The materials to prepare students for practical lessons of inorganic chemistry
LESSON № 7.
THEME. The speed and mechanisms of the chemical reactions. Catalysis. Chemical equilibrium
1. Chemical kinetics. Reaction rate. Rate Law of chemical reaction. Concentration of reactants dependence rate.
2. Order and molecularity of reaction
3. The temperature dependence rate. The vant’-Hoff’s rule. Arrhenius’ equation.
4. Chemical equilibrium. The law of chemical equilibrium.
5. Features of action of catalysts. Homogeneous, heterogeneous and microheterogeneous catalysis. Acid-basic catalysis. Autocatalysis. Mechanism of action of catalysts.
6. Enzymes as biological catalysts.
1. CHEMICAL KINETICS. REACTION RATE. CHEMICAL KINETICS. REACTION RATE. RATE LAW OF CHEMICAL REACTION. CONCENTRATION OF REACTANTS DEPENDENCE RATE.
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.
Chemical kinetics is that branch of chemistry which deals with the study of the speeds r the rates f chemical reactions, the factors affecting the rates of the reactions and the mechanism by which the reactions proceed.
The general outcome of such experiments is the observation that reaction rates depend on the composition and the temperature of the reaction mixture. The next few sections look at these observations in more detail.
Consider а reaction of the form: A + В ® С
in which at some instant the concentrations of the participants are [А], [В], and [С]. Measures of the rate of the reaction are the rate of formation of а product (С) and the rate of consumption of one of thy reactants (А or В). The rate of consumption of the reactant А is
the rate of formation of the product С is:
Both rates are positive. In the present case, the reaction stoichiometry implies that the rate of formation of С is equal to the rate of consumption of either А or В, because, whenever а С molecule is formed, one А molecule and one В molecule are destroyed. For а reaction with а more complicated stoichiometry, such as А + 2В ® 3С + D, the relation between the various rates of formation and consumption is more complicated.
More specifically and completely,
The ambiguity in the definition of rate is avoided if we define the rate of reaction u as
where vJ is the stoichiometric coefficient of substance J. Now there is а single rate for the entire equation.
The rate of chemical reaction is the change in the concentration of any one of the reactants or products per unit of time.
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:
By doing experiments involving a reaction between A and B, you would find that the rate of the reaction was related to the concentrations of A and B in this way:
This is called the rate equation for the reaction.
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.
The rate constant
Surprisingly, the rate constant isn’t actually a true constant! It varies, for example, if you change the temperature of the reaction, add a catalyst, or change the catalyst.
The rate constant is constant for a given reaction only if all you are changing is the concentration of the reactants. You will find more about the effect of temperature and catalysts on the rate constant on another page.
Note: If you want to follow up this further look at rate constants you might like to follow this link. Alternatively, you could visit it later via the rates of reaction menu.
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.
2. ORDER AND MOLECULARITY OF REACTION
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
First-order reactions: As shown in the Justification that follows, the first-order rate law for the consumption of а reactant А:
has the solution:
These two equations are versions of an integrated rate law, the integrated form of the rate equation.
Half-lives. А useful indication of the rate of а first-order chemical reaction is the half-life, of а substance, the tune it takes for its concentration to fall to half the initial value. The time for [А] to decrease from [А]0 to ½ [А]0 in а first-order reaction is given
Hence
(It is sometimes useful to remember that ln 2 = 0.693.) The main point to note about this result is that, for а first-order reaction, the half-life of а reactant is independent of its initial concentration. Hence, if the concentration of А at some arbitrary stage of the reaction is [А], then it will have fallen to ½[А] after а further interval of (ln 2)/k.
Second-order reactions. We show in the Justification that follows that the integrated form of the second-order rate law
is
Equation shows that to test for а second-order reaction we should plot 1/[А] against t and expect а straight line. The slope of the graph is k. Equation lets us predict the concentration of А at any tune after the start of the reaction. It shows that the concentration of А approaches zero more slowly than in а first-order reaction with the same initial rate:
that the half-life of а species А that is consumed in а second-order reaction is proportional to 1/[А]о.
Another type of second-order reaction is one that is first-order in each of two reactants А and В:
А+ В = Products
and the initial concentrations are [А]0 and [В]0, then it is shown in the following Justification that, at а time r after the start of the reaction, the concentrations satisfy the relation
Third-order reactions. An example will help to show how а pre-equilibrium assumption helps to elucidate а mechanism. The oxidation of nitrogen(II) oxide is found to be third-order overall;
А mechanism that accounts for the rate law and the temperature dependence is а pre-equilibrium:
followed by а simple bimolecular reaction
The reaction rate is obtained by combining the two equations into
which has the observed overall third-order form.
Some examples
Each of these examples involves a reaction between A and B, and each rate equation comes from doing some experiments to find out how the concentrations of A and B affect the rate of reaction.
Example 1:
In this case, the order of reaction with respect to both A and B is 1. The overall order of reaction is 2 – found by adding up the individual orders.
Note: Where the order is 1 with respect to one of the reactants, the “1” isn’t written into the equation. [A] means [A]1.
Example 2.
rate = k [B]2
This reaction is zero order with respect to A because the concentration of A doesn’t affect the rate of the reaction. The order with respect to B is 2 – it’s a second order reaction with respect to B. The reaction is also second order overall (because 0 + 2 = 2).
Example 3.
rate = k [A]
This reaction is first order with respect to A and zero order with respect to B, because the concentration of B doesn’t affect the rate of the reaction. The reaction is first order overall (because 1 + 0 = 1).
What if you have some other number of reactants?
It doesn’t matter how many reactants there are. The concentration of each reactant will occur in the rate equation, raised to some power. Those powers are the individual orders of reaction. The overall order of the reaction is found by adding them all up.
molecularity
Most reactions occur in а sequence of steps called elementary reactions, each of which involves only one or two molecules or ions. A typical elementary reaction is
Н+ Br2 ® HBr + Br
(We do not specify the phase of the species when giving the chemical equation for an elementary reaction.) This equation signifies that an H atom attacks а Br2 molecule to produce an HBr molecule and а Br atom.
The number of atoms, ions or molecules that must collide with one another simultaneously so as to result into а chemical reaction и called the molecularity of the reaction. Or other: the molecularity of an elementary reaction is the number of molecules coming together to react.
In case of simple reactions (also called elementary reactions), the molecularity is simply the sum of the molecules of the different reactants as represented by the balanced chemical equation. А few examples are given below:
(i) Decomposition of F2О2: The balanced equation is
F2О2 = F2 + O2
Hence the molecularity of the reactions 1 and the reaction is called Unimolecular.
(ii) Dissociation of HI: The balanced equation is
2НI = H2 + I2
Hence the molecularity is 2 and the reaction is called Bimolecular.
(iii) Reaction between NO and O2. The balanced equation is
2NO + О2 = 2NO2
Неnсе the molecularity is 3 and the reaction is called Тermolecular.
In case of complex reactions. i.e. reactions involving а number of atoms, ions or molecules of the reactants in the balanced equation (usually more than 3), the chances for all the atoms, ions or molecules of the reactants to come together and collide are чету гаге. Hence in such cases, the reactions are supposed to take place in а number of steps.
The slowest step is the rate determining step.
The situation is similar to that of а cycle factory. The different sections of the factory manufacture different parts like handles, paddies, rims, seats etc. The overall number of cycles manufactured per day depends upon the slowest working section of the factory.
The number of atoms, ion оr molecules taking part in the slowest step i.е. the rate determining step is called the molecularity of the reaction.
Thus it is obvious that the molecularity of а reaction must always be а whole number whereas the order of a reaction can be fractional also (as already mentioned). However in most of the reactions, the molecularity of the reaction is found to be same as the order. Hence the two are used without any distinction.
А series of step reactions or elementary processes proposed to account for the overall reaction is called the mechanism of the reaction.
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).
3. THE TEMPERATURE DEPENDENCE RATE. THE VANT’-HOFF’S RULE. ARRHENIUS’ EQUATION.
The Arrhenius parameters.
van’t Hoff’s rule the velocity of chemical reactions is increased twofold or more for each rise of 10°C in temperature; generally true only when temperatures approximate those normal for the reaction.
By 1890 it was common knowledge that higher temperatures speed up reactions, often doubling the rate for a 10-degree rise, but the reasons for this were not clear. Finally, in 1899, the Swedish chemist Svante Arrhenius (1859-1927) combined the concepts of activation energy and the Boltzmann disribution law into one of the most important relationships in physical chemistry.
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 Arrhenius equation is often written as
The factor А is called the pre-exponential factor or the frequency factor; Еa 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 Еa 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. However, it is still possible to express the strength of the dependence by defining an activation energy as
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:
Determining the Activation Energy of a Reaction
The rate of a reaction depends on the temperature at which it is run. As the temperature increases, the molecules move faster and therefore collide more frequently. The molecules also carry more kinetic energy. Thus, the proportion of collisions that can overcome the activation energy for the reaction increases with temperature.
The only way to explain the relationship between temperature and the rate of a reaction is to assume that the rate constant depends on the temperature at which the reaction is run. In 1889, Svante Arrhenius showed that the relationship between temperature and the rate constant for a reaction obeyed the following equation.
In this equation, k is the rate constant for the reaction, Z is a proportionality constant that varies from one reaction to another, Ea is the activation energy for the reaction, R is the ideal gas constant in joules per mole kelvin, and T is the temperature in kelvin.
The Arrhenius equation can be used to determine the activation energy for a reaction. We start by taking the natural logarithm of both sides of the equation.
We then rearrange this equation to fit the equation for a straight line.
y = mx + b
According to this equation, a plot of ln k versus 1/T should give a straight line with a slope of – Ea/R, as shown in the figure below.
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Practice Problem 11: Use the following data to determine the activation energy for the decomposition of HI: Temperature (K) Rate Constant (M/s) 573 2.91 x 10-6 673 8.38 x 10-4 773 7.65 x 10-2 |
By paying careful attention to the mathematics of logarithms, it is possible to derive another form of the Arrhenius equation that can be used to predict the effect of a change in temperature on the rate constant for a reaction.
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Practice Problem: Calculate the rate of decomposition of HI at 600°C. |
The Arrhenius equation can also be used to calculate what happens to the rate of a reaction when a catalyst lowers the activation energy.
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 shows the uncatalyzed path of а reaction contrasted with its catalyzed path.
4. CHEMICAL EQUILIBRIUM. THE LAW OF CHEMICAL EQUILIBRIUM.
The Law of Chemical Equilibrium
Two examples of an equilibrium constant dealt with in other sections, namely,
for the reaction cis-C2H2F2 = trans-C2H2F2
and
for the reaction N2O4 = 2NO2
are both particular examples of a more general law governing chemical equilibrium ingases. If we write an equation for a gaseous equilibrium in general in the form
aA(g) + bB(g)
then the equilibrium constant defined by the equation
is found to be a constant quantity depending only on the temperature and the nature of the reaction. This general result is called the law of chemical equilibrium, or thelaw of mass action.
EXAMPLE 1 Write expressions for the equilibrium constant for the following reactions:
a) 2HI(g) = H2(g) + I2(g)
b) N2(g) + 3H2(g) = 2NH3(g)
c) O2(g) + 4HCl(g) = 2H2O(g) + 2Cl2(g)
EXAMPLE 2 A mixture containing equal concentrations of methane and steam is passed over a nickel catalyst at 1000 K. The emerging gas has the composition [CO] = 0.1027 mol dm–3, [H2] = 0.3080 mol dm–3, and [CH4] = [H2O] = 0.8973 mol dm–3. Assuming this mixture is at equilibrium, calculate the equilibrium constant Kc for the reaction
CH4(g) + H2O(g)
Solution The equilibrium constant is given by the following equation:
5. FEATURES OF ACTION OF CATALYSTS. HOMOGENEOUS, HETEROGENEOUS AND MICROHETEROGENEOUS CATALYSIS. ACID-BASIC CATALYSIS. AUTOCATALYSIS. MECHANISM OF ACTION OF CATALYSTS
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.
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.
Types of catalytic reactions
Catalysts can be divided into two main types – heterogeneous and homogeneous.
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).
HETEROGENEOUS CATALYSIS
What is a phase?
If you look at a mixture and can see a boundary between two of the components, those substances are in different phases. A mixture containing a solid and a liquid consists of two phases. A mixture of various chemicals in a single solution consists of only one phase, because you can’t see any boundary between them.
You might wonder why phase differs from the term physical state (solid, liquid or gas). It includes solids, liquids and gases, but is actually a bit more general. It can also apply to two liquids (oil and water, for example) which don’t dissolve in each other. You could see the boundary between the two liquids.
If you want to be fussy about things, the diagrams actually show more phases than are labelled. Each, for example, also has the glass beaker as a solid phase. All probably have a gas above the liquid – that’s another phase. We don’t count these extra phases because they aren’t a part of the reaction.
How the heterogeneous catalyst works (in general terms)
Most examples of heterogeneous catalysis go through the same stages:
· One or more of the reactants are adsorbed on to the surface of the catalyst at active sites. Adsorption is where something sticks to a surface. It isn’t the same as absorption where one substance is taken up within the structure of another. Be careful! An active site is a part of the surface which is particularly good at adsorbing things and helping them to react.
· There is some sort of interaction between the surface of the catalyst and the reactant molecules which makes them more reactive. This might involve an actual reaction with the surface, or some weakening of the bonds in the attached molecules.
· The reaction happens. At this stage, both of the reactant molecules might be attached to the surface, or one might be attached and hit by the other one moving freely in the gas or liquid.
· The product molecules are desorbed. Desorption simply means that the product molecules break away. This leaves the active site available for a new set of molecules to attach to and react.
A good catalyst needs to adsorb the reactant molecules strongly enough for them to react, but not so strongly that the product molecules stick more or less permanently to the surface.
Examples of heterogeneous catalysis
· Silver, for example, isn’t a good catalyst because it doesn’t form strong enough attachments with reactant molecules. Tungsten, on the other hand, isn’t a good catalyst because it adsorbs too strongly.
· Metals like platinum and nickel make good catalysts because they adsorb strongly enough to hold and activate the reactants, but not so strongly that the products can’t break away.
· The hydrogenation of a carbon-carbon double bond
· The simplest example of this is the reaction between ethene and hydrogen in the presence of a nickel catalyst.
In practice, this is a pointless reaction, because you are converting the extremely useful ethene into the relatively useless ethane. However, the same reaction will happen with any compound containing a carbon-carbon double bond.
One important industrial use is in the hydrogenation of vegetable oils to make margarine, which also involves reacting a carbon-carbon double bond in the vegetable oil with hydrogen in the presence of a nickel catalyst.
Ethene molecules are adsorbed on the surface of the nickel. The double bond between the carbon atoms breaks and the electrons are used to bond it to the nickel surface.
Hydrogen molecules are also adsorbed on to the surface of the nickel. When this happens, the hydrogen molecules are broken into atoms. These can move around on the surface of the nickel. If a hydrogen atom diffuses close to one of the bonded carbons, the bond between the carbon and the nickel is replaced by one between the carbon and hydrogen.
That end of the original ethene now breaks free of the surface, and eventually the same thing will happen at the other end.
As before, one of the hydrogen atoms forms a bond with the carbon, and that end also breaks free. There is now space on the surface of the nickel for new reactant molecules to go through the whole process again.
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)
The reaction is believed to proceed through the following preequilibrium:
(In the equilibrium constant, the activity of H2O has been set equal to 1.) Because the second step is rate-determining, the rate law of the overall reaction is predicted to be:
in agreement with the observed dependence of the rate on the Br– concentration and the pH of the solution. The observed activation energy is that of the effective rate coefficient kK. In the absence of Br– ions the reaction cannot proceed through the path set out above, and а different and much higher activation energy is observed.
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 converters
Catalytic converters change poisonous molecules like carbon monoxide and various nitrogen oxides in car exhausts into more harmless molecules like carbon dioxide and nitrogen. They use expensive metals like platinum, palladium and rhodium as the heterogeneous catalyst.
The metals are deposited as thin layers onto a ceramic honeycomb. This maximises the surface area and keeps the amount of metal used to a minimum.
Taking the reaction between carbon monoxide and nitrogen monoxide as typical:
In the same sort of way as the previous example, the carbon monoxide and nitrogen monoxide will be adsorbed on the surface of the catalyst, where they react. The carbon dioxide and nitrogen are then desorbed.
The use of vanadium(V) oxide in the Contact Process
During the Contact Process for manufacturing sulphuric acid, sulphur dioxide has to be converted into sulphur trioxide. This is done by passing sulphur dioxide and oxygen over a solid vanadium(V) oxide catalyst.
This example is slightly different from the previous ones because the gases actually react with the surface of the catalyst, temporarily changing it. It is a good example of the ability of transition metals and their compounds to act as catalysts because of their ability to change their oxidation state. The sulphur dioxide is oxidised to sulphur trioxide by the vanadium(V) oxide. In the process, the vanadium(V) oxide is reduced to vanadium(IV) oxide.
The vanadium(IV) oxide is then re-oxidised by the oxygen.
This is a good example of the way that a catalyst can be changed during the course of a reaction. At the end of the reaction, though, it will be chemically the same as it started.
Homogeneous catalysis
This has the catalyst in the same phase as the reactants. Typically everything will be present as a gas or contained in a single liquid phase. The examples contain one of each of these
This involves the use of a catalyst in a different phase from the reactants. 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]
Typical examples involve a solid catalyst with the reactants as either liquids or gases.
Examples of homogeneous catalysis
The reaction between persulphate ions and iodide ions.
This is a solution reaction that you may well only meet in the context of catalysis, but it is a lovely example!
Persulphate ions (peroxodisulphate ions), S2O82-, are very powerful oxidising agents. Iodide ions are very easily oxidised to iodine. And yet the reaction between them in solution in water is very slow.
The reactioeeds a collision between two negative ions. Repulsion is going to get seriously in the way of that!
The catalysed reaction avoids that problem completely. The catalyst can be either iron(II) or iron(III) ions which are added to the same solution. This is another good example of the use of transition metal compounds as catalysts because of their ability to change oxidation state.
For the sake of argument, we’ll take the catalyst to be iron(II) ions. As you will see shortly, it doesn’t actually matter whether you use iron(II) or iron(III) ions.
The persulphate ions oxidise the iron(II) ions to iron(III) ions. In the process the persulphate ions are reduced to sulphate ions. The iron(III) ions are strong enough oxidising agents to oxidise iodide ions to iodine. In the process, they are reduced back to iron(II) ions again.
Both of these individual stages in the overall reaction involve collision between positive and negative ions. This will be much more likely to be successful than collision between two negative ions in the uncatalysed reaction. What happens if you use iron(III) ions as the catalyst instead of iron(II) ions? The reactions simply happen in a different order.
The destruction of atmospheric ozone. This is a good example of homogeneous catalysis where everything is present as a gas.
Ozone, O3, is constantly being formed and broken up again in the high atmosphere by the action of ultraviolet light. Ordinary oxygen molecules absorb ultraviolet light and break into individual oxygen atoms. These have unpaired electrons, and are known as free radicals. They are very reactive.
The oxygen radicals can then combine with ordinary oxygen molecules to make ozone. Ozone can also be split up again into ordinary oxygen and an oxygen radical by absorbing ultraviolet light.
This formation and breaking up of ozone is going on all the time. Taken together, these reactions stop a lot of harmful ultraviolet radiation penetrating the atmosphere to reach the surface of the Earth.
The catalytic reaction we are interested in destroys the ozone and so stops it absorbing UV in this way.
Chlorofluorocarbons (CFCs) like CF2Cl2, for example, were used extensively in aerosols and as refrigerants. Their slow breakdown in the atmosphere produces chlorine atoms – chlorine free radicals. These catalyse the destruction of the ozone.
This happens in two stages. In the first, the ozone is broken up and a new free radical is produced. The chlorine radical catalyst is regenerated by a second reaction. This can happen in two ways depending on whether the ClO radical hits an ozone molecule or an oxygen radical. If it hits an oxygen radical (produced from one of the reactions we’ve looked at previously). Or if it hits an ozone molecule
Because the chlorine radical keeps on being regenerated, each one can destroy thousands of ozone molecules.
AUTOCATALYSIS
The oxidation of ethanedioic acid by manganate(VII) ions.
In autocatalysis, the reaction is catalysed by one of its products. One of the simplest examples of this is in the oxidation of a solution of ethanedioic acid (oxalic acid) by an acidified solution of potassium manganate(VII) (potassium permanganate).
The reaction is very slow at room temperature. It is used as a titration to find the concentration of potassium manganate(VII) solution and is usually carried out at a temperature of about 60°C. Even so, it is quite slow to start with.
The reaction is catalysed by manganese(II) ions. There obviously aren’t any of those present before the reaction starts, and so it starts off extremely slowly at room temperature. However, if you look at the equation, you will find manganese(II) ions amongst the products. More and more catalyst is produced as the reaction proceeds and so the reaction speeds up.
You can measure this effect by plotting the concentration of one of the reactants as time goes on. You get a graph quite unlike the normal rate curve for a reaction.
Concentrations are high at the beginning and so the reaction is fast – shown by a rapid fall in the reactant concentration. As things get used up, the reaction slows down and eventually stops as one or more of the reactants are completely used up.
You can see the slow (uncatalysed) reaction at the beginning. As catalyst begins to be formed in the mixture, the reaction speeds up – getting faster and faster as more and more catalyst is formed. Eventually, of course, the rate falls again as things get used up.
Don’t assume that a rate curve which looks like this necessarily shows an example of autocatalysis. There are other effects which might produce a similar graph. For example, if the reaction involved a solid reacting with a liquid, there might be some sort of surface coating on the solid which the liquid has to penetrate before the expected reaction can happen.
A more common possibility is that you have a strongly exothermic reaction and aren’t controlling the temperature properly. The heat evolved during the reaction speeds the reaction up.
Examples of autocatalytic reactions
· Haloform reaction
· Tin pest
· Reaction of Permanganate with Oxalic Acid
· The mechanism of the above reaction
· Vinegar syndrome
· Binding of oxygen by hemoglobin
· The spontaneous degradation of aspirin into salicylic acid and acetic acid, causing very old aspirin in sealed containers to smell mildly of vinegar.
· The α-bromination of acetophenone with bromine.
Involvement in life processes
Autocatalysis plays a major role in the processes of life. Two researchers who have emphasised its role in the origins of life are Robert Ulanowicz and Stuart Kauffman.
Autocatalysis occurs in the initial transcripts of rRNA. The introns are capable of excising themselves by the process of two nucleophilic transesterification reactions. The RNA able to do this is sometimes referred to as a ribozyme. Additionally, the citric acid cycle is an autocatalytic cycle run in reverse.
Ultimately, biological metabolism itself can be seen as a vast autocatalytic set, in that all of the molecular constituents of a biological cell are produced by reactions involving this same set of molecules.
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.
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.
How to determine Reaction Rates from Balanced Equations
Using the equation
aA+bB -> cC +dD
where the lower case letters (i.e. a) represent the coefficient of the balanced equation and the upper case letters (i.e. A) represent the molecule.
It is seen that the Rate of Disappearance are: -∆[A]/∆t*1/a=-∆[B]/∆t*1/b
and the Rate of Formation are: ∆[C]/∆t*1/c=∆[D]/∆t*1/d
Since Rate of Disappearance and Rate of Formation are equal to one another
-∆[A]/∆t*1/a=-∆[B]/∆t*1/b=∆[C]/∆t*1/c=∆[D]/∆t*1/d
Measuring Reagents Versus Product
It does not matter whether an experimenter monitors the reagents or products because there is no effect on the overall reaction. However, since reagents decrease during reaction, and products increase, there is a sign difference between the two rates. Reagent concentration decreases as the reaction proceeds, giving a negative number for the change in concentration. The products, on the other hand, increase concentration with time, giving a positive number. Since he convention is to express the rate of reaction as a positive number, to solve a problem, set the overall rate of the reaction equal to the negative of a reagent’s disappearing rate. The overall rate also depends on stoichiometric coefficients.
It is worth noting that the process of measuring the concentration can be greatly simplified by taking advantage of the different physical or chemical properties (i.e.: phase difference, reduction potential, etc.) of the reagents or products involved in the reaction by using the above methods. We have emphasized the importance of taking the sign of the reaction into account in order to get a positive reaction rate. Now, we will turn our attention to the importance of stoichiometric coefficients.
Unique Average Rate of Reaction
A reaction rate can be reported quite differently depending on which product or reagent selected to be monitored.
Given a reaction:
rate of reaction =
This formula can also be written as:
rate of reaction =
=
=
=
Even though the concentrations of A, B, C and D may all change at different rates, there is only one average rate of reaction. To get this unique rate, choose any one rate and divide it by the stoichiometric coefficient. When the reaction has the formula:
The general case of the unique average rate of reaction has the form:
rate of reaction =
|
Example |
|
For the reaction,
Solutions:
a)rate of reaction =
b)rate of disappearance of A =
rate of reaction =
Average and Instantaneous Reaction Rate
Reaction rates have the general form of (change of concentration / change of time). There are two types of reaction rates. One is called the average rate of reaction, often denoted by (Δ[conc.] / Δt), while the other is referred to as the instantaneous rate of reaction, denoted as either:
The average rate of reaction, as the name suggests, is an average rate, obtained by taking the change in concentration over a time period, for example: -0.3 M / 15 minutes. This is an approximation of the reaction rate in the interval; it does not necessarily mean that the reaction has this specific rate throughout the time interval or even at any instant during that time. The instantaneous rate of reaction, on the other hand, depicts a more accurate value. The instantaneous rate of reaction is defined as the change in concentration of an infinitely small time interval, expressed as the limit or derivative expression above. Instantaneous rate can be obtained from the experimental data by first graphing the concentration of a system as function of time, and then finding the slope of the tangent line at a specific point which corresponds to a time of interest. Alternatively, experimenters can measure the change in concentration over a very small time period two or more times to get an average rate close to that of the instantaneous rate. The reaction rate for that time is determined from the slope of the tangent lines.
Example
From the graph, the blue line is the graphical representation of [A] over t. The red and green lines are tangent lines to the graph. Find the reaction rate at a) t=195 sec and b) t=395 sec.
Solutions:
To find the slopes, divide the y-intercept by the x-intercept.
1. The red line has a slope of
2. The green line has a slope of
The Activation Energy of Chemical Reactions
Only a small fraction of the collisions between reactant molecules convert the reactants into the products of the reaction. This can be understood by turning, once again, to the reaction between ClNO2 and NO.
ClNO2(g) + NO(g)
In the course of this reaction, a chlorine atom is transferred from one nitrogen atom to another. In order for the reaction to occur, the nitrogen atom in NO must collide with the chlorine atom in ClNO2.
Reaction won’t occur if the oxygen end of the NO molecule collides with the chlorine atom on ClNO2.
Nor will it occur if one of the oxygen atoms on ClNO2 collides with the nitrogen atom on NO.
Another factor that influences whether reaction will occur is the energy the molecules carry when they collide. Not all of the molecules have the same kinetic energy, as shown in the figure below. This is important because the kinetic energy molecules carry when they collide is the principal source of the energy that must be invested in a reaction to get it started.
The overall standard free energy for the reaction between ClNO2 and NO is favorable.
|
ClNO2(g) + NO(g) |
|
But, before the reactants can be converted into products, the free energy of the system must overcome the activation energy for the reaction, as shown in the figure below. The vertical axis in this diagram represents the free energy of a pair of molecules as a chlorine atom is transferred from one to the other. The horizontal axis represents the the sequence of infinitesimally small changes that must occur to convert the reactants into the products of this reaction.
To understand why reactions have an activation energy, consider what has to happen in order for ClNO2 to react with NO. First, and foremost, these two molecules have to collide, thereby organizing the system. Not only do they have to be brought together, they have to be held in exactly the right orientation relative to each other to ensure that reaction can occur. Both of these factors raise the free energy of the system by lowering the entropy. Some energy also must be invested to begin breaking the Cl-NO2 bond so that the Cl-NO bond can form.
NO and ClNO2 molecules that collide in the correct orientation, with enough kinetic energy to climb the activation energy barrier, can react to form NO2 and ClNO. As the temperature of the system increases, the number of molecules that carry enough energy to react when they collide also increases. The rate of reaction therefore increases with temperature. As a rule, the rate of a reaction doubles for every 10oC increase in the temperature of the system.
Purists might note that the symbol used to represent the difference between the free energies of the products and the reactants in the above figure is
Purists might also note that the symbol used to represent the activation energy is written with a capital “E“. This is unfortunate, because it leads students to believe the activation energy is the change in the internal energy of the system, which is not quite true. Ea measures the change in the potential energy of a pair of molecules that is required to begin the process of converting a pair of reactant molecules into a pair of product molecules.
Catalysts and the Rates of Chemical Reactions
Aqueous solutions of hydrogen peroxide are stable until we add a small quantity of the I– ion, a piece of platinum metal, a few drops of blood, or a freshly cut slice of turnip, at which point the hydrogen peroxide rapidly decomposes.
2 H2O2(aq)
This reaction therefore provides the basis for understanding the effect of a catalyst on the rate of a chemical reaction. Four criteria must be satisfied in order for something to be classified as catalyst.
· Catalysts increase the rate of reaction.
· Catalysts are not consumed by the reaction.
· A small quantity of catalyst should be able to affect the rate of reaction for a large amount of reactant.
· Catalysts do not change the equilibrium constant for the reaction.
The first criterion provides the basis for defining a catalyst as something that increases the rate of a reaction. The second reflects the fact that anything consumed in the reaction is a reactant, not a catalyst. The third criterion is a consequence of the second; because catalysts are not consumed in the reaction, they can catalyze the reaction over and over again. The fourth criterion results from the fact that catalysts speed up the rates of the forward and reverse reactions equally, so the equilibrium constant for the reaction remains the same.
Catalysts increase the rates of reactions by providing a new mechanism that has a smaller activation energy, as shown in the figure below. A larger proportion of the collisions that occur between reactants now have enough energy to overcome the activation energy for the reaction. As a result, the rate of reaction increases.
To illustrate how a catalyst can decrease the activation energy for a reaction by providing another pathway for the reaction, let’s look at the mechanism for the decomposition of hydrogen peroxide catalyzed by the I– ion. In the presence of this ion, the decomposition of H2O2 doesn’t have to occur in a single step. It can occur in two steps, both of which are easier and therefore faster. In the first step, the I– ion is oxidized by H2O2 to form the hypoiodite ion, OI–.
H2O2(aq) + I–(aq)
In the second step, the OI– ion is reduced to I– by H2O2.
OI–(aq) + H2O2(aq)
Because there is no net change in the concentration of the I– ion as a result of these reactions, the I– ion satisfies the criteria for a catalyst. Because H2O2 and I– are both involved in the first step in this reaction, and the first step in this reaction is the rate-limiting step, the overall rate of reaction is first-order in both reagents.
Reaction coordinate
Each potential-energy maximum corresponds to the formation of an activated complex. Note that 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).
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 pb 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 = kpBq
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 ofteot activated.
If we know the adsorption isotherm for А, we can express the rate law in terms of its partial pressure pA. For example, if the adsorption of А follows а Langmuir isotherm in the pressure range of interest, then the rate law would be
In А were а diatomic molecule that adsorbed as atoms.
When the partial pressure of А is high (in the sense KpA >> 1) there is almost complete surface coverage, and the rate is equal to kpB. Now the rate-determining step is the collision of В with the adsorbed fragments. When the pressure of А is low (KpA <<1), perhaps because of its reaction, the rate is equal to kKpapb now the extent of surface coverage is important in the determination of the rate.
The Langmuir ‑ Hinshelwood mechanism. In the Langmuir-Hinshelwood mechanism of surface-catalysed reactions, the reaction takes place by encounters between molecular fragments and atoms adsorbed on the surface. We therefore expect the rate
law to be second-order in the extent of surface coverage:
А+ В®P; u = kqAqB
Insertion of the appropriate isotherms for А and В then gives the reaction rate in terms of the partial pressures of the reactants. For example, if A and В follow Langmuir isotherms, and adsorb without dissociation, so that:
The parameters K in the isotherms and the rate constant k are all temperature-dependent, so the overall temperature dependence of the rate may be strongly non-Arrhenius.
Two important types of homogeneous catalysis are acid catalysis and base catalysis, and many organic reactions involve one or the other (and sometimes both). Brensted acid catalysis is the transfer of а proton to the substrate:
X + НА — + HX+ + А–; HX+ then reacts
It is the primary process in the solvolysis of esters, keto-enol tautomerism, and the inversion of sucrose, Brensted base catalysis is the transfer of а hydrogen con from the substrate to а base:
ХН+ В–®Х + ВН+; Х– then reacts
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 ieutral 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.
It is the primary step in the isomerization and halogenation of organic compounds, and of the Claisen and aldol reactions.
Autocatalysis. The phenomenon of autocatalysis is the acceleration of а reaction by the products. For example, in а reaction А® Р it may be found that the rate law is: u = k[A] [P] so the reaction rate increases as products are formed. (The reaction gets started because there are usually other reaction routes for the formation of some Р initially, which then takes part in the autocatalytic reaction proper.) An example of autocatalysis is provided by two steps in the Belousov — Zhabotinskii reaction (BZ reaction) that will figure in discussions later in the section:
BrO3– + НВrО2 + Н3О+ ®2ВrО2 + 2Н2О
2ВrО2+ 2Ce(III) + 2Н3О+ ®2HBrO2 + Ce(IV) + 2Н2О
The product HbrO2 is а reactant in the first step.
The industrial importance of autocatalysis (which occurs in а number of reactions, such as oxidations) is that the rate of the reaction can be maximized by ensuring that the optimum concentrations of reactant and product are always present.
Oscillating reactions. One consequence of autocatalysis is the possibility that the concentrations of reactants, intermediates, and products will vary periodically either in space or in time. Chemical oscillation is the analogue of electrical oscillation, with autocatalysis playing the role of positive feedback. Oscillating reactions are much more than а laboratory curiosity. While they are known to occur in only а few cases in industrial processes, there are many examples in biochemical systems where а cell plays the role of а chemical reactor. Oscillating reactions, for example, maintain the rhythm of the heartbeat. They are also known to occur in the glycolytic cycle, in which one molecule of glucose is used to produce (through enzyme-catalysed reactions involving ATP) two molecules of ATP. All the metabolites in the chain oscillate under some conditions, and do so with the same period but with different phases.
The Lotka — Volterra mechanism. We shall use an autocatalytic reaction of а particularly simple form that illustrates how these oscillations may occur. The actual chemical examples that have been discovered so far have а different mechanism, as we shall see. The Lotka-Volterra mechanism is as follows:
Steps (а) and (b) are autocatalytic. The concentration of А is held constant by supplying it to the reaction vessel as needed. (В plays no part in the reaction once it has been produced, and so it is unnecessary to remove it; in practice, though, it would normally be removed.) These constraints leave [Х] and [Y], the concentrations of the intermediates, as variables. Note that we are considering а steady-state condition, which is maintained by the flow of А into the reactor. This steady-state condition must not be confused with the steady-state approximation made earlier: in the present case we shall solve the rate equations exactly for the variable concentrations of Х and Y, but hold [А] at an arbitrary but constant value.
The Lotka ‑ Volterra equations can be solved numerically, and the results can be depicted in two ways. One way is to plot [Х] and [Y] against time. The same information can be displayed more succinctly by plotting one concentration against the other.
The periodic variation of the concentrations of the intermediates can be explained as follows. At some stage there may be only а little Х present, but reaction (а) provides more, and the production of Х autocatalyses the production of even more Х. There is therefore а surge of Х. However, as Х is formed, reaction (b) can begin. It occurs slowly initially because [Y] is small, but autocatalysis leads to а surge of Y. This surge, though, removes Х, so reaction (а) slows, and less Х is produced. Because less Х is now available, reaction (b) slows. As less Y becomes available to remove Х, Х has а chance to surge forward again, and so on.
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.
Н2 +D2 ®2HD
The reaction 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 ofteecessary. 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 caow 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 О2- 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.
Enzyme inhibitor
An enzyme inhibitor is a molecule, which binds to enzymes and decreases their activity. Since blocking an enzyme’s activity can kill a pathogen or correct a metabolic imbalance, many drugs are enzyme inhibitors. They are also used as herbicides and pesticides. Not all molecules that bind to enzymes are inhibitors; enzyme activators bind to enzymes and increase their enzymatic activity, while enzyme substrates bind and are converted to products in the normal catalytic cycle of the enzyme.
The binding of an inhibitor can stop a substrate from entering the enzyme’s active site and/or hinder the enzyme from catalyzing its reaction. Inhibitor binding is either reversible or irreversible. Irreversible inhibitors usually react with the enzyme and change it chemically (e.g. via covalent bond formation). These inhibitors modify key amino acid residues needed for enzymatic activity. In contrast, reversible inhibitors bind non-covalently and different types of inhibition are produced depending on whether these inhibitors bind to the enzyme, the enzyme-substrate complex, or both.
Many drug molecules are enzyme inhibitors, so their discovery and improvement is an active area of research in biochemistry and pharmacology. A medicinal enzyme inhibitor is often judged by its specificity (its lack of binding to other proteins) and its potency (its dissociation constant, which indicates the concentratioeeded to inhibit the enzyme). A high specificity and potency ensure that a drug will have few side effects and thus low toxicity.
Enzyme inhibitors also occur naturally and are involved in the regulation of metabolism. For example, enzymes in a metabolic pathway can be inhibited by downstream products. This type of negative feedback slows the production line when products begin to build up and is an important way to maintain homeostasis in a cell. Other cellular enzyme inhibitors are proteins that specifically bind to and inhibit an enzyme target. This can help control enzymes that may be damaging to a cell, like proteases or nucleases. A well-characterised example of this is the ribonuclease inhibitor, which binds to ribonucleases in one of the tightest known protein–protein interactions. Natural enzyme inhibitors can also be poisons and are used as defences against predators or as ways of killing prey.
Types of reversible inhibitors
Reversible inhibitors attach to enzymes with non-covalent interactions such as hydrogen bonds, hydrophobic interactions and ionic bonds. Multiple weak bonds between the inhibitor and the active site combine to produce strong and specific binding. In contrast to substrates and irreversible inhibitors, reversible inhibitors generally do not undergo chemical reactions when bound to the enzyme and can be easily removed by dilution or dialysis.
There are four kinds of reversible enzyme inhibitors. They are classified according to the effect of varying the concentration of the enzyme’s substrate on the inhibitor.
In competitive inhibition, the substrate and inhibitor cannot bind to the enzyme at the same time, as shown in the figure on the left. This usually results from the inhibitor having an affinity for the active site of an enzyme where the substrate also binds; the substrate and inhibitor compete for access to the enzyme’s active site. This type of inhibition can be overcome by sufficiently high concentrations of substrate (Vmax remains constant), i.e., by out-competing the inhibitor. However, the apparent Km will increase as it takes a higher concentration of the substrate to reach the Km point, or half the Vmax. Competitive inhibitors are often similar in structure to the real substrate (see examples below).
In uncompetitive inhibition, the inhibitor binds only to the substrate-enzyme complex, it should not be confused with non-competitive inhibitors. This type of inhibition causes Vmax to decrease (maximum velocity decreases as a result of removing activated complex) and Km to decrease (due to better binding efficiency as a result of Le Chatelier’s principle and the effective elimination of the ES complex thus decreasing the Km which indicates a higher binding affinity).
In mixed inhibition, the inhibitor can bind to the enzyme at the same time as the enzyme’s substrate. However, the binding of the inhibitor affects the binding of the substrate, and vice versa. This type of inhibition can be reduced, but not overcome by increasing concentrations of substrate. Although it is possible for mixed-type inhibitors to bind in the active site, this type of inhibition generally results from an allosteric effect where the inhibitor binds to a different site on an enzyme. Inhibitor binding to this allosteric site changes the conformation (i.e., tertiary structure or three-dimensional shape) of the enzyme so that the affinity of the substrate for the active site is reduced.
Non-competitive inhibition is a form of mixed inhibition where the binding of the inhibitor to the enzyme reduces its activity but does not affect the binding of substrate. As a result, the extent of inhibition depends only on the concentration of the inhibitor. Vmax will decrease due to the inability for the reaction to proceed as efficiently, but Km will remain the same as the actual binding of the substrate, by definition, will still function properly.
Quantitative description of reversible inhibition
Reversible inhibition can be described quantitatively in terms of the inhibitor’s binding to the enzyme and to the enzyme-substrate complex, and its effects on the kinetic constants of the enzyme. In the classic Michaelis-Menten scheme below, an enzyme (E) binds to its substrate (S) to form the enzyme–substrate complex ES. Upon catalysis, this complex breaks down to release product P and free enzyme. The inhibitor (I) can bind to either E or ES with the dissociation constants Ki or Ki’, respectively.
Competitive inhibitors can bind to E, but not to ES. Competitive inhibition increases Km (i.e., the inhibitor interferes with substrate binding), but does not affect Vmax (the inhibitor does not hamper catalysis in ES because it cannot bind to ES).
Non-competitive inhibitors have identical affinities for E and ES (Ki = Ki’). Non-competitive inhibition does not change Km (i.e., it does not affect substrate binding) but decreases Vmax (i.e., inhibitor binding hampers catalysis).
Mixed-type inhibitors bind to both E and ES, but their affinities for these two forms of the enzyme are different (Ki ≠ Ki’). Thus, mixed-type inhibitors interfere with substrate binding (increase Km) and hamper catalysis in the ES complex (decrease Vmax).
When an enzyme has multiple substrates, inhibitors can show different types of inhibition depending on which substrate is considered. This results from the active site containing two different binding sites within the active site, one for each substrate. For example, an inhibitor might compete with substrate A for the first binding site, but be a non-competitive inhibitor with respect to substrate B in the second binding site.
Examples of reversible inhibitors
As enzymes have evolved to bind their substrates tightly, and most reversible inhibitors bind in the active site of enzymes, it is unsurprising that some of these inhibitors are strikingly similar in structure to the substrates of their targets. An example of these substrate mimics are the protease inhibitors, a very successful class of antiretroviral drugs used to treat HIV. The structure of ritonavir, a protease inhibitor based on a peptide and containing three peptide bonds, is shown on the right. As this drug resembles the protein that is the substrate of the HIV protease, it competes with this substrate in the enzyme’s active site.
Enzyme inhibitors are often designed to mimic the transition state or intermediate of an enzyme-catalyzed reaction. This ensures that the inhibitor exploits the transition state stabilising effect of the enzyme, resulting in a better binding affinity (lower Ki) than substrate-based designs. An example of such a transition state inhibitor is the antiviral drug oseltamivir; this drug mimics the planar nature of the ring oxonium ion in the reaction of the viral enzyme neuraminidase.
However, not all inhibitors are based on the structures of substrates. For example, the structure of another HIV protease inhibitor tipranavir is shown on the left. This molecule is not based on a peptide and has no obvious structural similarity to a protein substrate. These non-peptide inhibitors can be more stable than inhibitors containing peptide bonds, because they will not be substrates for peptidases and are less likely to be degraded.
In drug design it is important to consider the concentrations of substrates to which the target enzymes are exposed. For example, some protein kinase inhibitors have chemical structures that are similar to adenosine triphosphate, one of the substrates of these enzymes. However, drugs that are simple competitive inhibitors will have to compete with the high concentrations of ATP in the cell. Protein kinases can also be inhibited by competition at the binding sites where the kinases interact with their substrate proteins, and most proteins are present inside cells at concentrations much lower than the concentration of ATP. As a consequence, if two protein kinase inhibitors both bind in the active site with similar affinity, but only one has to compete with ATP, then the competitive inhibitor at the protein-binding site will inhibit the enzyme more effectively.
Irreversible inhibitors
Types of irreversible inhibition
Irreversible inhibitors usually covalently modify an enzyme, and inhibition can therefore not be reversed. Irreversible inhibitors often contain reactive functional groups such as nitrogen mustards, aldehydes, haloalkanes, alkenes, Michael acceptors, phenyl sulfonates, or fluorophosphonates. These electrophilic groups react with amino acid side chains to form covalent adducts. The residues modified are those with side chains containing nucleophiles such as hydroxyl or sulfhydryl groups; these include the amino acids serine (as in DFP, right), cysteine, threonine or tyrosine.
Irreversible inhibition is different from irreversible enzyme inactivation. Irreversible inhibitors are generally specific for one class of enzyme and do not inactivate all proteins; they do not function by destroying protein structure but by specifically altering the active site of their target. For example, extremes of pH or temperature usually cause denaturation of all protein structure, but this is a non-specific effect. Similarly, some non-specific chemical treatments destroy protein structure: for example, heating in concentrated hydrochloric acid will hydrolyse the peptide bonds holding proteins together, releasing free amino acids.
Irreversible inhibitors display time-dependent inhibition and their potency therefore cannot be characterised by an IC50 value. This is because the amount of active enzyme at a given concentration of irreversible inhibitor will be different depending on how long the inhibitor is pre-incubated with the enzyme. Instead, kobs/[I] values are used, wherekobs is the observed pseudo-first order rate of inactivation (obtained by plotting the log of % activity vs. time) and [I] is the concentration of inhibitor. The kobs/[I] parameter is valid as long as the inhibitor does not saturate binding with the enzyme (in which case kobs = kinact).
Analysis of irreversible inhibition
As shown in the figure to the left, irreversible inhibitors form a reversible non-covalent complex with the enzyme (EI or ESI) and this then reacts to produce the covalently modified “dead-end complex” EI*. The rate at which EI* is formed is called the inactivation rate or kinact. Since formation of EI may compete with ES, binding of irreversible inhibitors can be prevented by competition either with substrate or with a second, reversible inhibitor. This protection effect is good evidence of a specific reaction of the irreversible inhibitor with the active site.
The binding and inactivation steps of this reaction are investigated by incubating the enzyme with inhibitor and assaying the amount of activity remaining over time. The activity will be decrease in a time-dependent manner, usually following exponential decay. Fitting these data to a rate equation gives the rate of inactivation at this concentration of inhibitor. This is done at several different concentrations of inhibitor. If a reversible EI complex is involved the inactivation rate will be saturable and fitting this curve will give kinact and Ki.
Another method that is widely used in these analyses is mass spectrometry. Here, accurate measurement of the mass of the unmodified native enzyme and the inactivated enzyme gives the increase in mass caused by reaction with the inhibitor and shows the stoichiometry of the reaction. This is usually done using a MALDI-TOF mass spectrometer. In a complementary technique, peptide mass fingerprinting involves digestion of the native and modified protein with a protease such as trypsin. This will produce a set of peptides that can be analysed using a mass spectrometer. The peptide that changes in mass after reaction with the inhibitor will be the one that contains the site of modification.
USES OF INHIBITORS
Enzyme inhibitors are found iature and are also designed and produced as part of pharmacology and biochemistry. Natural poisons are often enzyme inhibitors that have evolved to defend a plant or animal against predators. These natural toxins include some of the most poisonous compounds known. Artificial inhibitors are often used as drugs, but can also be insecticides such as malathion, herbicides such as glyphosate, or disinfectants such as triclosan.
Chemotherapy
The most common uses for enzyme inhibitors are as drugs to treat disease. Many of these inhibitors target a human enzyme and aim to correct a pathological condition. However, not all drugs are enzyme inhibitors. Some, such as anti-epileptic drugs, alter enzyme activity by causing more or less of the enzyme to be produced. These effects are called enzyme induction and inhibition and are alterations in gene expression, which is unrelated to the type of enzyme inhibition discussed here. Other drugs interact with cellular targets that are not enzymes, such as ion channels or membrane receptors.
An example of a medicinal enzyme inhibitor is sildenafil (Viagra), a common treatment for male erectile dysfunction. This compound is a potent inhibitor of cGMP specific phosphodiesterase type 5, the enzyme that degrades the signalling molecule cyclic guanosine monophosphate. This signalling molecule triggers smooth muscle relaxation and allows blood flow into the corpus cavernosum, which causes an erection. Since the drug decreases the activity of the enzyme that halts the signal, it makes this signal last for a longer period of time.
Another example of the structural similarity of some inhibitors to the substrates of the enzymes they target is seen in the figure comparing the drug methotrexate to folic acid. Folic acid is a substrate of dihydrofolate reductase, an enzyme involved in making nucleotides that is potently inhibited by methotrexate. Methotrexate blocks the action of dihydrofolate reductase and thereby halts the production of nucleotides. This block of nucleotide biosynthesis is more toxic to rapidly growing cells thaon-dividing cells, since a rapidly growing cell has to carry out DNA replication, therefore methotrexate is often used in cancer chemotherapy.
Drugs also are used to inhibit enzymes needed for the survival of pathogens. For example, bacteria are surrounded by a thick cell wall made of a net-like polymer called peptidoglycan. Many antibiotics such as penicillin and vancomycin inhibit the enzymes that produce and then cross-link the strands of this polymer together. This causes the cell wall to lose strength and the bacteria to burst. In the figure, a molecule of penicillin (shown in a ball-and-stick form) is shown bound to its target, the transpeptidase from the bacteria Streptomyces R61 (the protein is shown as a ribbon-diagram).
Drug design is facilitated when an enzyme that is essential to the pathogen’s survival is absent or very different in humans. In the example above, humans do not make peptidoglycan, therefore inhibitors of this process are selectively toxic to bacteria. Selective toxicity is also produced in antibiotics by exploiting differences in the structure of the ribosomes in bacteria, or how they make fatty acids.
Metabolic control
Enzyme inhibitors are also important in metabolic control. Many metabolic pathways in the cell are inhibited by metabolites that control enzyme activity through allosteric regulation or substrate inhibition. A good example is the allosteric regulation of the glycolytic pathway. This catabolic pathway consumes glucose and produces ATP, NADH and pyruvate. A key step for the regulation of glycolysis is an early reaction in the pathway catalysed by phosphofructokinase-1 (PFK1). When ATP levels rise, ATP binds an allosteric site in PFK1 to decrease the rate of the enzyme reaction; glycolysis is inhibited and ATP production falls. This negative feedback control helps maintain a steady concentration of ATP in the cell. However, metabolic pathways are not just regulated through inhibition since enzyme activation is equally important. With respect to PFK1, fructose 2,6-bisphosphate and ADP are examples of metabolites that are allosteric activators.
Physiological enzyme inhibition can also be produced by specific protein inhibitors. This mechanism occurs in the pancreas, which synthesises many digestive precursor enzymes known as zymogens. Many of these are activated by the trypsin protease, so it is important to inhibit the activity of trypsin in the pancreas to prevent the organ from digesting itself. One way in which the activity of trypsin is controlled is the production of a specific and potent trypsin inhibitor protein in the pancreas. This inhibitor binds tightly to trypsin, preventing the trypsin activity that would otherwise be detrimental to the organ. Although the trypsin inhibitor is a protein, it avoids being hydrolysed as a substrate by the protease by excluding water from trypsin’s active site and destabilising the transition state. Other examples of physiological enzyme inhibitor proteins include the barstar inhibitor of the bacterial ribonuclease barnase and the inhibitors of protein phosphatases.
Pesticides and herbicides
Many herbicides and pesticides are enzyme inhibitors. Acetylcholinesterase (AChE) is an enzyme found in animals from insects to humans. It is essential to nerve cell function through its mechanism of breaking down the neurotransmitter acetylcholine into its constituents, acetate and choline. This is somewhat unique among neurotransmitters as most, including serotonin, dopamine, and norepinephrine, are absorbed from the synaptic cleft rather than cleaved. A large number of AChE inhibitors are used in both medicine and agriculture. Reversible competitive inhibitors, such as edrophonium, physostigmine, and neostigmine, are used in the treatment of myasthenia gravis and in anaesthesia. The carbamate pesticides are also examples of reversible AChE inhibitors. The organophosphate insecticides such as malathion, parathion, and chlorpyrifos irreversibly inhibit acetylcholinesterase.
The herbicide glyphosate is an inhibitor of 3-phosphoshikimate 1-carboxyvinyltransferase, other herbicides, such as the sulfonylureas inhibit the enzyme acetolactate synthase. Both these enzymes are needed for plants to make branched-chain amino acids. Many other enzymess are inhibited by herbicides, including enzymes needed for the biosynthesis of lipids and carotenoids and the processes of photosynthesis and oxidative phosphorylation.
To discourage seed predators, pulses contain trypsin inhibitors that interfere with digestion.
Natural poisons
Animals and plants have evolved to synthesise a vast array of poisonous products including secondary metabolites, peptides and proteins that can act as inhibitors. Natural toxins are usually small organic molecules and are so diverse that there are probably natural inhibitors for most metabolic processes. The metabolic processes targeted by natural poisons encompass more than enzymes in metabolic pathways and can also include the inhibition of receptor, channel and structural protein functions in a cell. For example, paclitaxel (taxol), an organic molecule found in the Pacific yew tree, binds tightly to tubulin dimers and inhibits their assembly into microtubules in the cytoskeleton.
Many natural poisons act as neurotoxins that can cause paralysis leading to death and have functions for defence against predators or in hunting and capturing prey. Some of these natural inhibitors, despite their toxic attributes, are valuable for therapeutic uses at lower doses. An example of a neurotoxin are the glycoalkaloids, from the plant species in the Solanaceae family (includes potato, tomato and eggplant), that are acetylcholinesterase inhibitors. Inhibition of this enzyme causes an uncontrolled increase in the acetylcholine neurotransmitter, muscular paralysis and then death. Neurotoxicity can also result from the inhibition of receptors; for example, atropine from deadly nightshade (Atropa belladonna) that functions as a competitive antagonist of the muscarinic acetylcholine receptors.
Although many natural toxins are secondary metabolites, these poisons also include peptides and proteins. An example of a toxic peptide is alpha-amanitin, which is found in relatives of the death cap mushroom. This is a potent enzyme inhibitor, in this case preventing the RNA polymerase II enzyme from transcribing DNA. The algal toxin microcystin is also a peptide and is an inhibitor of protein phosphatases. This toxin can contaminate water supplies after algal blooms and is a known carcinogen that can also cause acute liver hemorrhage and death at higher doses.
Proteins can also be natural poisons or antinutrients, such as the trypsin inhibitors (discussed above) that are found in some legumes, as shown in the figure above. A less common class of toxins are toxic enzymes: these act as irreversible inhibitors of their target enzymes and work by chemically modifying their substrate enzymes. An example is ricin, an extremely potent protein toxin found in castor oil beans. This enzyme is a glycosidase that inactivates ribosomes. Since ricin is a catalytic irreversible inhibitor, this allows just a single molecule of ricin to kill a cell.
6. ENZYMES AS BIOLOGICAL CATALYSTS.
Enzymes are macromolecules that help accelerate (catalyze) chemical reactions in biological systems. This is usually done by accelerating reactions by lowering the transition state or decreasing the activation energy.
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.
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.
Figure. Enzyme decreases activation energy
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.
Properties of enzymes. By definition а catalyst is а substance that enhances the rate of а chemical reaction but is not permanently altered by the reaction. Catalysts perform this feat because they provide an alternative reaction path-way that requires less energy than the uncatalyzed reaction. During any chemical reaction the energy of the system increases until the transition state is reached. At this point, а high proportion of substrate molecules have become sufficiently energized to enter into an intermediate state that has а high probability of being converted into product. For example, the transition state of the reaction in which ethanol is oxidized to form acetaldehyde, might look like
Several mechanisms have been proposed that attempt to explain how catalysts work. According to one hypothesis, а catalyst-reactant complex is formed. As the reaction products are formed, they are released from the complex, and the catalyst is then available to bind new reactant molecules. А well-regarded variation of this proposal suggests that catalysts provide а surface on which reactants are adsorbed. As reactant molecules bind to the catalyst’s surface, they are oriented in such а manner that increases the likelihood of product formation. As products are formed, they leave their adsorption sites, which are then available for other reactant molecules. For example, in the hydrogenation of an alkene, а platinum catalyst serves as а surface upon which the reactants are adsorbed. It is believed that the catalyst facilitates the reaction because its binding of H2 results in the weakening and subsequent breakage of the Н — Н bond.
Catalysts also work because they decrease the activation energy required for а chemical reaction to proceed. It should be emphasized that these reactions occur spontaneously in the absence of the catalyst. No catalyst can affect reactions that are not thermodynamically favorable. Catalysts cannot alter the equilibrium point of а reaction, but they can accelerate the rate at which equilibrium is attained. (Note that а catalyst speeds up а reaction in both the forward and reverse directions.) Consider the following reversible reaction: А = В
In the absence of а catalyst the conversion of the reactant (А) into product (В) occurs at а certain rate. Because this is а reversible reaction, there is also а rate of conversion of В into А. The rate expression for the forward reaction is RF[A]n, and that for the reverse reaction is to RR [В]m. (The superscripts и and m represent the order of а reaction.) At equilibrium the rates for the forward and reverse reactions must be equal:
RF[A]n = RR [В]m
Which rearranges to:
The ratio of the forward and reverse rates is the equilibrium constant:
In equation (3), if m == 1 and RF = 1 . 10-3 and RR = 1 .10-6 s-1 then
At equilibrium, therefore, the ratio of products to reactants would be 1000 to 1. The difference between the catalyzed and uncatalyzed reactions is one of time. In the presence of the catalyst the equilibrium is attained in seconds or minutes instead of hours or days.
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.
Figure. Key Model Of Enzyme Specificity
Catalysis happens at the active siteof the enzyme. It contains the residues that directly participate in the making and breaking of bonds. These residues are called the catalytic groups. Although enzymes differ widely in structure, specificity, and mode of catalysis a number of generalizations concerning their active sites can be made:
1. The active site is a three dimensional cleft or crevice formed by groups that come from different parts of the amino acid sequence – residues far apart in the amino acid sequence may interact more strongly than adjacent residues in the sequence.
2. The active site takes up a relatively small part of the total volume of an enzyme. Most of the amino acid residues in an enzyme are not in contact with the substrate, which raises the question of why enzymes are so big. Nearly all enzymes are made up of more than 100 amino acid residues. The “extra” amino acids serve as a scaffold to create the three dimensional active site from theamino acids that are far apart in the primary structure. In many proteins the remaining amino acids also constitute regulatory sites, sites of interaction with other proteins, or channels to bring the substrate to the active sites.
3. Active sites are unique microenvironments. In all enzymes of known structure, substrate molecules are bound to a cleft or crevice. Water is usually excluded unless it is a reactant. The nonpolar microenvironment of the cleft enhances the binding of substrates as well as catalysis. Nevertheless, the cleft may also contain polar residues. Certain of these polar residues acquire special properties essential for substrate binding or catalyis.
4. Substrates are bound to enzymes by multiple weak interactions. Stated above
5. The specificity of binding depends on the precise defined arrangement of atoms in the active site. Because the enzyme and the substrate interact by means of short-range forces that require close contact, a substrate must have a matching shape to fit into the site. However, the active site of some enzymes assume a shape that is complementary to that of the substrate only after the substrate is bound. This process of dynamic recognition is called induced fit.
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.
Figure. Strycture of enzymes
Classification of enzymes. 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 ofteamed 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-knowames 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.
Figure. Diagrams to show the induced fit hypothesis of enzyme action
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. For example, enzymes containing free sulfhydryl groups can react with alkylating agents such as iodoacetate:
Glyceraldehyde-3-phosphate dehydrogenase, an enzyme in the glycolytic pathway, is inactivated by alkylation with iodoacetate. Enzymes that use sulfhydryl groups to form covalent bonds with metal cofactors are often irreversibly inhibited by heavy metals (е.g., mercury and lead). The anemia that occurs in lead poisoning is caused in part because of lead binding to а sulfhydryl group of ferrochelatase. Ferrochelatase catalyzes the insertion of Fe3+ into heme.
In reversible inhibition the inhibitor can dissociate from the enzyme because it binds through noncovalent bonds. The most common forms of reversible inhibition are competitive and noncompetitive .
Competitive inhibition. The structure of а competitive inhibitor closely resembles that of the enzyme’s normal substrate. Because of its structure, а competitive inhibitor binds reversibly to the enzyme’s active site. In so doing, the inhibitor forms an enzyme-inhibitor complex (El) that is equivalent to the ES complex:
The concentration of El complex depends on the concentration of free inhibitor and on the dissociation constant KI.
The inhibition is said to be competitive because the El complex readily dissociates. The empty active site is then available for substrate binding. Because по productive reaction occurs during the finite amount of time that the ЕI complex exists, the enzyme’s activity is observed to decline . The effect of а competitive inhibitor on activity is reversed by increasing the concentration of substrate. At high [S], all the active sites are filled with substrate, and reaction velocity reaches the value observed in the absence of inhibitor. Succinate dehydrogenase, an enzyme in the Krebs citric acid cycle, atalyzes the following redox reaction:
This reaction is inhibited by malonate. Malonate binds to the enzyme’s active site but cannot be converted to product. Succinate = Fumarate
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.
Most proteins, and therefore enzymes, are active only within a narrow pH range usually between 5 and 9. Several factors are influenced directly by the pH in which the reaction takes place.
· the binding of substrate to the enzyme
· the ionization states of the amino acid residues involved in the catalytic activity of the enzyme.
· the ionization of the substrate
· variation in the protein structure at extreme pH.
The graph of pH against the reaction rate is a bell shaped curve. The curve reprecents the ionization of certain amino acid residue that must be in a specific ionization state for enzyme activity. The inflection point of the curve is called the pK of the reaction and can identify amino acid residues essential to enzymatic activity.
Figure. Enzyme disturtions due to extreme pH
Effects of Temperature on ENZYMES.
Temperature can effect an enzyme in two ways. One is a direct influence on the reaction rate constant, and the other is in thermal denturization of the enzyme at elevated temperatures.
To relate the effect of temperature to the reaction rate constant, the Arrhenius equation is used:
where k is the rate constant, R is the gas law constant, A is the frequency factor and Ea is the activation energy of the reaction.
The temperature range over which enzymes show activity is limited between the melting point (0oC) and bioling point (100oC) of water. If a temperature is too low, there can be no noticable reaction rate since the enzyme is operating at a temperature far below its optimum. If the temperature at which the enzyme is operating at is well above 100oC, then thermal deactivization can occur.
Thermal deactivization of enzymes limits their useful lifetime in processing environments. Therefore, it is important in many preocess design and manufacturing levels to have the correct temperature of reaction. If the reaction temperature is too high, the enzymes will eventually deactivate in an irreversible way and thus halting the reaction from taking place. For many enzymes found within mammals, the optimum temperature is 37oC, but deactivization can occur as low as 45 to 55oC. Deactivization of enzymes may be irreversible or reversible.
References:
1. The abstract of the lecture.
2. intranet.tdmu.edu.ua/auth.php
3. Atkins P. W. Physical chemistry / P.W. Atkins. – New York, 1994. – P.299‑307.
4. Cotton F. A. Chemical Applications of Group Theory / F. A. Cotton. ‑ John Wiley & Sons : New York, 1990.
5. Girolami G. S. Synthesis and Technique in Inorganic Chemistry / G. S. Girolami, T. B. Rauchfuss, R. J. Angelici. ‑ University Science Books : Mill Valley, CA, 1999.
6. Russell J. B. General chemistry / J B. Russell. New York.1992. – P. 550‑599.
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Prepared by PhD Falfushynska H.
