Simply put, Chemical kinetics addresses the question: “how fast do reactions occur?” Chemistry can be described as the science that deals with making new substances from other substances. Or, we could say, it is taking molecules apart and putting together the atoms and fragments to form new molecules. So, if chemistry is about creating new substances out of old substances (i.e., chemical reactions), then there are two basic things that must be answered.
- Does the reaction wants to happen? This is about chemical thermodynamics (deals with the direction in which a process occurs).
- If the reaction wants to go, how fast will it go? This is the topic of chemical kinetics.
Considered to be one of the easiest of the science subjects, Chemical Kinetics is almost like an acid test for JEE aspirants. It’s crucial to have a sound knowledge of the subject and its formulae as it can give candidates the golden opportunity to crack the JEE and get into prestigious IITs. So, if you think it’s tricky to make it through this section in JEE, put all your worries away as Toppr.com brings you a quick revision of the related concepts.
Today, in this article, we’ll take a closer look at the term ‘Chemical Kinetics’ and also understand its relevance for JEE.
What is Chemical Kinetics?
The field of chemical kinetics was born from the law of mass action, which was formulated in 1864 by Peter Waage and Cato Guldberg. Chemical Kinetics is also reaction kinetics or simply “kinetics”. The rate of a chemical reaction usually consists units of.
The process of Chemical Kinetics also includes the analysis of conditions that directly affect the speed of a chemical reaction, reaction mechanisms, and transition states. It is used to form mathematical models to predict and describe a chemical reaction as well.
Understanding Rate Laws Of Chemical Kinetics
A rate law is an equation that determines how fast a reaction proceeds and how the reaction rate depends on the concentrations of the chemical species involved. A rate law is an equation of the form:
V = f ([A], [B], … [E])
Experimental data is often used for finding reaction rates, from which rate laws and chemical kinetics rate constants are derived by applying the law of mass action. Rate laws help simple calculations for zero-order reactions, first-order reactions, and second order reactions.
Ideally, procedures for reactions that take place within hours or minutes use of various techniques for determining concentration, such as spectroscopy and electrochemistry. When the reactions are very fast, they are studied spectroscopically. Spectroscopic procedures are available for monitoring reactions that are initiated by a rapid pulse of electromagnetic radiation and are over in a few femtoseconds (1 fs = 10−15 s).
Common Types of Rate Laws
- First Order Reactions
In a first-order reaction, the rate is directly proportional to the concentration of one of the reactants. That is,
v = rate = k[B],
where B is a reactant. If there’s a reaction that is known to be first order in B, such as
B + other reactants → products,
we can write the rate law as,
–d [B]=k [B]
The constant, k, in this rate equation, is the first order rate constant.
- Second Order Reactions
In a second-order reaction, the rate is proportional to the concentration squared. For example, possible second order rate laws could be written as:
Rate = k[B]2
Rate = k[A][B].
That is, the rate might be proportional to the square of the concentration of one of the reactants, or it might be proportional to the product of two different concentrations.
- Third Order Reactions
There are many unique ways to write a rate law for a third order reaction. One might have cases where
Rate = k[A]3,
Rate = k[A]2[B],
Rate = k[A][B][C],
and so on.
Rate laws for individual steps must be combined to derive laws for more complex chemical reactions. For these reactions:
- There is a rate-determining step that limits the kinetics.
- One can also use the Arrhenius equation and Eyring equations to experimentally determine activation energy.
- Steady-state approximations may be applied to simplify the rate law.
Example of Rate Law
The rate law shows the relationship between the reaction rate to the concentration of each reactant. Therefore, the rate of reaction completely depends on the concentration of a product; if a product has a higher concentration, the faster the rate law will be and vice-versa.
A real life example will be a scenario where you are deciding where to build infrastructure such as a school or hospital. So, you would need to consider how fast the population is changing in that particular area. This question will easily determine the amount (concentration of the product), which in this case, is the capacity of schools that need to be built.
Factors That Affect the Chemical Reaction Rate
The main factors that affect reaction rate are:
- the concentration of reactants (increasing concentration increases reaction rate)
- temperature (increasing temperature increases reaction rate, up to a point)
- presence of catalysts (catalysts offer a reaction a mechanism that requires a lower activation energy, so the presence of a catalyst increases the rate of a reaction)
- the physical state of reactants (reactants in the same phase may come into contact via thermal action, but surface area and agitation affect reactions between reactants in different phases)
- pressure (for reactions involving gases, raising pressure increases the collisions between reactants, increasing reaction rate)
Theories Of Reaction Rates:
1) Collision Theory
If two molecules need to collide in order for a reaction to take place, the factors that influence the ease of collisions will be important. The more energy there is available to the molecules, the faster they will move around, and the more likely they are to bump into each other. Higher temperatures ought to lead to more collisions and a greater frequency of reactions between molecules.
2) Transition-State Theory
Transition state theory (TST) provides a more accurate alternative to the previously used Arrhenius equation and the collision theory. The transition state theory attempts to provide a greater understanding of activation energy, Ea, and the thermodynamic properties involving the transition state. Collision theory of reaction rate, although intuitive, lacks an accurate method to predict the probability factor for the reaction. The theory assumes that reactants are hard spheres rather than molecules with specific structures. In 1935, Henry Eyring helped develop a new theory called the transition state theory to provide a more accurate alternative to the previously used Arrhenius equation and the collision theory. The Eyring equation involves the statistical frequency factory, v, which is fundamental to the theory.
According to TST, between the state where molecules are reactants and the state where molecules are products, there is a state known as the transition state. In the transition state, the reactants are combined in a species called the activated complex. The theory suggests that there are three major factors that determine whether a reaction will occur:
- The concentration of the activated complex
- The rate at which the activated complex breaks apart
- The way in which the activated complex breaks apart: whether it breaks apart to reform the reactants or whether it breaks apart to form a new complex, the products.
Chemical Reactions On The Basis Of Rate Of Reaction
The rate of reaction measures how much product is formed in a certain time. The mass of a solid product is often measured in grams while the volume of a gaseous product is often measured in cm3. The time period chosen may depend upon the rate of the reaction. For example, it may be a few seconds for a fast reaction or a few minutes for a slow reaction.
Chemical reactions can be categorized as slow or fast based on the rate of a chemical reaction.
Slow Chemical Reactions: There are chemical reactions that are fast, but there are few that are slow and take place naturally and in the course of time the difference is notable. These are those chemical reactions which take place at very slow rate. These reactions can take days, months or even years to complete. In general the reactions between covalent compounds are slow. For example rusting of iron or fermentation of sugar into ethyl alcohol and carbon dioxide can take place in several hours or even in several days. In the same way weathering of rocks takes place in millions of years.
Fast Chemical Reactions: These are those chemical reactions which take place at a very fast rate. These reactions can take place in seconds or in minutes. In general the reactions between ionic compounds are fast. For example, combustion of LPG gas in kitchen takes place in a few seconds so it is a fast reaction. In the same way reaction of an acid like HCl with a base such as NaOH takes place in seconds to produce a salt NaCl and water is very fast reaction. This reaction is also called neutralization reaction.
Measuring Slow Reactions:
The best way to study exceedingly slow reactions is to change the conditions so that the reactions occur in a reasonable time. Increasing the temperature, which can have a strong effect on the reaction rate, is one possibility. If the temperature of a hydrogen-oxygen mixture is raised to about 500 °C (900 °F), reaction then occurs rapidly, and its kinetics has been studied under those conditions. When a reaction occurs to a measurable extent over a period of minutes, hours, or days, rate measurements are straightforward. Amounts of reactants or products are measured at various times, and the rates are readily calculated from the results. Many automated systems have now been devised for measuring rates in this way.
Measuring Fast Reactions:
Some processes are so fast that special techniques have to be used to study them. There are two difficulties with fast reactions. One is that the time that it takes to mix reactants or to change the temperature of the system may be significant in comparison with the half-life, so that the initial time cannot be measured accurately. The other difficulty is that the time it takes to measure the amounts of substances may be comparable with the half-life of the reaction. The methods used to overcome these difficulties fall into two classes: flow methods and pulse and probe methods.
In flow methods, two gases or solutions are introduced rapidly into a mixing vessel, and the resulting mixture then flows rapidly along a tube. Concentrations of reactants or products may then be measured—for example, by spectroscopic methods—at various positions along the tube, which correspond to various reaction times. A modification of this method is the stopped-flow technique, in which the reactants are forced rapidly into a reaction chamber; the flow is then suddenly stopped, and the amounts are measured by physical methods after various short times. These flow methods are limited by the time it takes to mix gases or solutions and are not suitable if the half-life is less than about a hundredth of a second.
These mixing difficulties were overcome by pulse and probe methods. The principle of these is that a short pulse, usually of radiation, is given to a chemical system and is then followed by a probe, usually involving radiation that provides spectroscopic evidence of what occurred after the initial pulse. The first of these methods, developed in 1949 by British chemists R.G.W. Norrish and George Porter, was the flash-photolysis method, for which Norrish and Porter won the Nobel Prize for Chemistry in 1967. In this technique a flash of light of high intensity but short duration brings about the formation of atomic and molecular species, the reactions of which can be studied kinetically by spectroscopy. In the earliest experiments the duration of the flash was about a millisecond (ms; 1 ms = 10–3 second), but in the next four decades the duration was reduced by more than 11 powers of 10, to just a few femtoseconds. A nanosecond (ns; 1 ns = 10–9 second) flash is adequate for studying almost any purely chemical reaction where there is a change in chemical identity. Any chemical reaction, however, involves processes of a purely physical nature, such as energy redistribution and the breakdown of transient species, which occur in the femtosecond range.
Composite Reaction Mechanism:
Various lines of evidence are used to determine if a reaction occurs in more than one step. Suppose that the kinetic equation for the reaction does not correspond to the balanced equation for the reaction.
A simple example is the reaction between hydrogen and iodine chloride, with the formation of iodine and hydrogen chloride: H2 + 2ICl → I2 + 2HCl.
To make the equation balance, the reaction must be written as shown, with two iodine chloride molecules reacting with a single hydrogen molecule. If this reaction occurred in a single elementary step, the rate would be proportional to the first power of the hydrogen concentration and the square of the iodine chloride concentration.
Instead, however, the rate is found to be proportional to both concentrations to the first power, so that it is a second-order reaction:v = k[H2][ICl].
This can be explained if there is initially a slow reaction between one hydrogen molecule and one of iodine chloride: H2 + ICl → HI + HCl (slow) followed by a rapid reaction between the hydrogen iodide formed and an additional molecule of iodine chloride: HI + ICl → HCl + I2 (fast).
If the second reaction is fast, the hydrogen iodide is removed as fast as it is formed. The rate of the second reaction therefore has no effect on the overall rate, which is the rate of the first step. This mechanism therefore explains the kinetic behaviour but does not prove it; other, more complicated schemes could be devised, but, until there is further evidence, it is expedient to accept the simple mechanism. This is an example of a consecutive reaction, which occurs in two steps, with the intermediate playing a role.
Another piece of evidence for a composite mechanism is the detection of reaction intermediates. In such a case, a reaction scheme must be devised that will account for these intermediates. Sometimes an intermediate can be a fairly stable substance. In other cases the intermediates are unstable species such as atoms and free radicals (fragments of molecules) that subsequently undergo rapid reactions. Free radicals can be detected by spectroscopy and other means. When organic molecules are raised to high temperatures, they decompose into smaller molecules, and organic free radicals have often been detected as intermediates. In an explosion, such as that between hydrogen and oxygen, free radicals such as hydroxyl can be detected.
Composite reaction mechanisms are of various kinds. Aside from the simple consecutive schemes, there are some special mechanisms that give rise to oscillatory behaviour: the amount of a product continuously rises and falls over a period of time. The conditions for this behaviour are that there must be at least two species involved in the reaction and there must be feedback, which means that products of the reaction affect the rate. There are also reaction mechanisms that give rise to what is technically known as chaos, or catastrophe. With such reactions it is impossible to predict the outcome. Chaotic conditions also require that there be feedback and that at least three species be involved.
Sometimes a complex reaction mechanism involves a cycle of reactions such that certain intermediates consumed in one step are regenerated in another.
For example, the accepted mechanism of the reaction between hydrogen and bromine, which can be written as
H2 + Br2 → 2HB
includes the steps
1. Br + H 2 → HBr + H
2. H + Br 2 → HBr + Br
In the first of these steps a bromine atom is consumed, but in the second a bromine atom is regenerated. This pair of reactions can thus occur with the production of two molecules of hydrogen bromide, the product of the reaction, without loss of bromine atoms. This pair of reactions is called a cycle of reactions, and it can occur a number of times, in which case the reaction is referred to as a chain reaction. The two reactions in which bromine is regenerated are known as the chain-propagating steps. The average number of times the pair of steps is repeated is known as the chain length.
One necessary condition for a proposed reaction mechanism to be correct is that it must account for the overall kinetic behaviour of the reaction—in particular, for the dependence of the reaction rate on the reactant concentrations. For any proposed reaction mechanism, it is possible to write down equations for the rate of each step in terms of the reactant concentration and then to solve the equations for the overall rate. A practical difficulty arises, since no exact mathematical solution is possible for all except the simplest of mechanisms. If one has values for the rate constants, solutions can be obtained with a computer, but explicit rate equations provide more insight into the reactions. One therefore looks for approximate solutions of the equations. One of these is provided by the steady-state treatment, which is applicable if (and only if) the intermediates are species that can be present only at low concentrations. If this condition is satisfied by an intermediate, the rate of change of its concentration during the course of reaction is always small and, as a good approximation, can be assumed to be zero, which means that the intermediate exists in a steady state. This approximation may safely be applied to atoms and free radicals present as reaction intermediates. With this approximation it is usually possible to obtain a reliable approximate equation for the overall reaction rate in terms of reactant concentrations. If this agrees with the experimental behaviour, the mechanism is accepted.
One situation to which the steady-state treatment does not apply is when a reaction is an explosion. Explosions occur because the concentration of intermediates does not remain steady during the course of reaction but rises to a high value so that the reaction goes out of control. This occurs if the reaction mechanism involves a special kind of chain called a branching chain.
In the hydrogen-oxygen explosion, for example, the following reaction is known to occur:
H + O2 → OH + O
In this step a single chain carrier hydrogen atom has produced two chain carriers: a hydroxyl group and an oxygen atom. The number of chain carriers increases rapidly and leads to an explosion.
Applications of Chemical Kinetics in Real Life
- The study of chemical kinetics unravels the deep secret of many things in our day-to-day lives. As per the mathematical models, chemical reaction kinetics provide chemists and chemical engineers with tools to understand and describe chemical processes such as food decomposition, microorganism growth, stratospheric ozone decomposition, and the complex chemistry of biological systems in a better way.
- Chemical kinetics can be used in designing or modifying chemical reactors and thus optimize product yield, separate products efficiently, and eliminate environmentally-dangerous by-products.
- Even while performing catalytic cracking of heavy hydrocarbons into gasoline and light gas, Chemical Kinetics models can be used to find out the temperature and pressure at which the highest yield of heavy hydrocarbons into gasoline will occur.
- As we know, popcorn is different than the regular corn. For making delicious popcorn, once the kernels are heated, the water begins to boil inside and expands from a liquid into a gas. The gas pushes against the hard shell and the kernel explodes or “pops” when the pressure increases. The kernels that do not “pop” are generally ones that have little water inside them to create enough pressure to explode. So, the underlying concept is:
- as the temperature increases, liquids tend to become gases.
- Liquids tend to be denser than gases.
- As the temperature increases so do the pressure in a fixed volume.
Important Topics in Chemical Kinetics for JEE
Chemical Kinetics is one of the most important topics in chemistry, and can really help increase your score for JEE drastically. Questions from the concept of order of reaction, how to determine order of reaction along and integrated rate laws have been repeatedly asked in IIT-JEE in the past. Here are some more topics of Chemical Kinetics that should be studied without fail:
- Rate of Reaction
- Molecularity & Order of Reaction
- Methods for Determination of Order of a Reaction
- Factors Affecting Rate of Reaction
- Zero Order Reaction
- First Order Reactions
- Parallel and Sequential Reactions
- Collision Theory of Reaction Rate
The Right Way to Approach a Question on Chemical Kinetics
Here’s an example:
Question: Calculate the half-life for a first-order reaction if 68% of the reactant is reacted within 66 s.
Answer: 68% of the reactant has been consumed implies that 32% remains, hence using the integrated equation for 1st order reactions: ln A = -kt + ln Ao ln 0.32Ao = – k (66) + ln Ao k = 0.0172642 s-1. This calculation is done so that you can know how much substance remains, not how much is used up.
We don’t really think that students should have problems with this topic as it is well-covered in the JEE syllabus [Read also: Difficult Topics From JEE Syllabus]. Reading the books prescribed by CBSE might suffice for most students. So, try to make this one of your strongest topics as it will help you in understanding many other concepts in chemistry as well.
You may also like: