1. CHEMICAL KINATICS

2. Plan 1. Basic concepts 2. Classification of processes
3. The rate of a chemical reaction
4. Influence of a reagent concentration on a reaction rate
5. Influence of temperature on a reaction rate
6. Phenomenon of catalysis

3. The study of reaction rates is called
chemical kinetics
The reaction rate is a development of a reaction in time

4. Chemical kinetics studies rates and mechanisms of chemical reactions
Thermodynamics is a science about macrosystems
The chemical kinetics examines mechanisms of reactions to investigate separate particles
Kinetics and thermodynamics together gives complete representation about chemical reactions

5. Classification of Chemical Reactions
1) Homogeneous reactions occur throughout the volume of reactants
2) Heterogeneous reactions occur on the phase boundary between reactants

6. Classification According to the Mechanism of a Reaction
A reaction mechanism is a description of a path, or a sequence of steps, by which a reaction occurs at the molecular level
A simple reaction has single step
(reagents  products)
A complex (composite) reaction is a sequence of steps
(reagents  intermediate products  final products)

7. Molecularity of Reactions
With respect to the number of the molecules which take part in the reaction at the same time the simple reactions can be divided in:
Monomolecular
N2O4 = 2NO2
Bimolecular
NO + H2O = NO2 + H2
Trimolecular
2NO + Cl2 = 2NOCl

8. 2N2O5 = 4NO2 + O2 1) N2O5 = N2O3 + O2 2) N2O3 + N2O5 = 4NO2
The complex (composite)reactions (several steps in reactions) can be divided with respect to their mechanism process:
consecutive reactions
2N2O5 = 4NO2 + O2
1) N2O5 = N2O3 + O2
2) N2O3 + N2O5 = 4NO2
parallel reactions
3KClO4 + KCl
4KClO3
4KCl + 6O2

9. H2+Cl2 = 2HCl H2 + Cl• = HCl + H• H• + Cl2 = HCl + Cl• Example:
Chain Reactions
Under light quantum the active particles are formed and cause the large number of transformations of an initial molecules
Example:
H2+Cl2 = 2HCl
Cl2 = 2Cl•
H2 + Cl• = HCl + H•
H• + Cl2 = HCl + Cl• h

10. Chain (ramified) Reactions
Н2 + 0,5О2 = Н2О
Origin of a chain : Н2+ О2 = 2ОН
Development of a circuit : ОН + Н2 = Н2О + Н
Branching of a chain: Н+О2 = ОН + О
О + Н2 = ОН + Н
Breakage of chain: ОН + ОН  Н2О2
О + О  О2

11. Slow steps Step 1: A → B (fast) Step 2: B → C (slow)
Often, one step in a mechanism is much slower than any other.
In this case the slow step is rate-determining
Step 1: A → B (fast)
Step 2: B → C (slow)
Step 3: C → D (fast)
A → D

12. Vhomog. = =  Vheter. = Reaction rate n Vt C t n St
This is the number of interaction among molecules occurring in a unit of time in a unit of a volume for homogeneous reactions or on a unit of an interface of phases for heterogeneous reactions:
Vhomog. = = 
Vheter. = n
Vt C t n
St

13. Rate is a function of concentration change
The interactions of atoms and molecules cannot be fixed, therefore the rate of a reaction is expressed as the rate of disappearance of some reactant or appearance of some product.
It is better to express the rate as the changing of concentration of a reagent or a product for the certain time term and also can be used any other parameters such as weight, pressure, volume, coloring, electric conductivity and others

14. Average and Instantaneous RateС2
Average rate
Instantaneous
rate:
C t С1 t1 t2
dC dt C t

15. Vt = - = - = = dCB bdt dCC cdt dCD ddt dCA adt aA + bB = cC + dD
In general for any reaction
aA + bB = cC + dD
the rate of a chemical reaction can be represented by any one of the following expression:
Vt = - = = =
dCB
bdt
dCC
cdt
dCD
ddt
dCA
adt
The rate of a chemical reaction is maximal in a start and minimal to the end of the reaction.

16. The factors influencing to a reaction rate
Nature of a substance
Concentrations of substances
Temperature
Catalysts
Size of particles of solid matter influence on the rate of heterogeneous reactions
Sizes and form of a reactionary vessel influence on the rate of chain reactions

17. Influence of a Nature and Concentration of Reagents on the Reaction Rate
The mass action law
The rate of a simple reaction at a constant temperature is proportional to the multiplication of a concentration of reactants in a degree of stoichiometric coefficients
In general case:
aA + bB + dD
V = kC C  C  x А y B z D

18. V = k C •C Kinetic equation For elementary reaction: аА + bВ = сС +dD
V – reaction rate
k – rate constant
CA и CB – molar concentrations of reactants
а and b – kinetic order of a reaction on substances A and B accordingly a A b B

19. V=kСC2H5OH V = k С2HI V = k C2NOCCl2
Examples of the records of the kinetic equations of the elementary reactions
1) C2H5OH = C2H4 + H2O
V=kСC2H5OH
2) 2HI = H2 + I2
V = k С2HI
3) 2NO + Cl2 = 2NOCl
V = k C2NOCCl2
Overall order of an elementary reaction is equal to its molecularity

20. Graphical determination of n
V = = f(С)
Graphical determination of n dc dt
v а) n= v b) n= v c) n>1
c c c

21. The kinetic equation of a composite reaction
аА + bВ=сС + dD
V = k C C
x & y – The small whole or fractional numbers, which are determined in an experiment (does not coincide with coefficients in the equation) x A y B

22. Example: The reaction 2H2(g)+ O2(g) =2H2O(g) have carried out at one pressure, and then at pressure in 10 times greater.
How has the reaction rate changed, if the kinetic equation of complex chain reaction looks like: V = k [H2]0,4 • [O2]0,3
Solution:
If pressure increases in 10 times, [H2] & [O2] increases in 10 times too, then
V1 = k(10[H2])0,4(10[O2])0,3 = 100,7 = in 5 times

23. Rate Constant a A
Physical meaning of k goes from V = k C  C
If the concentration of reagents CA = CB = 1 mol/L then this is a specific reaction rate
The rate constant when temperature does not change depends only on a nature of substances and does not depend on concentrations
Measuring units of К n=0, [K] = [моль/лс]
n=1, [K] = [1/с]
n=2, [K] = [л/мольc] b B

24. Rate of Heterogeneous Reactions
It depends on specific surface and concentration of reactants in a gas phase or in a solution
V=kSsp(reag.)С(reag.)
Example: CaO(solid)+CO2(g)=CaCO3(solid)
V=kSsp(CaO)С(CO2)
Ssp(CaO) – specific surface of the oxide

25. Specific surface in a reaction changes a little, therefore it is combined with a rate constant of a reaction
Example: to write down the kinetic equation of the heterogeneous reaction:
C(solid) + O2(g) = CO2 (g)
To explain, why coal grinded before combustion on thermal power stations?
Answer:V = kSsp(C)С(O2) or V =ki С(O2)

26. Equilibrium Constant аА+bВ сС+dD
For elementary reversible reaction:
аА+bВ сС+dD
V = Vforward–Vback = k forward C C – k back C C
In the equilibrium state:
Vforward = Vback; kforward[A]a[B]b = kback[C]c[D]d a A b B c C d D

27. The dependence of the rate from temperature
(Van’t-Hoff’s rule )
The rate of simple reaction increases from 2 to 4 as the temperature increases on 10 degrees :
Т  Т0 ,  - temperature
coefficient

28. Arrhenius’s Activation Theory
The chemical reaction can take place only at collision of active particles, i.e. particles must have characteristic energy for the given reaction which is necessary for overcoming of forces of pushing away between electronic shells of particles
The activation energy (Еа, kJ/mol) is an such energy of molecules which is above average statistical supply of molecules energy allowing them to realize chemical interaction

29. When Еа is great, reaction rate is small and vice versa
According to the molecular-kinetic theory of gases for each system there is a threshold of energy (Ea), which is sufficient for the proceeding of a reaction
Ea varies from 0 up to 500 kJ/mol
When Еа is great, reaction rate is small and vice versa
Arrhenius’s equation:
A – frequency factor, it shows some characteristic of a complete collision number
is a share of collisions which give a probability to be a
reaction

30. When temperature increases the share of molecules having energy more than Ea is increased too
It results in the increase of chemical rate of a reaction

31. Power Profile of Exothermal Reaction
А+В
А...В
С...D Еа H
Reaction way Е Е3 Е1 Е2
А…В – promoted complex
Е1, Е2,,,, Е3 - average energy of the molecules of reagents, products and the molecules transitive condition
Еа = Е3 - Е1 – activation energy
Еа` - activation energy of reversible reaction
ЕIа

32. Determination of Activation Energy

33. Catalysis
The catalysis is the phenomenon of acceleration of a reaction under action of substances which not spent in the reaction
The catalytic reactions are reactions, in which the path of the reactions changes at invariable reagents and products
The catalyst is a substance, that increases the rate of a reaction without being consumed by it
Ea of intermediate stages with participation of the catalyst is less, than Ea of the reactions without the catalyst

34. Then at 500 К the reaction rate increases in 1012 times:
2HI = H2+ I2; Еа=184 kJ/mol
Еa=69 kJ/mol in the presence of catalyst (Pt)
Then at 500 К the reaction rate increases in 1012 times:

35. Homogeneous catalysis
(A catalyst and a reagent form one phase)
Example: In technology of H2SO4 production one uses production SO3 by oxidation of SO2
NO2 – catalyst; all substance are gases
1) SO2 + NO2 = SO3 + NO
2) NO + 1/2О2 = NO2
SO2 + 1/2О2 = SO3

36. Heterogeneous catalysis
Production of H2SO4 with Pt catalyst
SO2 (г) + 1/2О2 (г) = SO3 (г)
The efficiency of heterogeneous catalysts is greater than homogeneous
The reaction rate in homogeneous catalysis depends on concentration of the catalyst, and for heterogeneous one depends on its specific surface