2. WHAT IS CHEMICAL
KINETICS ?
• Chemical kinetics, also known as reaction kinetics, is
the branch of physical chemistry that is concerned
with understanding the rates of chemical reactions.
It is to be contrasted with thermodynamics, which
deals with the direction in which a process occurs but
in itself tells nothing about its rate. Chemical kinetics
includes investigations of how different experimental
conditions can influence the speed of a chemical
reaction and yield information about the reaction's
mechanism and transition states, as well as the
construction of mathematical models that also can
describe the characteristics of a chemical reaction.
3. RATE OF A CHEMICAL
REACTION
• The speed or rate of a chemical
reaction is the change in concentration
of a reactant or product per unit time.
To be specific, it can be expressed in
terms of: the rate of decrease in
concentration of any of the reactants.
the rate of increase in the
concentration of any of the products.
4. The factors that affect
reaction rates are:
• surface area of a solid reactant.
• concentration or pressure of a
reactant.
• temperature.
• nature of the reactants.
• presence/absence of a catalyst.
5. DEPENDENCE OF RATE
ON CONCENTRATION
• The rate of a chemical reaction at a
given temperature may depend on
the concentration of one or more
reactants and products. The
representation of rate of reaction
in terms of concentration of the
reactants is known as rate law. It is
also called as rate equation or rate
6. ORDER OF
REACTION
• The Order of Reaction refers to the
power dependence of the rate on the
concentration of each reactant. Thus,
for a first-order reaction, the rate is
dependent on the concentration of a
single species. ... The order of reaction
is an experimentally determined
parameter and can take on a fractional
value.
7. MOLECULARITY OF A
REACTION
• The molecularity of a reaction is
defined as the number of molecules or
ions that participate in the rate
determining step. A mechanism in which
two reacting species combine in the
transition state of the rate-determining
step is called bimolecular.
8. INTEGRATED RATE
EQUATION
• The rate law is a differential
equation, meaning that it describes
the change in concentration of
reactant (s) per change in time.
Using calculus, the rate law can be
integrated to obtain an integrated
rate equation that links
concentrations of reactants or
products with time directly.
9. ZERO ORDER
REACTION
• Zero order reaction means that the rate of the reaction is
proportional to zero power of the concentration of
reactants. Consider the reaction, R → P
• Rate = -d[R]/dT = k[R]⁰
• As any quantity raised to power zero is unity
• Rate = = -d[R]/dT
• d[R] = – k dt
• Integrating both sides [R] = – k t + I
• where, I is the constant of integration.
• At t = 0, the concentration of the reactant R = [R]ₒ,
where [R]ₒ is initial concentration of the reactant.
• [R]ₒ= –k × 0 + I
• [R]ₒ = I
• Substituting the value of I in the equation
10. • [R] = -kt + [R]ₒ
• Comparing (with equation of a straight line, y = mx + c, if we plot [R]
against t, we get a straight line with slope = –k and intercept equal to
[R]ₒ
•
• Zero order reactions are relatively uncommon but they occur under
special conditions. Some enzyme catalysed reactions and reactions
which occur on metal surfaces are a few examples of zero order
reactions. The decomposition of gaseous ammonia on a hot platinum
surface is a zero order reaction at high pressure
• 2NH₃(g) (at 1130k in presence of pd as catalyst)------->N₂(g) + 3H₂(g)
In this reaction, platinum metal acts as a catalyst. At high pressure, the
metal surface gets saturated with gas molecules. So, a further change in
reaction conditions is unable to alter the amount of ammonia on the
surface of the catalyst making rate of the reaction independent of its
concentration. The thermal decomposition of HI on gold surface is another
example of zero order reaction
11. FIRST ORDER REACTION
• In this class of reactions, the rate of the reaction is
proportional to the first power of the concentration of
the reactant R.
• For example,
• R → P
• Rate = d[R]/Dt = -k[R]
or
-k dT
• Integrating this equation, we get
• ln [R] = – kt + I
• Again, I is the constant of integration and its value can be
determined easily.
• When t = 0, R = [R]ₒ, where [R]ₒs the initial concentration
of the reactant.
12. • Therefore, equation can be written as
• ln [R]ₒ= –k × 0 + I
• ln [R]ₒ= I
• Substituting the value of I in equation
• ln[R] = -kt + ln[R]ₒ
• Rearranging this equation
• lnR/Rₒ=kt
or
K=(1/t)ln[R]ₒ/[R]
• At time t1
• *ln[R]₁= – kt1 + *ln[R]ₒ
• At time t2
• ln[R]₂ = – kt₂+ ln[R]ₒ
• where, [R]₁ and [R]₂ are the concentrations of the
reactants at time t1 and t2 respectively
• ln[R]₁– ln[R]₂= – kt₁ – (–kt₂)
13. • On simplifying we get
• lnR/Rₒ =Kt
• Taking antilog of both sides
• [R] = [R]ₒ e⁻ᴷᵗ
• Comparing with y = mx + c, if we plot ln
[R] against t we get a straight line with
slope = –k and intercept equal to ln [R]ₒ
14. HALF LIFE OF A
REACTION
• Another approach to describing reaction
rates is based on the time required for the
concentration of a reactant to decrease to
one-half its initial value. This period of
time is called the half-life of the reaction,
written as t1/2. Thus the half-life of a
reaction is the time required for the
reactant concentration to decrease from
[A]0 to [A]0/2. If two reactions have the
same order, the faster reaction will have a
shorter half-life, and the slower reaction
will have a longer half-life.
15. • The half-life of a first-order reaction under a
given set of reaction conditions is a constant.
This is not true for zeroth- and second-order
reactions. The half-life of a first-order reaction
is independent of the concentration of the
reactants.
This becomes evident when we rearrange the
integrated rate law for a first-order reaction to
produce the following equation:
• ln[A]ₒ/[A]=kt (1)
• ln[A]ₒ/[A]=kt
• Substituting [A]0/2 for [A] and t1/2 for t (to
indicate a half-life) into Equation 11 gives
• ln[A]0[A]0/2=ln2=kt1/2 (2)
16. • Substituting ln2≈0.693ln2≈0.693 into
the equation results in the expression
for the half-life of a first-order
reaction:
• t1/2=0.693k (3)
