is lost, and an alkene is formed. The alkene migrates to а Brensted acid site, where it accepts а proton and attaches to the surface as а carbocation. This carbocation can undergo several different reactions. It can break into two, isomerize into а more highly branched form, or undergo varieties of ring-closure. Then it loses а proton, escapes &от the surface, and migrates (possibly through the gas) as an alkene to а metal part of the catalyst where it is hydrogenated. We end up with а rich selection of smaller molecules that can be withdrawn, fractionated, and then used as raw materials for other products.

INHIBITORS.

Inhibitors, once inappropriately called "negative catalysts," are substances which, when added to а reaction mixture, slow down the reaction. Inhibitors can act in а number of ways. One kind of inhibition occurs when the added substance combines with а potential catalyst, rendering it inactive and thus slowing the rate. For example, inhibition of а surface-catalyzed reaction can occur when foreign molecules bond at the active sites, blocking them from substrate molecules. Such inhibition is frequently called poisoning and the inhibitor, а poison.

ENZIMS

Incroduction. Life is inconceivable without enzymes. Most of the thousands of biochemical reactions that sustain living processes would occur at imperceptible rates in the absence of enzymes. The remarkable properties of enzymes include enormous catalytic power and а high degree of

One of the most important functions of proteins is their role as catalysts. Recall that living processes consist almost entirely of biochemical reactions. Without catalysts these reactions would not occur fast enough to sustain the living state.

То proceed at an acceptable rate, most chemical reactions require an initial input of energy. In the laboratory the energy required for reactions to proceed is usually supplied in the form of heat. Heating а reaction mixture increases the reaction rate for the following reason. At temperatures above absolute zero - 273.1⁰С), all molecules possess vibrational energy, which increases as the molecules are heated. Consider the following reaction: А + В = С

As the temperature rises, the likelihood of collisions between vibrating molecules (i.е., between А and В) increases. А chemical reaction occurs when the colliding molecules possess а minimum amount of energy called the activation energy. Not all collisions result in chemical reactions, because only а fraction of the molecules have sufficient energy to enter into the reaction (i.е., to break bonds or rearrange atoms into the product moIecuIes). Another way of increasing the likelihood of collisions, thereby increasing the formation of product, is to increase the concentration of the reactants.

In living systems the aforementioned strategies are not feasible. Elevated temperatures are harmful to delicate biological structures, and reactant concentrations are usually quite low. Living organisms circumvent these problems by using enzymes.

Enzymes have several remarkable properties. First, the rates of enzymatically catalyzed reactions are often phenomenally high. (Rate increases by factors of 10⁶ or greater are common.) Second, in marked contrast to inorganic catalysts the enzymes have а high degree of specificity with respect to the react ions they catalyze. The formation of side products is also гаге. Finally, because of their complex structures, enzymes are capable of being regulated. This is an especially important consideration in living organisms that must conserve energy and гаи materials.

Because enzymes are involved in so many aspects of living processes, any understanding of biochemistry depends on an appreciation of these remarkable catalysts.

Even in the presence of an inorganic catalyst, most laboratory reactions require an input of energy. In addition, most of these catalysts are nonspecific, that is, the accelerate а wide variety of reactions. Enzymes perform their work at mild temperatures and are quite specific in the reactions that each one catalyzes. The difference between inorganic catalysts and enzymes is directly related to their structures.in contrast to inorganic catalysts, each type of enzyme molecule contains а unique intricately shaped binding surface called an active site. Reactant molecules, called substrates, bind to the enzyme's active site, which is typically а small cleft or crevice on an otherwise large protein molecule. The active site is not just а binding site, however. Many of the amino acid side chains that line the active site actively participate in the catalytic process.

The lock-and-key model of enzyme action, originally introduced by Emil Fischer in 1890, accounts for enzyme specificity in the following way. Each enzyme binds to а single type of substrate because the active site and the substrate have complementary structures. The substrate's overall shape and charge distribution allow it to enter and interact with the enzyme's active site. In а modern variation by Daniel Koshland of the lock-and-key model, called the induced-fit model, the flexible structure of proteins is taken into account. In this model, substrate does not fit precisely into а rigid active site. Instead, noncovalent interactions between the enzyme and substrate cause а change in the three-dimensional structure of the active site. As а result of these interactions the shape of the active site conforms to the shape of the substrate.

Although the catalytic activity of some enzymes depends only on interactions between active site amino acids and the substrate, other enzymes require nonprotein components for their activities. Enzyme cofactors may be ions, such as Mg²⁺or Zn²⁺, or complex organic molecules, referred to as coenzymes. An enzyme that lacks an essential cofactor is called an apoenzyme. Intact enzymes with their bound cofactors are referred to as ho1oenzymes.

Some enzymes have another remarkable feature. Their activities can be regulated to an extraordinary extent. Regulation is necessary to the maintenance of a stable intracellular environment. For example, adjustments in the rates of enzymecatalyzed reactions allow cells to respond effectively to changes in the concentrations of various nutrients. Organisms use а variety of techniques to control enzyme activities. In some mechanisms, enzymes are regulated directly, principally through the binding of activators or inhibitors. Моте indirect methods involve the regulation of enzyme synthesis.

Classification of enzimes. In the early days of biochemistry, enzymes were named at the whim of their discoverers. Often, enzyme names provided по clue to their function (е.g., trypsin), or several names were used for the same enzyme. Enzymes were often named by adding the suffix "-ase" to the пате of the substrate. For example, urease catalyzes the hydrolysis of urea. To eliminate confusion, the International Union of Biochemistry (КВ) instituted а systematic naming scheme for enzymes. Each enzyme is now classified and named according to the type of chemical reaction it catalyzes. In this scheme an enzyme is assigned а four number classification and а two-part пате called а systematic паше. In addition, a shorter version of the systematic name, called the recommended name, is suggested by the IUB for everyday use. Because many enzymes were discovered before the institution of the systematic nomenclature, тапу of the old well-known names have been retained.

The following are the six major enzyme categories:

1. Oxidoreductases. Oxidoreductases catalyze various types of oxidation-reduction reactions. Subclasses of this group include the dehydrogenases, oxidases, oxygenases, reductases, peroxidases, and hydroxylases.

2. Transferases. Transferases catalyze reactions that involve the transfer of groups from one molecule to another. Examples of such groups include amino, carboxyl, carbonyl, methyl, phosphoryl, and acyl (RC=0). Common trivial names for the transferases often include the prefix "trans." Examples include the transcarboxylases, transmethylases, and transaminases.

3. Hydrolases. Hydrolases catalyze reactions in which the cleavage of bonds is accomplished by the addition of water. The hydrolases include the esterases, phosphatases, and peptidases.

4. Lyases. Lyases catalyze reactions in which groups (е.g., Н₂O, CO₂, and NH₃) are removed to form а double bond or added to а double bond. Decarboxylases, hydratases, dehydratases, deaminases, and synthases are examples of lyases.

5. Isomerases. This is а heterogeneous group of enzymes. lsomerases catalyze several types of intramolecular rearrangements. The epimerases catalyze the inversion of asymmetric carbon atoms. Mutases catalyze the intramolecular transfer of functional groups.

6. Ligases. Ligases catalyze bond formation between two substrate molecules. The energy for these reactions is always supplied by ATP hydrolysis. The names of many ligases include the term synthetase. Several other ligasesare called carboxylases.

Enzyme Inhibition.

The activity of enzymes can be inhibited. Study of the methods by which enzymes are inhibited have practical applications. For example, many clinical therapies and biochemical research tools are based on enzyme inhibition.

А variety of substances have the ability to reduce or eliminate the catalytic activity of specific enzymes. Inhibition may be irreversible or reversible. Irreversible inhibitors usually bond covalently to the enzyme, often to а side chain group in the active site.

Noncompetitive Inhibition.

In noncompetitive inhibition the inhibitor binds to the enzyme at а site other than the active site. Both ЕI and EIS complexes form. Inhibitor binding causes an alteration in the enzyme's three-dimensional configuration that prevents the reaction from occurring. For example, АМР is а noncompetitive inhibitor of fructose bisphosphate phosphatase, the enzyme that catalyzes the conversion of fructose-1,6-bisphosphate to fructose-6-phosphate. Noncompetitive inhibition is not reversed by increasing the concentration of substrate.

Catalysis.

However valuable kinetic studies are, they reveal little about how enzymes catalyze biochemical reactions. Biochemists use а variety of other techniques to investigate the catalytic mechanisms of enzymes. (А mechanism is а description of the specific steps that occur as а chemical reaction takes place.) The goal of enzyme mechanism investigations is to relate enzyme activity to the structure and function of the active site. Methods that are used to provide insight into catalytic mechanisms include Х-ray crystallography, chemical inactivation of active site side chains, and studies using simple model compounds as substrates and as inhibitors.

Catalytic Mechanisms.

Despite extensive research, the mechanisms of only а few enzymes are known in significant detail. However, it has become increasingly clear that enzymes utilize the same catalytic mechanisms as nonenzymatic catalysts. The significantly higher catalytic rates that enzymes achieve are largely Же to the fact that their active sites possess structures that are uniquely suited to promote catalysis.

Several factors contribute to enzyme catalysis. The most important of these are (1) proximity and strain effects, (2) electrostatic effects, (3) acid base catalysis, and (4) covalent catalysis. Each factor will be described briefly.

Proximity and Strain Effects. For а biochemical reaction to occur, the substrate must come into close proximity to catalytic functional groups (side chain groups involved in а catalytic mechanism) within the active site. In addition, the substrate must be precisely oriented in relation to the catalytic groups. Once the substrate is correctly positioned, а change in the enzyme's conformation may result in а strained enzyme-substrate complex. This strain helps to bring the enzyme-substrate complex into the transition state. In general, the more tightly the active site is able to bind the substrate while it is in its transition state, the greater the rate of the reaction.

Electrostatic Effects.

Recall that the strength of electrostatic interactions is related to the capacity of surrounding solvent molecules to reduce the attractive forces between chemical groups. Because water is largely excluded from the active site as substrate binds, the local dielectric constant is often low. The charge distribution in the relatively anhydrous active site may influence the chemical reactivity of the substrate. In addition, weak electrostatic interactions, such as those between permanent and induced dipoles in both the active site and the substrate, are believed to contribute to catalysis. А more efficient binding of substrate causes а lowering in the free energy of the transition state, which results in an acceleration of the reaction.

Acid-Base Catalysis.

Chemical groups can often be made more reactive by the addition or removal of а proton. Enzyme active sites contain side chain groups that act as proton donors or acceptors. Transfers of protons are а common feature of chemical reactions. For example, consider the hydrolysis of an ester: Because water is а weak nucleophile, ester hydrolysis is relatively slow in neutral solution. Ester hydrolysis takes place much more rapidly if the pH is raised. As hydroxide ion attacks the polarized carbon atom of the carbonyl group, and а tetrahedral intermediate is formed. As the intermediate breaks down, а proton is transferred from а nearby water molecule. The reaction is complete when the alcohol is released. However, hydroxide ion catalysis is not practical in living systems. Enzymes use several functional groups that behave as general bases to aid in the efficient transfer of protons. Such groups can be precisely positioned in relation to the substrat. Ester hydrolysis can also be catalyzed by а general acid. As theoxygen of the ester’s carbonyl group binds to the proton, the carbon atom becomes more positive. The ester then becomes тоге susceptible to the nucleophilic attack of а water molecule.

Because such groups are only weakly ionizable, they are referred to as general acids or general bases. (The terms general acid and general base refer to substances that are capable of releasing а proton or accepting а proton, respectively. Enzymes almost always use general acids or general bases in preference to protons or hydroxide groups. For the sake of simplicity, however, the symbols Н⁺ and ОН^- are often used in illustrations of reaction mechanisms.) For example, the side chain of histidine (referred to as an imidazole group) often participates in catalytic mechanisms. It does so because its рК, is approximately 6. Therefore the histidine side chain ionizes within the hysiological pH range. The protonated form of histidine is а general acid. Once it loses its proton (and becomes а conjugate base), histidine is а general base.

Covalent Catalysis.

In some enzymes а nucleophilic side chain group forms an unstable covalent bond with the substrate. The enzyme-substrate complex then undergoes further reaction to form product. А class of enzymes called the serine proteases use the - СН₂ - ОН group of serine as а nucleophile to hydrolyze peptide bonds. (Examples of the serine proteases include the digestive enzymes trypsin and chymotrypsin and the blood- clotting enzyme thrombin.) During the first step, the nucleophile attacks the carbonyl group. As the ester bond is formed, the peptide bond is broken. The resulting highly reactive intermediate is hydrolyzed in а second reaction by water.

Several other amino acid side chains may act as nucleophiles. The sulfhydryl group of cysteine, the carboxylate groups of aspartate and glutamate, and the imidazole group of histidine can play this role.

Irreversible inhibitors usually bind covalently to enzymes. In reversible inhibition the inhibitor can dissociate from the enzyme. The most common types of reversible inhibition are competitive and noncompetitive. The kinetic properties of allosteric enzymes are not explained by the Michaelis-Menten model. Most allosteric enzymes are composed of subunits called protomers. The binding of substrate or effector to one protomer affects the binding properties of other protomers. Enzymes use the same catalytic mechanisms as nonenzy- matic catalysts. Several factors contribute to enzyme catalysts: proximity and strain effects, electrostatic effects, acid-base catalysis, and covalent catalysis. Each enzyme mechanism results from the simultaneous use of various combinations of these factors.

Enzymes are biological catalysts. They enhance reaction rates because they provide an alternative reaction pathway that re quires less energy than an uncatalyzed reaction. In contrast to some inorganic catalysts, most enzymes catalyze reactions at mild temperatures. In addition, enzymes are specific in regard to the types of reactions they catalyze. Each type of enzyme contains а unique, intricately shaped binding surface called an active site. Substrate binds to the enzyme's active site, which is а small cleft or crevice in an otherwise large protein molecule. In the look-and-key model of enzyme action the structures of the enzyme's active site and the substrate are complementary. In the induced fit model the protein molecule is assumed to be flexible.

Each enzyme is currently classified and named according to the type of reaction it catalyzes. There are six major enzyme categories: oxidoreductases, transferases, hydrolases, lyases, isomerases, and ligases.

Enzyme inhibition may be reversible or irreversible. Active site amino acid side chains are primarily responsible for catalyzing proton transfers and nucleophilic substitutions. Nonprotein cofactors (metals and coenzymes) are used by enzymes to catalyze other types of reactions.

Enzymes are sensitive to environmental factors such as temperature and pH. Each enzyme has an optimum temperature and an optimum pH.

The chemical reactions in living cells are organized into а series of biochemical pathways. Control of biochemical path-ways is achieved primarily by adjusting the concentrations andactivities of enzymes. This control is accomplished by utilizing various combinations of the following mechanisms: genetic control, covalent modification, allosteric regulation, and compartmentation.