1. Chemical Kinetics

2. Fundamental questions:
Will it take place? Thermodynamics
If it does, how long will it take to reach completion or equilibrium?
Chemical kinetics: is the study of the speeds, or rates, of chemical reactions
A2(g) + B2(g) AB(g)

3. Outline Why study kinetics? Factors affecting rates Measuring rates
reactant concentration
temperature
action of catalysts
surface area.
Concentration vs. Rate (Rate Laws)
Concentration vs. Time (Integrated Rate Laws)
Theories about Rxn Rates
Activation Energy
Mechanisms
Catalysts

4. The rate of the reaction is a measure of how fast the changes are taking place.
Rate of appearance of AB =
D[AB] Dt
Rate of disappearance of A2 =
-D[A2] Dt
Unit:
mol/liter S
= mol / L  s = mol dm-3 s-1

5. Factors affecting reaction rates
The nature of the reactants
2 NO+ O NO2 fast reaction at 25°C
2 CO + O CO2 very slow at 25°C
2. The concentration of reactants Rate equation or rate law
The rate of reaction is proportional to the rate of disappearance of reactants:
-D[A2] Dt
= k [A] mol dm-3 s-1
Rate constant
(in the simplest case: A products)

6. Rate Laws: Basic Assumptions
aA + bB  cC + dD
Simple rate laws depend only on the concentrations of the reactants, not the products.
Rate = k[A]x[B]y
k = rate constant, units depend on the values of x and y.
x and y = the orders of A and B respectively
x + y = overall order of the rxn
3 unknowns (x, y, k)  3 experiments

7. Slope (Rate) Changes w/ Time
Rate is a function of concentration.
When the concentration is high, the rate is large.
Concentration is a function of time.
When the reaction starts the concentration changes the most.
2HI  H2 + I2

8. Factors affecting reaction rates
Zero order reaction:
2 N2O N2 + O2 Au
rate = k
First order reaction:
2 N2O NO2 + O2
rate = k [A]
rate = k [N2O5]

9. Second order reaction: 2 A products
A+B products
rate = k [A]2
rate = k [A] [B]
Third order reaction:
3 A products
rate = k [A] [B] [C]
rate = k [A]3
A + B + C products
rate = k [A]2 [B]
2 A + B products

10. rate = k [A]m [B]n [C]p …. order = m + n + p + …
The order of a reaction is given by the sum of the exponents of the conc. terms in the rate equation.
rate = k [A]m [B]n [C]p ….
order = m + n + p + …

11. Temperature and Rate Constants (k)
Since the rate law has no temperature term in it, the rate constant must depend on temperature.
The size of k depends on T.
Greater T gives a larger k
Consider the reaction H2(g) + I2(g)  2HI(g).

12. 3. The temperature of the reaction. Collision theory of reaction rates
No of molecules
Minimum energy required for reaction Ea
effective collisions
most probable energy at t1 t2

13. n = n0e k = Ae log k = log A - ln k = ln A - -Ea/RT
Maxwell-Boltzmann distribution law
k = Ae
-Ea/RT
Arrhenius equation
ln k = ln A - Ea RT
log k = log A -
2.303 RT
Ea = activation energy
A = frequency factor
K = rate constant

14. Arrhenius Equation K = rate constant A = frequency factor
Ea = activation energy
T = temperature
R = J/mol K

15. Frequency Factor Solve for A T (K) k (L/mol s) 400 7.8 410 10 420 14
430 18
Effective collisions / 6.02 x 1023

16. Temperature Increases Rate
Most reactions speed up as temperature increases. (E.g. food spoils when not refrigerated.)
When two light sticks are placed in water: one at room temperature and one in ice, the one at room temperature is brighter than the one in ice.
The chemical reaction responsible for chemi-luminescence is dependent on temperature: the higher the temperature, the faster the reaction and the brighter the light.

17. Half-Life
Half-lives are typically reported for radioactive materials, medications, toxins…etc.
238U has a t½ = 5 x 109 y
234P has a t½ = 7 h
Amount of time required for half of the compound to react.
Half life equations are found from the integrated rate laws.
Find the time required for the concentration to be one half the initial concentration.

18. The Collision Model Low T Collision High T Collision
Observations: rates of reactions are affected by concentration and temperature.
Concentration terms already accounted for in rate laws
R = k[A]x
Goal: develop a model that explains why rates of reactions increase as temperature increases.
Low T Collision
High T Collision

19. The Rules of Reaction Mechanisms
Elementary step: any process that occurs in a single step.
Molecularity: the number of molecules present in an elementary step.
Unimolecular: one molecule in the elementary step,
Bimolecular: two molecules in the elementary step, and
Termolecular: three molecules in the elementary step.
It is not common to see termolecular processes (statistically improbable).

20. Unimolecular reaction
A products
A unimoleculare reaction is first order
Biomolecular reaction
A + B products
2A products
Termolecular reaction
A + B + C products
2A + B products
3A products
They are not common!

21. Energy of Collision
Chemical rxns usually occur with bonds breaking and bonds forming.
Collision energy supplies the energy needed to break the bonds.
With increased temperature, collision energy increases and the rate of rxn increases.

22. Effective Collisions The right molecules must collide.
The collision must be energetic enough.
The molecules must collide in the proper orientation.
All of the above conditions must be met for a rxn to occur!
There are far more collisions between the wrong molecules or in the wrong orientation or with the wrong energy!
Some estimate that only 1 in 1018 collisions are effective.

23. Molecular Orientation
If H2 and I2 collide, they can collide along any possible trajectory.
However, only one trajectory leads to a rxn.

24. Top of the Peak ‡ H I At the top of the peak, we have…
Maximum potential energy
Transition state
Point where reactants change to products
Bonds breaking and bonds forming
Intermediate structures
The [activated complex]‡ H I
Potential Energy
2 HI
H2 + I2
Progress of Rxn

25. Activation Energy
In order to form products, bonds must be broken in the reactants.
Bond breakage requires energy
Arrhenius: molecules must posses a minimum amount of energy to react..
Ea is the minimum energy required to initiate a chemical reaction.

26. Progress of Reaction Diagrams
Relates the PE of the rxn to time.
For example, an exothermic rxn between A + B is shown on the right.
Progress of Rxn
Potential Energy
Energy of Collision
Activation Energy Ea
Reactants
Energy of Rxn
Enthalpy DH
Note: DH and Ea are unrelated!
100% of Ea is returned!
Products

27. Endothermic System is Reversed
Products are higher in PE than reactants.
Progress of Rxn
Potential Energy
Energy of Collision
Activation Energy Ea
Products
Energy of Rxn
Enthalpy DH
Note: Less than 100% of Ea is returned.
Reactants

28. Reaction Mechanisms 2NO2Cl(g) Cl2(g) + 2NO2(g)
Rate Laws of Multistep Mechanisms
Rate-determining step: is the slowest of the elementary steps.
Therefore, the rate-determining step governs the overall rate law for the reaction.
Mechanisms with an Initial Fast Step
It is possible for an intermediate to be a reactant. …
2NO2Cl(g) Cl2(g) + 2NO2(g)

29. The rate equation for a reaction must be determined by experimentation
Reaction mechanisms:
The rate equation for a reaction must be determined by experimentation E1 E2
Two step mechanism
Rate-determining step
(intermediate)

30. 4. The surface area of a solid
Increasing the surface area of a solid reactant will increase the rate of reaction (explosion of flour dust)
5. The catalysis
A catalyst is a substance that increases the rate of a chemical reaction without being used up in the reaction
reactants
products
Uncatalysed reaction
Catalysed reaction

31. Catalysis A catalyst changes the rate of a chemical reaction.
There are two types of catalyst:
homogeneous, and
heterogeneous.
Chlorine atoms are catalysts for the destruction of ozone.
Homogeneous Catalysis
The catalyst and reaction is in one phase.
Hydrogen peroxide decomposes very slowly:
2H2O2(aq)  2H2O(l) + O2(g).

32. homogeneous
Catalyst heterogeneous
positive (accelerate…)
negative or inhibitor (retard…)
Promoters catalytic poisons

33. Catalysis Homogeneous Catalysis 2H2O2(aq)  2H2O(l) + O2(g).
In the presence of the bromide ion, the decomposition occurs rapidly:
2Br –(aq) + H2O2(aq) + 2H+(aq)  Br2(aq) + 2H2O(l).
Br2(aq) is brown.
Br2(aq) + H2O2(aq)  2Br –(aq) + 2H+(aq) + O2(g).
Br – is a catalyst because it can be recovered at the end of the reaction.
Generally, catalysts operate by lowering the activation energy for a reaction.

34. Catalysis
Homogeneous Catalysis

35. Catalysis
Heterogeneous Catalysis

36. Catalysis Enzymes Enzymes are biological catalysts.
Most enzymes are protein molecules with large molecular masses (10,000 to 106 amu).
Enzymes have very specific shapes.
Most enzymes catalyze very specific reactions.
Substrates undergo reaction at the active site of an enzyme.
A substrate locks into an enzyme and a fast reaction occurs.

37. 1. S + E ES 2. ES E + P Natural catalysts = enzymes (very specific)
Michaelis – Menten mechanism:
S E ES
Substrate
(reactant)
Enzyme
Enzyme-substrate
complex
ES E + P
Product

38. At low concentrate: Rate of disappearence of S = k [S]
Reaction rate
Concentration of substrate
Zero order
First order
At low concentrate: Rate of disappearence of S = k [S]
At high concentrate: Rate of disappearence of S = k’ [S]0 = k’
The enzyme is saturated!

39. Catalysis
Enzymes