This document discusses the topic of chemical kinetics, which deals with the rates and mechanisms of chemical reactions. It provides definitions for key concepts such as reaction rate, factors that affect reaction rates like concentration and temperature, and order of reactions. Catalysts are also explained as substances that speed up reactions without being consumed. The document then covers chemical equilibrium, including definitions of the equilibrium constant, Le Chatelier's principle, and the relationship between the equilibrium constant and Gibbs free energy.

1. 1
Chemical Kinetics

2. Chemical Kinetics is the branch of physical chemical which
deals with the rate and mechanisms of chemical reactions.

4. The reaction rate (R) for a reactant or product in a particular
reaction is intuitively defined as how fast or slow a reaction
takes place:
R=
t
C
∆
∆
±
С
t
R – the increase in C(x) of
products of a reaction per unit
of time or the decrease in C(x)
of reactants per unit of time.
1
2
1. Forward
2. Backward

5. Factors That Affect Reaction Rates
1) the physical state (nature) of the reactants,
2) the concentrations of the reactants,
3)the temperature at which the reaction occurs,
4) and whether or not any catalysts are present in the
reaction.

6. 2. Concentration plays a very important role in
reactions, because according to the collision theory of
chemical reactions, molecules must collide in order to
react together.

7. As the concentration of the reactants increases, the
frequency of the molecules colliding increases, striking each
other more frequently by being in closer contact at any given
point in time.

8. The rate of a chemical reaction at any instant at a given t° is
proportional to the active mass at that instant of each of the reactants
present in the systems; the active mass in a homogeneous system, is
defined as the number of moles of the substance present per unit
volume (mol/l – molar concentration).
The Law of Mass Action.
The effect of mass on the rate of reaction was studied by Guldberg
and Waage who in 1864 stated Law of mass action as follows.

9. Homogeneous reactions are chemical reactions in
which the reactants are in the same phase.
2CO(g) + O2(g) 2CO2 (g) R1 =k1 [CO]2
[O2]
R2 =k2 [CO2]2

10. Heterogeneous reactions have reactants in two or more
phases. Reactions that take place on the surface of a catalyst
of a different phase are also heterogeneous.
! Solid state is not on account
S(s) + O2 (g) SO2 (g) R1 =k1[O2]
R2 =k2[SO2]

11. 3. Effect of Temperature.
Vant – Hoff’s Rule.
A rise in t° leads to a tremendous increases in reaction rate.
The rate of many reactions increases 2-4 times for a 10°C
rise in t°.

13. Ea
Most reactions involving neutral molecules
cannot take place at all until they have
acquired the energy needed to stretch, bend,
or otherwise distort one or more bonds. This
critical energy is known as the activation
energy (Ea) of the reaction.

14. Finally, in 1899, the Swedish chemist Svante Arrhenius (1859-
1927) combined the concepts of activation energy and the
Boltzmann disribution law into one of the most important
relationships in physical chemistry:

15. It is often convenient to estimate the activation energy from
experiments at only two temperatures.

16. Activation energy plots.
Activation energy diagrams can describe both
exothermic and endothermic reactions:

17. A catalyst is usually defined as a substance that
speeds up a reaction without being consumed by it.
More specifically, a catalyst provides an alternative,
lower activation energy pathway between reactants
and products. Most biochemical processes that
occur in living organisms are mediated by enzymes,
which are catalysts made of proteins.
4. Effect of catalysis.

19. catalysts affect the forward and reverse rates
equally; this means that catalysts have no effect on the
equilibrium constant.

20. Catalysts are conventionally divided into two categories:
homogeneous and heterogeneous. Enzymes, natural
biological catalysts, are often included in the former group,
but because they share some properties of both but exhibit
some very special properties of their own, we will treat them
here as a third category.

21. Mechanism of catalysis.
There are many types of mechanisms of catalysis .
The catalyst combines with one of the reactants to form a more
reactive but unstable intermediate compound which then reacts with
other reactants to yield the products.
1) 2NO + O2
2NO2
2) NO2
+SO2
SO3
+ NO cat.
2SO2
+ O2
2SO3
NO

22. Molecularity - is the member of molecular taking part in a
simple reaction or as the number of molecules taking part in
the rate determining step (slowest step) of a complex
chemical reaction.
By molecularity reactions may be classify to:
Unimolecular : Ca CO3
t°
CaO + CO2
Dimolecular : H2
+ Cl2
2HCl
Trimolecular : 2CO + O2
2CO2

23. The Order of an elementary reaction. The term order of a
reaction is used to denote the dependence of
experimentally determined reaction rate on concentration.
It is the sum of the exponents of concentration term
in the experimentally observed rate reaction.

24. The First order reactions.
The rate of the reaction will be directly proportional to the
concentration of the reactant.
First order main formulas:
C0 -initial concentration
C τ – final concentration
τ – time
K=time-1
R = kC

25. Half – life of a reaction (τ0.5)
is the time it takes for the reactant concentration to
decrease to one half of its initial value.

26. The First order reactions.
The rate of the reaction will be directly proportional to the
concentration of the reactant.
First order main formulas:
R = kC
K= 1/time
K = τ – time
C0
– initial concentration
Cτ – final concentration
Half – life of a reaction (τ0.5) is the time it takes for the reactant
concentration to decrease to one half of its initial value.
Chemical
Equilibrium

27. Chemical Equilibrium is the state reached by a reaction
mixture when the rates of forward and backward reactions
have become equal.

28. The equilibrium constant Kc
is the value obtained for the equilibrium
constant expression when equilibrium
concentrations are substituted.

29. A reversible reaction is a chemical
reaction that results in an equilibrium
mixture of reactants and products. For a
reaction involving two reactants and two
products this can be expressed
symbolically as: 3H2 + N2 ↔ 2 NH3

30. At Equilibrium the rate of the forward reaction
is equal to the rate of the backward:
R1 = R2

32. Le – Chatelier,
s Principle: when a system in
chemical equilibrium is disturbed by a change of t°,
p or C, the system shifts in equilibrium composition
in a way that tends to counteract this change of
variable ( or: if system at equilibrium is subjected to
a change (such as: t°, p, C), the system will tend to
adjust itself so as to neutralize the effect of change).

33. T = absolute temperature, ln = natural logarithm,
ΔG0
= change of reaction in Gibbs Free energy,
Kc = equilibrium constant.
Thermodynamic equilibrium.
ΔG0
= - lnKc R T

34. Thank you!