The document discusses various topics related to chemical kinetics including: - Rate of reaction is defined as the change in concentration of reactants or products over time. Rate laws relate the rate of reaction to the concentrations of reactants. - Reaction order refers to the sum of powers in the rate law. Molecularity is the actual number of reacting species. Pseudo-orders occur when one reactant is in excess. - Rate constants have different units for different reaction orders. Integrated rate laws and the half-life method can be used to determine the order of a reaction experimentally. - Collision theory states that molecules must collide with sufficient energy and correct orientation to react. Physical factors like temperature, solvent, and

1. Chemical kinetics
V2
(Dr.) Mirza Salman Baig
Assistant Professor (Pharmaceutics)
AIKTC, School of Pharmacy,New Panvel
Affiliated to University of Mumbai (INDIA)

2. Rate of reaction
• A → B
• Rate of reaction is defined as the change
in concentration of any of reactant or
product per unit time.
– Rate of reaction = Rate of disappearance of A
– Rate of reaction = Rate of appearance of B
• Unit = M/s
• Rate= dA/dt

3. Rate of reaction
• Rate of reaction = - d[A] /dt
• Rate of reaction = + d[B] /dt

4. Rate law
• The Equation which shows how rate is
related to concentration
Reaction
• 2NO + 2H2 → N2 + 2H2O
Rate Law
• Rate of reaction = k[H2] [NO]2

5. Rate Law
• The rate of reaction is directly proportional
to the reactant concentration, each
concentration being raised to some power.
• 2A + B → Product
• Rate =k[A]m[B]n
(k = specific rate constant)

6. Order of reaction
• Order of reaction is defined as sum of
powers of concentration in rate law.
• Rate =k[A]m[B]n
• Order of reaction in above case is (m+n)
• It is the number of concentration of
chemical species which determine the rate
of reaction depends.

7. Molecularity of reaction
Number of molecules (chemical species) which
take part in chemical reaction. It is number
(integer) not a fraction.
Types
• Unimolecular
cis to trans
• Bimolecular
A+B ---> C
• Termolecular
A+B+C --> D

8. Reaction
Order
• Sum of powers of
concentration terms in
rate law
• Experimentally
determined
• It can have fractional
value
• It can assume zero
value
Molecularity
• Number of reacting
species in a reaction
• Therotical concept
• It is always a whole
number
• It cannot have zero
value

9. Order of reaction wrt A

10. Overall order of reaction

11. First Order Reaction
• A--> Product
• At time t=0 concentration of reactant A is a
mole/lit,
• After time t, x mole of A have changed, the
final concentration after time t will be (a-x)
• For first order reaction dx/dt is directly
proportional to (instantaneous)
concentration of reactant...
• dx/dt= k(a-x)
• dx/(a-x)= kdt

12. First Order Reaction
• dx/(a-x)= kdt …. (1)
• Integrating above equation..
• ∫dx/(a-x)= ∫kdt
• -ln (a-x) = kt + I …. (2)
If t=0, x=0 then I= -ln a
Substituting I in equation (2)
• ln [a/(a-x)]= kt or k = (1/t) ln [a/(a-x)]
• Converting it to common log
• k = log

13. First Order Reaction
• After integration
• K= 2.303/t . Log a/(a-x)
• k= 2.303/(t2-t1) . log (a-x1)/(a-x2)
• t1 and t2 are time interval at which x1 and
x2 amount of reactant changed
respectively

14. Second Order Reaction
Reaction A--> Product
Initial conc a 0
Final conc a-x x
• For second order reaction rate of reaction
is proportional to square of concentration
of reactants
• dx/dt= k (a-x)2

15. Second Order Reaction
• dx/dt= k (a-x)2
• dx/(a-x)2=k dt
• On integration ∫ dx/(a-x)2= ∫k dt
• 1/(a-x) = kt +I
At t=0, x=0.... I = 1/a

16. Second Order Reaction
• Substituting value of I in above equation
• 1/(a-x) = kt + 1/a
• kt = 1/(a-x) - 1/a
• k = .

17. Zero Order reaction
• Rate of reaction is independent of
concentration of reactant/product
• Time A --> Product
• Initial conc 0 a 0
• Final conc t a-x x

18. Zero Order reaction
• Rate of reaction = -d[A]/dt= k0[A]0
• dx/dt= k0 (a-x)0
• x=k0t
• k0 is rate constant (or specific rate
constant) of zero order reaction
• Rate constant is the rate of reaction at all
concentrations x/t

19. Pseudo order reaction
• Experimental order of reaction which is not
actual is known as pseudo order
• Reaction A+B → Product
• If B is in excesses, its concentration will
practically constant and only concentration
of A will affect rate of reaction hence rate
law will be..
• Rate = k' [A]...

20. Pseudo order reaction
(Example)
• Ethyle acetate+ Water (excesses) -->
Acetic acid + EtOH

21. Units of rate constant
• Units of rate constant for different order
reaction are different
• For Zero order
• k= d[A]/dt
• k= mol/lit . 1/time
• k= mol lit-1 time-1

22. Units of rate constant
• For first order
• k= 2.303/t . log a/(a-x)
• k= 2.303/t . log [A]0/[A]t
• k= 1/time
• k= time-1

23. Units of rate constant
• For second order reaction
• k= 1/t . x/a(a-x)
• k= 1/t . x/[A]0([A]0-x)
• k= 1/time . concentration/concentration2
• k= 1/time . 1/concentration
• k= 1/time . 1/(mol/lit)
• k= mol-1 lit time-1

24. How to determine order of reaction
• Using integrated rate equations
• Graphical Method
• Using half life period
• Oswald isolation method

25. Using integrated rate equations
• Perform the reaction using different initial
concentration of of reactant (a) and note
the concentration (a-x) at regular time
intervals (t)
• These values are then substituted in
integrated rate equations (first order,
second order, zero order)
• The rate equation which yield constant
value for k corrosponds to that order of
reaction.

26. Graphical Method
• For straight line y=mx+c
Log (a-x) = -kt + Log a
• First Order
• If the plot of log (a-x) vs t is a straight line,
the reaction follows first-order .
t

27. Graphical Method
• For straight line y=mx+c
1/(a-x) = kt + 1/a
• Second Order
• If the plot of 1/(a-x) vs t is a straight line, the
reaction follows second order
t

28. Graphical Method
x=k0t

29. Half life reaction
• It is defined as the time required as for the
decrease in concentration of reactant to
half of its initial value.
• When x= a/2, then time is termed as t1/2
• We can substitute this value in equation of
reaction

30. Using Half life method
• Seperate experiments should be
performed using different initial
concentration. Let initial concentration be
A1 and A2 and corresponding half life t1
and t2
• Half life for nth order reaction is
• t1/2→ 1/ [A] n-1

31. Using Half life method
• Zero order: Half life is directly proportional
to initial concentration of the reactants
• First order: Half life is independent of the
initial concentration of reactants
• Second order: Half life is inversly
proportional to initial cincentration of
reactants

32. Order of
reaction
Equation t1/2 (Half life)
0 x=k0t t1/2= a/2k0
1 k= 2.303/t . log a/(a-x) t1/2= 0.693/k
2 k = 1/t . x/a(a-x) t1/2= 1/ak

33. Collision Theory of Reaction Rate
• According to this theory, chemical reaction
take place only when there is collision
between reacting molecules.
• Colliding molecules must possess enough
kinetic energy
• A-A + B-B --> 2 A-B
• Molecules must collide with correct
orientation

34. Energy of activation Ea
• Minimum amount of
energy (Ea)
necessary to cause
reaction between
two.
• Only molecules that
collide with kinetic
energy higher than
Ea could react.

35. Correct orientation

36. Limitations of collision theory
• It applies to simple gaseous reactions only
• Values of rate constant (k) calculated by
Arrhenius equation are in agreement with
simple bimolecular reaction only (not for
complex molecules)
• Theory consider only (linier) kinetic energy
but not rotational, vibrational energy.
• Theory don’t speaks about breaking of
bonds involved in it.

37. Physical and chemical factors
influencing the chemical
degradation of pharmaceutical
product:
• Temperature
• Solvent & Dielectric constant
• Ionic strength
• Specific acid base catalysis
• General acid base catalysis

38. Effect of Temperature
(Arrhenius equation)
• Rate of reaction increase 3-folds by
increase in temperature by about 10oC
• Relation between rate constant,
temperature and Ea .....Arrhenius equation
• k= Ae-Ea/RT
• R= gas constant
• T = Temperature in Kelvin
• A= Factor related to frequency of collision
• Ea= Energy of activation
• e= mathematical const (Napier’s ) ≈ 2.71828

39. Effect of Temperature

40. Effect of Solvent
• Rate of reaction will be fast if product is
more soluble than reactant, in reaction
solvent.
• Rate of reaction will be slow if reactant are
more soluble in solvent than product.
• A→ B

41. Effect of solvent (Dielectric constant)
• Reactions involving ions of opposit charge are
accelerated by solvents with low dielectric constant
(ability of a substance to store electrical energy in an
electric field)
• Ex. Rate of hydrolysis of sulphate ester is greater in
low dielectric constant solvent like methylene chloride
than in water.
• Acid inversion of sucrose, rate of reaction increases if
dielectric constant (of water ) reduced (by adding
dioxane)
• Reaction between similar charged ions is favoured
by high dielectric constant solvents
• Reaction between neutral molecules which produce
highly polar transition state is favoured by high
dielectric constant solvents

42. Effect of Ionic strength
• Ionic strength may affect inter-ionic
attraction
• Increase in ionic strength expected to
decrease (not favorable) rate of reaction
between oppositely charged ions
• Increase in ionic strength expected to
Increase (favorable) rate of reaction
between similar charged ions... as per
Debye-Hukel equation

43. Catalyst
• It is a substance that influence speed of reaction
without itself being altered chemically.
• When catalyst decrease rate of reaction then it is
known as negative catalyst.
Mechanism :
• It may form intermediate complex with reactant
• May produce free radicles like *CH3
• Homogeneous catalyst→ Catalyst and reactant
are in same phase
• Heterogeneous catalyst→ ex. Finely divided
platinum

44. Specific acid base catalysis
• Specific acid base catalysis refer to catalysis by
hydrogen ion (H+) or Hydroxyl Ion (OH-) hence
(H+ and OH-) terms involved in rate law.
• Ex. Rate of hydrolysis of ester ethylacetate
• CH3-COOC2H5 → CH3COOH + C2H5-OH
• It is studied at constant pH (buffered soln), rate
of disappearance of ethyl acetate (ester) will
apparently First order.
• If reaction is studied at different pH (in acidic
range) then different first order rate constant k is
observed at different pH.

45. Specific acid base catalysis
• Observed rate depend on concentration of
ester and [H+], therefore it is actually
second order reaction.
• Observed rate constant (kobs) at different
pH in acidic region, is proportional to [H+]
• kobs = kacid [H+]
• logkobs = log kacid + log [H+]
• logkobs = log kacid - pH

46. Specific acid base catalysis
• Log kobs = log kacid - pH
• It suggest log kobs vs pH is straight line
with slop -1 and y-intercept log kacid
• If we study same reaction in alkaline pH
then we observe different rate constant at
different pH
• log kobs = log kbase + log [OH-]
• kobs = kacid[H+] + kbase [OH-]

47. 1 st order Hydrolysis of atropin

48. General acid base catalysis
• Acid or base catalysis is not restricted to
effect of [H+] or [OH-]
• Other components of buffer (which is
generally used in pharma formulations)
may cause catalysis.
• Undissociated acid and base can often
produce catalytic effect
• Metal can also serve as catalyst
• Ex. Mutarotation of glucose in acetate
buffer is catalysed by [H+], [OH-], Acetate
ion [CH3COO-], undissociated acid
[CH3COOH] undissociated acetic acid.

49. Problem
Q1. Solution of drug contained 500unit/ml of
drug when prepared. It was analyzed after
40 days and found to contain 300 unit/ml.
Assume decomposition is first order, at
what time will the drug have decomposed
to one half of its original concentration?

50. Ans 1
• a= 500 unit/ml
• (a-x) = 300 unit/ml
• t= 40 days
• k= 2.303/t . log a/(a-X)
• k= 2.303/40 . log 500/300
• k= 0.0576 × log 1.667
• k= 0.0576 × 0.222
• k= 0.01278 day-1

51. • t1/2 = 0.693/k .... for first order
• t1/2 = 0.693/ 0.01278
• t1/2 = 54.23 Days

52. Problem
• Q2 Consider the reaction: 2B → C + 3D. In
one experiment it was found that at 300 K the rate
constant is 0.134 L/(mol.s). A second experiment
showed that at 450 K, the rate constant was 0.569
L/(mol.s). Determine the activation energy for the
reaction

53. • Ans 2.
• Ea = 10.8 kJ

54. Thank You