1. Chemical Kinetics Part 2
Chapter 13 1 1 1 1

2. The Change of
Concentration with Time

3. Zero-Order Reactions (or zeroth order)
Goal: convert rate law into a convenient equation to give concentrations as a function of time.
For a zero order rxn, the rate is unchanged or is independent of the concentration of a reactant.
However, you must have some of the reactant for the rxn to occur!

4. Zero-Order Reactions (or zeroth order)
One example of a rxn which is 0 order is:
2HI(g) H2(g) + I2(g)
The rate law for this rxn has been determined experimentally and is:
rate = k[HI]0 = k or rate = k
What are the k units?
Rate = M/s so k units are M/s or M•s-1

5. Zero-Order Reactions (or zeroth order)
But the rate is also equal to the change in [reactant] over the change in time:
But if rate = k, this means that this is true:

6. Zero-Order Reactions (or zeroth order)
We rearrange this equation:
We then integrate:

7. Zero-Order Reactions (or zeroth order)
This eq. means that a graph of [HI] vs. time is a straight line with a slope of -k and a y-intercept of [HI]0.
Here are typical 0-order graphs:

8. Zero-Order Reactions (or zeroth order)
We can find the half-life, t1/2, for a 0-order rxn.
The t1/2 is defined as the time it takes for half of the reactant to disappear.
But this is the time required for [A] to reach
0.5[A]0
Mathematically, this is:

9. First-Order Reactions
For a first order rxn, the rate doubles as the concentration of a reactant doubles.
We can show that
A plot of ln[A]t versus t is a straight line with slope -k and y-intercept ln[A]0.
In the above we use the natural logarithm, ln, which is log to the base e.

10. First-Order Reactions

11. Half-Life for 1st-Order Rxns
Half-life is the time taken for the concentration of a reactant to drop to half its original value.
That is, half life, t1/2 is the time taken for [A]0 to reach ½[A]0.
Mathematically,
Note the half-life is independent of the [reactant]0.

12. The Change of Concentration with Time
Half-Life

13. Second-Order Reactions
For a second order reaction with just one reactant
A plot of 1/[A]t versus t is a straight line with slope k and intercept 1/[A]0
For a second order reaction, a plot of ln[A]t vs. t is not linear.

14. Second-Order Reactions

15. Second-Order Reactions
We can show that the half life is:
The half-life of a 2nd-order rxn changes as the rxn progresses.
Each half-life is twice as long as the one before!
This makes these problems harder (and less common).

16. Second-Order Reactions
A reaction can also have a rate constant expression of the form:
rate = k[A][B]
This is second order overall, but has first order dependence on A and B.
This is more complicated and you won’t have to solve for half-lives of these rxns.

17. Temperature and Rate

18. Temperature and 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 chemiluminescence is dependent on temperature: the higher the temperature, the faster the reaction and the brighter the light.
As temperature increases, the rate increases.

19. Temperature and Rate As temperature increases, the rate increases.
Since the rate law has no temperature term in it, the rate constant must depend on temperature.
Consider the first order reaction CH3NC → CH3CN.
As temperature increases from 190°C to 250°C the rate constant increases from 2.52 x 10-5 s-1 to 3.16 x 10-3 s-1.
A rule of thumb is that for every 10°C increase in temperature, the rate doubles!
The temperature effect is quite dramatic. Why?

20. Temperature and Rate

21. The Collision Model
Observations: rates of reactions are affected by concentration and temperature.
Goal: develop a model that explains why rates of reactions increase as concentration and temperature increases.
The collision model: in order for molecules to react they must collide.
The greater the number of collisions the faster the rate.

22. The Collision Model
The more molecules present, the greater the probability of collision and the faster the rate.
The higher the temperature, the more energy available to the molecules and the faster the rate.
However, not all collisions lead to products. In fact, only a small fraction of collisions lead to product.
In order for reaction to occur the reactant molecules must collide in the correct orientation and with enough energy to form products.
These are called effective collisions.

23. The Collision Model

24. The Collision Model

25. Orientation Factor in Effective Collisions
The orientation of a molecule during collisions is critical in whether a rxn takes place.
Consider the reaction between Cl and NOCl:
Cl + NOCl →NO + Cl2
If the Cl collides with the Cl of NOCl then the products are Cl2 and NO.
If the Cl collided with the O of NOCl then no products are formed.

26. Orientation Factor in Effective Collisions

27. Activation Energy
Arrhenius: molecules must possess a minimum amount of energy to react. Why?
In order to form products, bonds must be broken in the reactants.
Breaking bonds always requires energy.
Activation energy, Ea, is the minimum energy required to initiate a chemical reaction.
It is also called the Energy of Activation.

28. Activation Energy

29. Activation Energy
Consider the rearrangement of methyl isonitrile to form acetonitrile:
In H3C-N≡C, the C-N≡C bond bends until the C-N bond breaks and the N≡C portion is perpendicular to the H3C portion. This structure is called the activated complex or transition state.
The energy required for the above twist and break is the activation energy, Ea.
Once the C-N bond is broken, the N≡C portion can continue to rotate forming a C-C≡N bond.

30. Activation Energy

31. Activation Energy
The change in energy for the reaction is the difference in energy between CH3NC and CH3CN.
The activation energy is the difference in energy between reactants, CH3NC and transition state.
The rate depends on Ea.
The higher the Ea, the slower the rate!

32. Activation Energy
Notice that if a forward reaction is exothermic (CH3NC → CH3CN), then the reverse reaction is endothermic (CH3CN →CH3NC).
What is the ΔH and the Ea for the reverse rxn?
Is Ea rev just -Ea?

34. Activation Energy How does the Ea relate to temperature?
At any particular temperature, the molecules (or atoms) have an average kinetic energy.
However, some molecules have less energy while others have more energy than the average value.
This gives us an energy distribution curve where we plot the fraction of molecules with a given energy.
We can graph this for different temperatures as well.

36. Activation Energy
We can see on the graph that some molecules do have enough kinetic energy to react.
This is called f, the fraction of molecules with an energy ≥ Ea.
The equation for f is:

37. Activation Energy
Molecules with an energy ≥ Ea have sufficient energy to react.
What happens to the kinetic energy as we increase the temperature?
It increases!
So, as we increase the temperature, more molecules have an energy ≥ Ea.
So more molecules react per unit time, and the rate increases.

38. Arrhenius Equation
Arrhenius discovered that most rxn-rate data obeyed an equation based on 3 factors:
The number of collisions per unit time.
The fraction of collisions that occur with the correct orientation.
f, the fraction of colliding molecules with an energy ≥ Ea.
From this, he developed the Arrhenius Equation.

39. Arrhenius Equation
In the above, k is the rate constant: it depends on temperature!
R is the Ideal Gas Constant, 8.314J/mol•K
Ea is the Energy of Activation in J
T is the temperature in Kelvin
A is the frequency factor

40. Arrhenius Equation
A is related to the frequency of collisions & the probability that a collision occurs with the correct orientation.
This is related to the molecular size, mass, and shape.
Usually the larger or more complicated the shape, the lower A is.
Important: Both Ea and A are rxn-specific!

41. Arrhenius Equation How do we find Ea and A? By experiments!
You will do this in the lab!
If we have data from 2 different temperatures, we can find Ea mathematically:

42. Arrhenius Equation But we can’t find A with only 2 temperatures.
If we have data from 3 or more different temperatures, we can find Ea and A graphically.
According to the Arrhenius Equation:
If we graph lnk vs. 1/T, we get a straight line with a slope of -Ea/R and a y-intercept of lnA.

43. Arrhenius Equation
Here’s a typical graph of the Arrhenius Equation.