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Chemical kinetics

This document summarizes key concepts in chemical kinetics including: 1. Chemical kinetics deals with the rates of chemical reactions and factors that affect rates. Rates can be expressed in terms of changes in concentrations over time. 2. Reaction rates depend on temperature, catalysts, and concentrations as expressed in rate laws. Rate constants are measures of reaction speeds. 3. Reactions can have different orders depending on which concentrations influence the rate. Pseudo-order reactions appear higher order but actually follow lower order kinetics.

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Chemical kinetics
From this chapter in board exam • Either 12 marks or 7 marks • 5 marks numerical sure
Chemical kinetics: The branch of chemistry which deal with the study of speeds or rates of chemical reaction, the factors affecting the rate of the reaction and the mechanism by which the reaction proceed is known as chemical kinetics.
Concept of the reaction rate: The rate of the reaction is a measure of the speed with which reactants are converted into products. In terms of amount, the rate of the reaction is defined as the change in amount of reactant or product per unit time. Mathematically, Rate of the reaction= change in amount of reactants or products time taken for the changed
In terms of concentration, the rate of the reaction is defined as the change in the concentration of the reactant( or product) per unit time. Rate of reaction= change in concentration of reactants or products time taken for change In this case, unit of the rate is molL-1time-1 i.e molL-1s-1 or molL-1min-1 or molL-1hr-1.
Therefore rate of reaction is defined as the rate of disappearance (or decrease) in the concentration of the reactants or the rate of appearance (or increase) in the concentration of the products. For any hypothetical reaction, A → B Rate of reaction= decrease disappearance in concentration of reactants(A) time taken or, = increase appearance in concentration of products(B) time taken Or, rate of the reaction = - ∆[A] ∆T = + ∆[B] ∆T = change in concentration time needed for the change
Types of the reaction: Depending upon the time interval over when the change in amount of reactant or product is measured the reaction rates can be classified into two main classes: 1. Average rate of the reaction 2. Instantaneous rate of the reaction
1. Average rate of the reaction: It is defined as the rate of change of concentration of the reactants (or products) over a specified measurable period of time. For a reaction, A → B If [A]1 and [A]2 be the concentration of the reactants A at time t1 and t2 respectively, then the rate of the reaction in terms of the reactant in time interval (t1 – t2) is given by, Average rate of reaction = change in concentration of A time interval = A 2 − A 1 t2 −t1 = - ∆[A] ∆t
Here [A]1 > [A]2 Therefore negative is used. Similarly, if [B]1 and [B]2 be the concentration of product B at time t1 and t2 respectively, then the rate of reaction in terms of product B is given by, Average rate of the reaction = change in concentration of B (products) time interval = B 2 − B 1 t2 −t1 = + ∆[B] ∆t Here,[B]2 > [B]1, therefore positive sign is used Figure……….. in board
2. Instantaneous rate of the reaction: It is defined as the rate of change in concentration of any one of the reactant or product at particular time. It gives the tendency of the reaction at a particular point of time during the course. Mathematically, Instantaneous rate =[Average rate]∆t → 0 𝑑𝑥 𝑑𝑡 = Where dx= very small change in concentration dt= very small change in time For instantaneous rate, t1 ≈ t2 ∆t → 0.
Calculation of instantaneous rate of reaction: To know the rate of the reaction at any time , a tangent is drawn to the curve at the point corresponding to that time and it is extended on either side so as to cut the axes. Then Rate of reaction= change in concentration time Figure ……………….in board
Expressing the rate of the reaction: Example, aA + bB → cC + dD The expression for the rate of the reaction is Rate, R = - 1 a d[A] dt = - 1 b d[B] dt = + 1 c d[C] dt = + 1 d d[D] dt Rate of disappearance of A, - 𝐝[𝐀] 𝐝𝐭 = a × rate of the reaction Rate of disappearance of B, - 𝐝[𝐁] 𝐝𝐭 = b × rate of the reaction Rate of appearance of C, + 𝐝[𝐂] 𝐝𝐭 = c × rate of the reaction Rate of appearance of D, + 𝐝[𝐃] 𝐝𝐭 = d × rate of the reaction
Numericalfromelementarychemicalcalculations
Rate law and rate law equation: For a general reaction, aA + bB → products Then according to the law of mass action, Rate ∝ A a B b Or, rate = K[A]a[B]b where K is constant called velocity constant. However, experimentally is observed that the rate of this reaction may or may not depend upon all the ‘a’ concentration terms A and all the ‘b’ concentration terms of B. Suppose, experimentally the rate of the reaction is found to depend upon ‘x’ concentration terms of A and ‘y’ concentration terms of B.
Then we may write, Rate ∝ A x B y where, x is the order with respect to A and y is the order with respect to B. Rate = K[A]x[B]y…………………..(i) Where [A] and [B] are the molar concentration of A and B respectively and K is constant called rate constant. The above expression is called rate law. The rate law is defined as “ The mathematical expression which represents the experimentally measured rate of the reaction in terms of the concentration of the reacting species which influence the rate of the reaction.”
Again, if all concentrations are taken as unity i.e. [A]=[B]= 1 mol/L Then rate = K(rate constant) Hence, rate constant ,may be defined as the rate of the reaction when the concentration of each reactant is taken as unity.
Characteristics of the rate constant(K) i. Rate constant is a measure of the rate of reaction. Higher the value of rate constant faster is the reaction. ii. The value of rate constant ‘K’ is different for different reactions. Hence it is also called specific constant. iii. For a particular reaction, the value of rate constant is independent of the concentration of the reactants.(zero order reaction)
Order of the reaction: The sum of the concentration terms on which the rate of a reaction actually depends as observed experimentally is called order of the reaction. Or, The power to which the concentration of the reactants is raised in the rate law equation is called order of the reaction. Examples ,let us consider a general reaction, mA + nB → product ………….in board
Differences between rate of reaction and rate constant. Rate of reaction Rate constant 1. It is the change in concentration of the reactant or product per unit time. 1. It is a constant of proportionality in the rate law equation and is equal to the rate of the reaction when the molar concentration of each reactants is unity. 2. The rate of reaction at any instant of time depends upon the molar concentration. 2. The rate constant is constant for a particular reaction at a particular temperature and may or may not depend upon the concentration of the reactants. 3. Its units are always molL-1time-1 3. Its units depend upon the order of the reaction
Molecularity of a reaction: The number of atoms, ions or molecules that must collide with one another simultaneously so as to result into a chemical reaction. Here molecularity of a reaction is simply the number of reacting species taking part in a simple reaction i.e. single step reaction. For example, 1. Decomposition of O2F2: The balanced equation is O2F2 → O2 + F2 Hence the molecularity of the reaction is 1 and is called unimolecular. 2. Dissociation of HI: The balanced equation is 2HI → H2 + I2 Hence the molecularity is 2 and the reaction is called bimolecular reaction.
3. Reaction between NO and O2: The balanced equation is 2NO + O2 → 2NO2 Hence the molecularity is 3 and the reaction is called termolecular. However there are many reactions in which all the atoms, ions or molecules as represented in the balanced chemical reaction may not come together to collide simultaneously. For example in the reaction 5Br- + BrO3 - + 6H+ → 3Br2 + 3H2 It is impossible for all 12 ions of the reactant to come together simultaneously to collide. These reactions are supposed to take part in a sequence of number of steps. Each of these step reactions is called elementary reaction. Molecularity can be defined only for an elementary reaction. It has no meaning for the complex reactions.
Differences between molecularity and order of the reaction: (VVI) Molecularity of the reaction Order of the reaction 1. It is the number of reacting species taking part in a simple reaction. 2. It is the theoretical value determined from the balanced single step reaction. 1. It is the sum of the powers of the concentration terms in the rate law equation. 2. It is the experimentally determined value. 3. It has natural number values i.e. 1, 2, 3 etc. 3. Order of the reaction can be fractional values also. 4. Molecularity of the reaction cannot be zero. 4. It can be zero 5. Molecularity of a complex reaction has no meaning. It is expressed for each elementary step. 5. Order is same for the whole reaction no matter it is simple or a complex reaction.
Pseudo-order reaction: The reactions which appear to be higher order but follow lower order kinetics are called pseudo-order reactions. In case of elementary reactions (which complete in one step), the order and molecularity are identical. Let us consider a general reaction. A + B → Products In which the reactant B is present in a large excess. Since it is elementary, its rate law can be written as Rate = K[A][B] As B is present in large excess, its concentration remains practically constant during the course of the reaction. Thus the rate law can be written as Rate = K’[A]
Where the new rate constant K’=K[B]. Thus the actual order of the reaction is second order but in practice it is of first order. Therefore the reaction is said to be pseudo-first order. Examples: a. Hydrolysis of an ester: CH3COOC2H5 + H2O(Large excess) → CH3COOH + C2H5OH The rate law for the reaction can be written as Rate =K[CH3COOC2H5][H2O] Since water is present in large excess, its concentration remains practically constant during the course of the reaction. Thus above rate law can be written as Rate = K’[CH3COOC2H5] The reaction is actually second order but in practice it follows first- order kinetics. Thus, it is a pseudo-first order reaction.
Concept of activation energy: The excess energy which must be supplied to the reactants to undergo chemical reactions is called activation energy (Ea). It is equal to the difference between the threshold energy needed for the reaction and the average kinetic energy of all the reacting molecules. Activation energy= Threshold energy - Average kinetic energy of the reacting molecules (The minimum kinetic energy which the colliding molecules must have in order that the collision between them may be affective is called threshold energy.)
Figure in board: Activated complex: It is an intermediate state that is formed during the conversion of reactants into products.
According to the concept of activation energy, the reactants do not change directly into the product. The reactants first absorb energy equal to the activation energy and thereby form activated complex. At this state, the molecules have energy at least equal to the threshold energy. This means that the reaction involves some energy barrier which must be overcome before products are formed. The energy barrier is known as activation energy barrier. Example:
Effect of catalyst: The use of catalyst provides an alternative path for a reaction by lowering activation energy( positive catalyst) or highering(increasing) the activation energy(negative catalyst). Figure in board
A catalyst itself is not used in a reaction. So, a small amount of catalyst is sufficient to catalyse the reaction. After the completion of the reaction, the chemical composition of the catalyst remain unchanged. The phenomenon of catalyst is specific i.e. a particular catalyst can catalyse a particular reaction.
Effect of temperature on the reaction rate: Rate of the reaction increase on increasing the temperature. In general, it has been found that, an increase in temperature by 10oc doubles the rate of the reaction. In 1889 Arrhenius has derived a simple relation between rate constant and the temperature of the system. K = Ae-Ea/RT ………………………(i) where A =collision frequency Ea= activation energy R= universal gas constant T= Kelvin temperature (Collision frequency: The number of collision that take place per second per unit volume of the reaction mixture is called collision frequency)
In equation (i) the power consists of negative. It means with decrease in activation energy, the value of the rate constant increases and vice versa. Similarly, with increase in temperature the value of the rate constant (K) also increases. It is because with increase in temperature, the KE of reacting molecules also increases. So more number of molecules can cross the energy barrier and the rate of reaction increases.
Collision theory of bimolecular reaction: According to this theory, a chemical reaction takes place only by collision between the reacting molecules. But not all collision are effective. Only a small fraction of the collision can be change reactant into product. Two main criteria for the reaction are: 1. The colliding molecules must possesses sufficient kinetic energy to cause the reaction. 2. The reacting molecules must collide with proper orientation.
To explain collision theory let us consider a general reaction: A-A + B-B → 2AB To give the product, there must be breakdown of bonds between the atoms in the reactant molecules. The energy for the breakage of bonds comes from the kinetic energy possessed by the reacting molecules from collision. Fig……… Only those molecules which collide with K.E greater than or equal to Ea are able to cross the barrier and give product. The molecules colliding with KE less than Ea fail to cross the energy barrier. Such type of collision between the molecules are unproductive.
The reacting molecules must collide with proper orientation. The correct orientation is that which ensures direct contact between the atoms involved in the breaking and forming the bonds. Effective collision: The collision between reacting molecules resulting in the formation of product is called effective collision.