Chemical Kinetics

## Rate of Reaction

**Average Rate**

Average Rate = - ∆[A]/∆t = ∆[B]/∆t

**Instantaneous Rate**

Instantaneous Rate = - d[A]/dt = d[B]/dt

For every 10 K rise in temperature , instantaneous rate increases by 2 - 3 times . Temperature co-efficiebt of rate lies between 2 and 3 .

Rate is directly proportional to concentration and pressure .

## Rate Law

aA ----> bB

Rate = k[A]x[B]y

k is independent of initial concentration of reactants , but depends upon the temperature and type of reaction .

Unit of Reaction for nth order reaction= (mol)1-n (dm3)n-1 s-1

## Order of Reaction

The x y in the above example is called as partial order of reaction . x + y is the overall order of reaction .

## Integrated Rate Laws

The equations which are obtained by integrating the differential rate laws and which give a simple relationship between concentration , time and rate constant , are called as Integrated rate laws .

k = (2.303/t)( log10[A0]/[At])

Half life of a reaction is the time taken for the concentration of reactants to become half the original value .

t1/2 = 0.693 / k

After every half life period , the concentration becomes half .

## Graphical Representation of First Order Reactions

The integrated rate laws can be represented in the slope intercept form . y = mx + c .

1) Rate v/s Concentration

Rate = k[At]

y = Rate , x = [At]

Slope = k  ; y-intercept = 0

2) Concentration v/s time

3) log10(At) v/s time

log10[At] = - kt / (2.303) + log10[A0]

y = log10[At] ; x= time

Slope = - k / 2.303 ; y-intercept = log10[Ao]

4) log10([Ao]/[At]) v/s time

kt /2.303 = ( log10[A0]/[At])

y = log10([Ao]/[At])  ; x = time

Slope = k / 2.303 ; y-intercept = 0

## Zero Order Reactions

Zero Order reactions are those reactions , whose rate does not depend on the concentration of reactants .

## Pseudo First Order Reaction

Pseudo First Order Reactions are those reactions which do not have the order of 1 , but behave like first order reactions .

## Molecularity

Molecularity of a reaction is the number of molecules participating in a reaction . It can also be the number of reactant molecules .

## Reaction Intermediate & Catalyst

The reaction intermediate is formed in the first step and consumed in the second step .

The reaction intermediate is consumed in the first step and formed in the second step .

A catalyst increases the rate of reaction and can be recovered back unchanged after the completion of reaction . Catalysts provide an alternative path of lower activation energy , so that more reactant molecules are able to reach the energy barrier (f) .

## Arrhenius Equation

The Arrhenius Equation gives us the temperature dependence of of reaction rates .

log10 k = log10 A - Ea/2.303 RT

log10(k2/k1) = Ea/ 2.303 R [T2-T1/T1T2]

## Collision Theory

The basic requirement of collision theory is that the reactant molecules should come together and collide in order for the reaction to occur .

## Activation Energy

The activation Energy (Ea) is the minimum kinetic energy required for a molecular collision to lead to reaction .

## Orientation

If the molecules are not properly oriented , they do not react even at the activation energy . The molecules must be oriented relative to each other such that the groups reacting or bonds to be shifted are relatively close .