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 = log10A - 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.