1. REACTION KINETICS
Kenneth E. Schnobrich

2. WHAT IS REACTION KINETICS
It is the study of the factors that can influence the RATE or SPEED of a reaction.
Rate describes how fast a reactant is used or how fast a product is formed in a reaction under a specified set of conditions (temperature, pressure, etc).

3. Factors that Effect Rate
Nature of the Reactants
Concentration
Temperature
Surface Area (Particle Size)
Catalyst
Pressure (Gases only)

4. COLLISION THEORY
Before a reaction can take place there must be COLLISIONS between the reacting species (atoms, molecules, ions)
Successful collisions are those that result in a new product being formed (not all collisions are successful)

5. Successful Collisions
For a collision to be successful it must have:
A sufficient amount of energy when the collision takes place
The reacting particles must be in the proper orientation so that bonds can be broken and new bonds made.

6. Collision Theory H2 and I2 molecules moving randomly
H2I2 intermediate as a result of collision (unstable)
HI product, the result of “successful” collisions – also some unreacted H2 and I2

7. NATURE OF THE REACTANTS
Since bonds must be broken and formed in most cases - the greater the number of bonds broken and formed the slower the rate of reaction.
Reactions involving ions in aqueous solutions are almost instantaneous - the ions are already present so bonds are formed easily.

8. Nature of Reactants (cont.)
In covalent compounds the bonds are more difficult to break and form so the reactions vary in rate.
The more complex the reaction generally the slower the rate of reaction.
CH4 + 2O2 -> CO2 + 2H2O(fast)
2C8H O2 -> 16CO2 + 18H2O(slow)

9. Nature of Reactants + 2 + 2 2 + 25 16 + 18 4 bonds 4 bonds 4 bonds
Double bond
Double bonds + 2 + 2
4 bonds
4 bonds
4 bonds
4 bonds
Broken
Made
Double bond
Double bonds 2 + 25 16 + 18
50 bonds
50 bonds
64 bonds
36 bonds
Broken
Made

10. CONCENTRATION
Since reactions require collisions - the greater the number of reacting particles the greater the chances for successful collisions.

11. CONCENTRATION (cont.)
In conclusion - the greater the concentration the faster the rate of reaction (more collisions -> more successful collisions)
Analogy - students in the hallways
Analogy – cars on the highway

12. TEMPERATURE
Temperature is a measure of average Kinetic Energy - the greater the temperature, the faster the particles are moving (have more energy when they collide) - this assures more successful collisions (bonds can be broken)
Analogy - students in the hallways running to class
Analogy – a posted minimum speed limit of 95 on the highway

13. Temperature
All reactions require a certain minimum activation energy to get them started.
Activated
Complex
Activation
Energy (forward)
Activation
Energy (reverse)
Heat of
Reactants
Potential Energy
Heat of Rx
(DH) = -
Heat of
Products
Reaction Pathway

14. Temperature T1 T1 + 25° At T1+ 25° there are more molecules
with sufficient activation energy
#molecules
T1 + 25°
Activation
Energy
Energy

15. SURFACE AREA
The greater the size of the sample the smaller the surface area. The smaller we make the particle size the greater the surface area.
Mg + 2HCl -> MgCl2 + H2
Increased possibility for collisions
Hydrogen ions hit the outer layer of atoms
…but not those in the
Center of the lump
With the smaller particle size, hydrogen
Ions can reach more Mg atoms

16. Surface Area More Particle interaction 2 inches 2 inches 2 inches
Total Surface (6 sides)
Total Surface (48 sides)
Total Surface (384 sides)
4 in x 4 in = 16 in2 x 6 = 96 in2
2 in x 2 in = 4 in2 x 6 = 24 in2 x 8 = 192 in2
1 in x 1in =1 in2 x 6 = 6 in2 x 64 = 384 in2

17. Catalyst
Catalysts are chemical agents that speed up a reaction without changing either the initial reactants or final products and is not consumed by the reaction.
Heterogeneous catalysts – exist in a different phase than the reaction mixture (like the catalyst in your car’s exhaust system).
Homogeneous catalysts – exist in the same phase as the reaction mixture and frequently are part of an active intermediate to speed up the reaction.
Catalysts lower the activation energy required for a reaction and thus increase reaction rate.

18. Catalysts
Catalysts work in several ways to lower the activation energy –
They can provide a surface for the reaction (putting H2O2 on a cut, the blood cells provide a surface for the decomposition of the H2O2)
They can provide proper orientation for a successful collision (enzymes frequently act this way)
They can stretch bonds and make them easier to break (Au used to weaken bonds between N & O in N2O)
They can form an unstable intermediate (decomposition of H2O2 by NaI(aq))
In each case mentioned above they lower the activation energy needed for the reaction to take place.

19. Catalysts Potential Energy Reaction Pathway Heat of Reactants
Activation Energy with a catalyst
Activation
Energy (forward)
Activation
Energy (reverse)
Heat of
Reactants
Potential Energy
Heat of Rx
(DH) = -
Heat of
Products
Reaction Pathway

20. Equilibrium
In most systems there are two reactions occurring, the forward reaction and the reverse reaction.
When the rate of the forward reaction and the reverse reaction are equal we say the system is in equilibrium (at a specified set of conditions).
At equilibrium the concentrations of both reactants and products stay constant.

21. Equilibrium
To indicate a system has reached equilibrium, double arrows are frequently used or an sign could be used, the arrows are preferred.
There are two types of systems to be studied –
Phase Equilibrium
Chemical Equilibrium

22. Equilibrium
There are also factors that affect a system in equilibrium –
Concentration
Temperature
Pressure/Volume (gases)
The underlying rule is referred to as LeChatlier’s Principle

23. Equilibrium and Rate Rate of forward rx decreases over time
Rate of reaction
Rates are equal (Equilibrium)
Rate of reverse rx increases over time
Time

24. Equilibrium and Concentration
reactants
products
Equilibrium
Equilibrium
Equilibrium
Concentration [ ]
Concentration [ ]
Concentration [ ]
Time
Time
Time
Case 1
Case 2
Case 3
At equilibrium the concentrations of reactants and products are constant and almost the same
At equilibrium the concentrations of reactants are greater than the concentration of the products
At equilibrium the concentrations of reactants are less than the concentration of the products
** In each case, at equilibrium, the concentrations remain constant

25. Phase Equilibrium
Phase equilibrium refers to a case where there is an equilibrium between two different phases of the same substance.
H2O(l) = H2O(g) in a sealed container
The rate of vaporization is equal to the rate of condensation
Polar Water Molecule

26. Phase Equilibrium
In a saturated solution at a specific temperature you have another example of phase equilibrium -
Rate of solution = the rate of crystallization
Chloride ion (Cl-)
Sodium ion (Na+)
Polar Water Molecule

27. Chemical Equilibrium
In a chemical reaction when the rate of the forward reaction is equal to the rate of the reverse reaction at a specified set of conditions the system is in equilibrium,
N2(g) + 3H2(g) = 2NH3(g) N2 H2
Rates are equal (Equilibrium)
Rate of reaction
NH3
Time

28. LeChatlier’s Principle
LeChatlier’s Principle states that a system will always try to reach equilibrium. If a stress, such as concentration change, temperature change, or pressure change (or volume change in systems involving gases) is placed on the system, the system will temporarily favor the forward or the reverse reaction until equilibrium is re-established.
To indicate the reaction that is favored as a result of a stress we will use arrows of unequal length. In this case we are indicating that the forward reaction is favored (or there is a shift to the right)

29. D In Concentration
An increase in concentration of either the reactant(s) or product(s) will favor the reaction that uses the substance added.
In the reaction for the Haber Process, let’s say we increase the concentration of N2(g)
N2(g) + 3H2(g) NH3(g)
Note that the forward reaction is favored (there is a shift to the right)
Eventually, over time, equilibrium will be re-established at a new equilibrium point
Suppose H2(g) is removed (the concentration is decreased)
Note the reverse reaction is now favored to produce more H2(g) (there is a shift to the left)
Eventually, over time, equilibrium will be re-established at a new equilibrium point

30. D In Concentration
Given the reaction, CH4(g) + 2O2(g) = CO2(g) + 2H2O(g) fill in the following table
Stress
Shift
R/L
[CH4]
[O2]
[CO2]
[H2O]
Inc [CH4]
Dec [O2]
Inc [H2O]
Inc [O2]
Dec [CO2]

31. D In Temperature
An increase in temperature will always favor the endothermic reaction. A decrease in temperature favors the exothermic reaction.
2H2(g) + O2(g) = 2H2O(g) kJ …..exothermic reaction
A temperature increase would favor the reverse reaction
2H2(g) + O2(g) H2O(g)
There would be a shift to the left
The [H2] and [O2] would increase and the [H2O] would decrease

32. D In Temperature
Given the reaction, N2(g) + 3H2(g) = 2NH3(g) + heat……exothermic
Stress
Shift
R/L
[N2]
[H2]
[NH3]
Inc T
Dec T

33. D In Pressure/Volume (Gases)
A change in pressure is comparable to a change in volume. If the volume decreases the pressure increases. If the volume increases the pressure decreases.
An increase in pressure always favors the reaction that produces fewer mols of gas in the system.
A decrease in pressure favors the reaction that produces more mols of gas in the system.
After the pressure is changed equilibrium will be re-established at a new equilibrium point.

34. D In Pressure/Volume (Gases)
Given the reaction –
N2(g) + 3H2(g) NH3(g)
4 mols of reactants
1 mol of products
Pressure Increase
Fewer molecules N2 and
H2 , more NH3 , fewer molecules overall
Reaction shifts to the right

35. D In Pressure/Volume (Gases)
In cases where the number of mols of gas are the same for both reactants and products a pressure/volume change has no effect.
Example - H2(g) + I2(g) HI(g)
Complete the following table for the reaction:
N2(g) + 3H2(g) NH3(g)
2 mols
2 mols
Stress
Shift (R/L)
[N2]
[H2]
[NH3]
Inc. P
Inc. V

36. Equilibrium Constants
The equilibrium constant is known by several names but in general it is labeled, Keq. The various versions of equilibrium constant describe the make up of the system at various temperatures.
Given a general reaction (at equilibrium) –
wA + xB = yC + zD (w,x,y,z would all represent numerical coefficients)
By definition Keq = [C]y[D]z/[A]w[B]x - in other words products over reactants.
If the Keq = 1.8 x 10-5 that would tell us that there are very few products in the system at equilibrium. The smaller the value for Keq the fewer the products.

37. Equilibrium Constants
For saturated solutions of salts the equilibrium constant is frequently called the Ksp (solubility product constant). The Ksp gives us an indication of how soluble a salt is in solution at equilibrium.
AgCl has a Ksp of 1.8 x 10-10, that should indicate to you that in the saturated solution of AgCl there are very few Ag+(aq) and Cl-(aq) ions. AgCl(s) Ag+(aq) + Cl-(aq).
For acids and bases they frequently talk about Ka (for acids) and Kb (for bases). The Ka for CH3COOH is 1.85 x That should tell us there are few H+(aq) and CH3COO-(aq) ions in solution (it is a weak acid). CH3COOH(aq) H+(aq) + CH3COO-(aq).