Reaction Kinetics, Chemistry tutorial


Define: It is a stream of chemistry which mainly deals by the rate of chemical reactions, with factors affecting such rates, and by applications of rate studies to explain the method of reactions.

Chemical kinetics is mainly concerned by the rates of chemical reactions, that is, by the quantitative explanation of how fast chemical reactions take place and the factors influencing these rates. The chemist employs kinetics as a tool to comprehend the basic features of reaction pathways, a subject which continues to evolve by ongoing research. The applied chemist employs this understanding to work out new and/or better ways of accomplishing desired chemical reactions. This might comprise enhancing the outcome of desired products or making a better catalyst. The chemical engineer employs kinetics for reactor design in chemical reaction or process engineering.

Chemical kinetics is basically the study and conversation of chemical reactions with respect to reaction rates, effect of different variables, re-arrangement of atoms and preparation of intermediates and so on. There are numerous topics to be discussed and each of such topics as a tool for the study of chemical reactions. By the way, the study of motion is termed as kinetics, from the Greek kinesis, meaning movement.

At the macroscopic level, we are concerned in amounts reacted, formed, and the rates of their formation. At molecular or microscopic level, the given considerations should as well be made in the conversation of chemical reaction methods.

  • Molecules or atoms of reactants should collide by one other in chemical reactions.
  • The molecules should encompass adequate energy to initiate the reaction.
  • In certain cases, the orientation of the molecules throughout the collision should as well be considered.

Reaction Rates:

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

With respect to the reaction rates, we might deal by average rates, instantaneous rates and initial rates based on the experimental conditions.

Thermodynamics and kinetics are the two factors which influence reaction rates. The study of energy gained or liberated in chemical reactions is known as thermodynamics, and such energy data are known as thermodynamic data. Though, thermodynamic data encompass no direct correlation with the reaction rates, for which the kinetic factor is possibly more significant. For illustration, at room temperature (that is, a broad range of temperatures), thermodynamic data points out that diamond shall transform to graphite, however in reality, the conversion rate is extremely slow that most of the people assume that diamond is forever.

Rate Laws:

The rate of a chemical reaction is, possibly, it's most significant property as it dictates whether a reaction can take place throughout a lifetime. Recognizing the rate law, an expression associating to the rate to the concentrations of reactants, can assists a chemist in adjusting the reaction conditions to obtain a more appropriate rate. Whenever there are two competing reactions for a single reagent, one can, recognizing the rate law, favor the exclusive preparation of a single product.

To get this type of knowledge regarding reactions, we will initially define what rate signifies. We will then derive the rate law expression. By employing the process of initial rates, we will discuss how to find out the form and order of the rate law. Subsequently, we will probe rate laws in depth and introduce the integrated rate law as the alternative form of the simple rate law which allows us the other, simpler, experimental process to find out the order of the rate law. The integrated rate law will as well let us to find out the half-lives of chemical reactions.

Terms related to Reaction Rates:

Absorbance Spectroscopy: The analytical method which shines light of one specific wavelength via a sample cell to test what percentage of the incoming light is absorbed. Via comparing that absorbance datum to absorbance for cells of known concentration, we can compute the unknown concentration.

Half-Life: The time essential for the utilization (consumption) or decay of half of the limiting reagent.

Integrated Rate Law: The integral (that is, a calculus operation) of the rate law, this form of the rate law exhibits the dependence of the concentration of reactants on time of reaction.

Kinetics: The study of the rate and procedure of chemical reactions.

Method of Initial Rates: A sequence of experiments in which the concentration of one reagent at a time is varied and the initial rate (that is, rate at time zero) of the reaction is evaluated. By evaluating the change in concentration to the change in rate, it is likely to find out the order of the reaction in each and every reagent.

Order: In case of rate law of a reaction, the power to which the concentration of a reagent is increased. Or, the sum of powers on the concentration terms in the rate law.

Quench: To quickly stop a reaction via one of several signifies that comprise flash freezing and adding an inhibitor. Or to rapidly drive a reaction for completion, by adding the limiting reagent in surplus amount.

Rate: The speed of a reaction computed in amount or reagent consumed or product generated per unit time.

Rate Constant: It is the proportionality constant in the rate law expression. This factor is an assessment of the intrinsic reactivity of the reaction however is not constant with respect to the temperature.

Rate Law: It is an expression of the dependence of the rate of a reaction on the concentrations of reactants.

Factors affecting Reaction Rates:

Most of the factors affect rates of chemical reactions, and these are concluded below.

1) Nature of Reactants:

Formation of salts, acid-base reactions and exchange of ions are the examples of fast reactions. Reactions in which large molecules are made or break apart are generally slow. Reactions breaking strong covalent bonds are as well slow.

2) Temperature:

Generally, the higher the temperature, the faster is the reaction. The temperature effect is illustrated in terms of activation energy.

3) Concentration Effect:

Dependency of reaction rates on the concentrations is known as rate laws. Rate laws are the expressions of rates in terms of concentrations of reactants. Remember that rate laws can be in differential forms or in integrated forms. They are known as differential rate laws and integrated rate laws. The given is a short summary regarding rate laws.

Rate laws: Differential and integrated rate laws.

Integrated rate laws: First Order Reactions and Second Order Reactions

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

4) Heterogeneous reactions: The reactants are present in more than one stage or phase. For heterogeneous reactions, the rates are influenced by surface areas.

5) Catalysts: These are the substances utilized to facilitate reactions. By the nature of term, catalysts play significant roles in the chemical reactions.

Units of rate constant:

A point that frequently seems to cause endless confusion is the fact that the units of rate constant based on the form of the rate law in which it appears that is, a rate constant appearing in the first order rate law will encompass different units from a rate constant appearing in the second order or third order rate law. This follows instantly from the fact that the reaction rate for all time consists of the similar units of concentration per unit time that should match the overall units of a rate law in which concentrations increased to varying powers might appear. The good news is that it is extremely straightforward to find out the units of a rate constant in any particular rate law. Below are some illustrations.

i) Let suppose the rate law ν = k[H2][I2]. If we replace with units into the equation, we get

(mol dm-3 s-1) = [k] (mol dm-3) (mol dm-3)

Here, the notation [k] signifies 'the units of k'. We can reorganize this expression to determine the units of the rate constant, 'k'.

[k] = (mol dm-3 s-1)/(mol dm-3) (mol dm-3) = mol-1 dm3 s-1

ii) We can apply the similar treatment to a first order rate law, for illustration

ν = k [CH3N2CH3].

(mol dm-3 s-1) = [k] (mol dm-3)

 [k] = (mol dm-3 s-1)/(mol dm-3) = s-1

iii) As a final example, consider the rate law ν = k [CH3CHO]3/2.

(mol dm-3 s-1) = [k] (mol dm-3)3/2

[k] = (mol dm-3 s-1)/(mol dm-3)3/2 = mol-1/2 dm3/2 s-1

Determining the rate law from experimental data:

A kinetics experiment comprises of measuring the concentrations of one or more reactants or products at a number of various times throughout the reaction. We will evaluation some of the experimental methods used to make such measurements.

i) Isolation method:

The isolation process is a method for simplifying the rate law in order to find out its dependence on the concentration of a single reactant. Once the rate law has been simplified, the differential or integral procedures illustrated in the given subsections might be employed to find out the reaction orders.

The dependence of reaction rate on the selected reactant concentration is isolated by having all other reactants present in a big excess, in such a way that their concentration remains essentially constant all through the course of the reaction. As an illustration, consider a reaction A + B → P, in which B is present at a concentration 1000 times more than A. Whenever all of species A has been employed, the concentration of B will merely have changed via 1/1000, or 0.1%, and therefore 99.9% of the original B will still be present. It is thus a good approximation to treat its concentration as constant all through the reaction.

This very much simplifies the rate law as the (constant) concentrations of all the reactants present in large excess might be combined by the rate constant to result a single effective rate constant. For illustration, the rate law for the reaction considered above will become:

ν = k [A]a[B]b ≈ k [A]a[B]0b = keff[A]a with keff = k[B]0b

Whenever the rate law consists of contributions from a number of reactants, a sequence of experiments might be carried out in which each and every reactant is isolated in turn.

ii) Differential methods:

Whenever we encompass a rate law which based only on the concentration of one species, either as there is merely a single species reacting, or as we have utilized the isolation process to manipulate the rate law, then the rate law might be written:

ν = k[A]a

log ν = log k + a log[A]

The plot of logν against log[A] will then be a straight line having a slope equivalent to the reaction order, 'a', and an intercept equivalent to log k. There are mainly two ways in which to get data to plot in this manner.

a) We can evaluate the concentration of the reactant [A] as the function of time and use this data to compute the rate, ν = -d[A]/dt, as a function of [A]. A plot of logν vs log[A] then results the reaction order with respect to A.

b) We can prepare a sequence of measurements of the initial rate ν0 of the reaction by various initial concentrations [A]0. These might then be plotted as above to find out the order, a.

This is a generally used method termed as the initial rates procedure.

iii) Integral methods:

If we have evaluated the concentrations as a function of time, we might compare their time dependence by the suitable integrated rate laws. Again, this is the most straightforward if we have simplified the rate law in such a way that it depends on just one reactant concentration. The differential rate law given in equation ν = k[A]a will give mount to various integrated rate laws based on the value of a. The most generally encountered ones are:

Zeroth order integrated rate law: [A] = [A]0 - kt

                    A plot of [A] vs. t will be linear, having a slope of -k.

First order integrated rate law: ln[A] = ln[A]0 - kt

                    A plot of ln[A] vs. t will be linear having a slope of -k.

Second order integrated rate law: 1/[A] = 1/[A]0 + 2kt

                    A plot of 1/[A] vs. t will be linear having a slope of 2k.

If none of such plots yield in a straight line, then more complex integrated rate laws should be tried. 

iv) Half lives:

The other way of finding out the reaction order is to investigate the behavior of the half life as the reaction carries on. Particularly, we can assess a sequence of successive half lives. t = 0 is employed as the beginning time from which to assess the first half life, t1/2(1). Then t1/2(1) is employed as the start time from which to assess the second half life, t1/2(2) and so forth.

=> Zeroth order t1/2 = [A]0/2k

As at t1/2(1), the new starting concentration is 1/2[A]0, successive half lives will reduce via a factor of two for a Zeroth order reaction.

=> First order t1/2 = ln2/k 

There is no dependence of the half life on concentration; therefore t1/2 is constant for a first order reaction.

=> Second order t1/2 = 1/k[A]0

The inverse dependence on concentration signifies that successive half lives will two times for the second order reaction.

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