1. Chemical Kinetics CHAPTER 14
Chemistry: The Molecular Nature of Matter, 6th edition
By Jesperson, Brady, & Hyslop

2. CHAPTER 14 Chemical Kinetics
Learning Objectives:
Factors Affecting Reaction Rate:
Concentration
State
Surface Area
Temperature
Catalyst
Collision Theory of Reactions and Effective Collisions
Determining Reaction Order and Rate Law from Data
Integrated Rate Laws
Rate Law  Concentration vs Rate
Integrated Rate Law  Concentration vs Time
Units of Rate Constant and Overall Reaction Order
Half Life vs Rate Constant (1st Order)
Arrhenius Equation
Mechanisms and Rate Laws
Catalysts
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

3. CHAPTER 14 Chemical Kinetics
Lecture Road Map:
Factors that affect reaction rates
Measuring rates of reactions
Rate Laws
Collision Theory
Transition State Theory & Activation Energies
Mechanisms
Catalysts
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

4. CHAPTER 14 Chemical Kinetics Factors that Affect Reaction Rates
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

5. Kinetics: Study of factors that govern
The Speed at Which Reactions Occur
Kinetics: Study of factors that govern
How rapidly reactions occur and
How reactants change into products
Rate of Reaction:
Speed with which reaction occurs
How quickly reactants disappear and products form
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

6. A B Kinetics The Speed at Which Reactions Occur
Reaction rate is measured by the amount of product produced or reactants consumed per unit time.
[B] concentration of products will increase over time
[A] concentration of reactants will decrease over time
A B
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

7. 1. Chemical nature of reactants
Kinetics
Factors Affecting Reaction Rates
1. Chemical nature of reactants
What elements, compounds, salts are involved?
What bonds must be formed, broken?
What are fundamental differences in chemical reactivity?
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

8. Factors Affecting Reaction Rates
Kinetics
Factors Affecting Reaction Rates
Ability of reactants to come in contact
(Reactants must meet in order to react)
The gas or solution phase facilitates this
Reactants mix and collide with each other easily
Homogeneous reaction
All reactants in same phase
Occurs rapidly
Heterogeneous reaction
Reactants in different phases
Reactants meet only at interface between phases
Surface area determines reaction rate
Increase area, increase rate; decrease area, decrease rate
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

9. 3. Concentrations of reactants
Kinetics
Factors Affecting Reaction Rates
3. Concentrations of reactants
Rates of both homogeneous and heterogeneous reactions affected by [X ]
Collision rate between A and B increase if we increase [A] or increase [B ].
Often (but not always) reaction rate increases as [X ] increases
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

10. 4. Temperature Kinetics Factors Affecting Reaction Rates
Rates are often very sensitive to temperature
Raising temperature usually makes reaction faster for two reasons:
Faster molecules collide more often and collisions have more energy
Most reactions, even exothermic reactions, require energy to occur
Rule of thumb: Rate doubles if temperature increases by 10 °C (10 K)
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

11. 5. Presence of Catalysts Kinetics Factors Affecting Reaction Rates
Substances that increase rates of chemical reactions without being used up
Rate-accelerating agents
Speed up rate dramatically
Rate enhancements of 106 not uncommon
Chemicals that participate in mechanism but are regenerated at the end
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

12. CHAPTER 14 Chemical Kinetics Measuring Reaction Rates
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

13. Rate = ratio with time unit in denominator Rate of Chemical Reaction
Rates
Measuring Rate of Reaction
Rate = ratio with time unit in denominator
Rate of Chemical Reaction
Change in concentration per unit time.
Always with respect to a given reactant or product
[reactants] decrease with time
[products] increase with time
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

14. Measuring Rate of Reaction
Rates
Measuring Rate of Reaction
Concentration in M units
Time in s units
Units on rate:
[product] increases by 0.50 mol/L per second  rate = M/s
[reactant] decreases by 0.20 mol/L per second  rate = M/s
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

15. Always positive whether something is increasing or decreasing in [X ]
Rates
Rate of Reaction
Always positive whether something is increasing or decreasing in [X ]
Reactants
Reactant consumed
So [X ] is negative
Need minus sign to make rate positive
Products
Produced as reaction goes along
So [X ] is positive
Thus rate already positive
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

16. Measuring Rate of Reaction
Rates
Measuring Rate of Reaction
Coefficients indicate the relative rates at which reactants are consumed and products are formed
Related by coefficients in balanced chemical equation
Know rate with respect to one product or reactant
Can use equation to determine rates with respect to all other products and reactants.
A B  C D
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

17. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)
Rates
Rate of Reaction: Example
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)
O2 reacts 5 times as fast as C3H8
CO2 forms 3 times faster than C3H8 consumed
H2O forms 4/5 as fast as O2 consumed
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

18. NaClO + 2 HCl → Cl2 + H2O + NaCl
Group
Problem
Clorox bleach is sodium hypochlorite. It should never be mixed with acids, (like vinegar) because it forms chlorine gas:
NaClO + 2 HCl → Cl2 + H2O + NaCl
If Chlorine gas (Cl2) is formed at a rate of
5.0 x 10-4 mol/Ls what rate is HCl consumed?
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

19. Change of Reaction Rate with Time
Rates
Change of Reaction Rate with Time
Generally reaction rate changes during reaction, it isn’t constant
Often initially fast when lots of reactant present
Slower and slower as reactants are depleted
Why?
Rate depends on the concentration of the reactants
Reactants being used up, so the concentration of the reactants are decreasing and therefore the rate decreases
Measured in 3 ways:
instantaneous rate, average rate, initial rate
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

20. Instantaneous & Initial Reaction Rate
Rates
Instantaneous & Initial Reaction Rate
Instantaneous rate
Slope of tangent to curve at some specific time
Initial rate
Determined at time = 0
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

21. Rates
Average Rate of Reaction
Average Rate: Slope of line connecting starting and ending coordinates for specified time frame
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

22. Example Reporting Different Types of Rates
Concentration vs. Time Curve for 0.005M phenolphthalein reacting with 0.61 M NaOH at room temperature
Rate at any time t = negative slope
(or tangent line) of curve at that point

23. Rates
Example Reporting Different Types of Rates
[P] (mol/L)
Time (s)
0.005
0.0045
10.5
0.004
22.3
0.0035
35.7
0.003
51.1
0.0025
69.3
0.002
91.6
0.0015
120.4
Initial rate = Average rate between first two data points
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

24. Instantaneous Rate at 51.1 s
Rates
Example Reporting Different Types of Rates
Instantaneous Rate at 51.1 s
(90,0.0028)
(160,0.0018)
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

25. Average Rate between 0 and 120.4 s
Rates
Example Reporting Different Types of Rates
Average Rate between 0 and s
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

26. Group
Problem
A reaction was of NO2 decomposition was studied. The concentration of NO2 was found to be M at 5 minutes and at 10 minutes the concentration was M. What is the average rate of the reaction between 5 min and 10 min?
A M/min
B × 10–3 M/min
C × 10–3 M/min
D. 7.1 × 10–3 M/min
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

27. CHAPTER 14 Chemical Kinetics
Rate Laws
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

28. = k[A]m[B]n Rate Laws Rates Based on All Reactants A + B  C + D
Rate Law or Rate expression
k is the rate constant
Dependent on Temperature & Solvent
m and n = exponents found experimentally
No necessary connection between stoichiometric coefficients (, ) and rate exponents (m, n)
Usually small integers
Sometimes simple fractions (½, ¾) or zero
= k[A]m[B]n
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

29. rate=0.045 M–1 s–1 [0.2][0.3] rate=0.0027 M/s  0.003 M/s Rate Laws
Rates Based on All Reactants
Below is the rate law for the reaction
2A +B → 3C
rate= M–1s–1 [A][B]
If the concentration of A is 0.2 M and that of B is 0.3 M, and the reaction is 1st order (m & n = 1) what will be the reaction rate?
rate=0.045 M–1 s–1 [0.2][0.3]
rate= M/s  M/s
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

30. Exponents specify the order of reaction with respect to each reactant
Rate Laws
Order of Reactions
Rate = k[A]m[B]n
Exponents specify the order of reaction with respect to each reactant
Order of Reaction
m = [A]1 1st order in [A]
m = [A]2 2nd order in [A]
m = [A]3 3rd order in [A]
m = [A]0 0th order in [A]
[A]0 = 1  means A doesn't affect rate
Overall order of reaction = sum of orders (m and n) of each reactant in rate law
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

31. 5Br– + BrO3– + 6H+  3Br2 + 3H2O Rate Laws
Order of Reactions: Example
5Br– + BrO3– + 6H+  3Br H2O
x = 1 y = 1 z = 2
1st order in [BrO3–]
1st order in [Br –]
2nd order in [H+]
Overall order = = 4
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

32. Order of Reaction & Units for k
Rate Laws
Order of Reaction & Units for k
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

33. The following rate law has been observed:
Group
Problem
The following rate law has been observed:
Rate = k[H2SeO][I–]3[H+]2. The rate with
respect to I– and the overall reaction rate is:
A. 6, 2
B. 2, 3
C. 1, 6
D. 3, 6
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

34. If we know rate and concentrations, can use rate law to calculate k
Rate Laws
Calculating k
If we know rate and concentrations, can use rate law to calculate k
From Text Example of decomposition of HI at 508 °C
Rate= 2.5 × 10–4 M/s
[HI] = M
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

35. Determining Exponents in Rate Law
Rate Laws
Determining Exponents in Rate Law
Experimental Determination of Exponents
Method of initial rates
If reaction is sufficiently slow
or have very fast technique
Can measure [A] vs. time at very beginning of reaction
Before it slows very much, then
Set up series of experiments, where initial concentrations vary
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

36. 3A + 2B  products Rate = k[A]m[B]n Expt. # [A]0, M [B]0, M
Rate Laws
Determining Rate Law Exponents: Example
3A B  products
Rate = k[A]m[B]n
Expt. #
[A]0, M
[B]0, M
Initial Rate, M/s 1
0.10
1.2  10–4 2
0.20
4.8  10–4 3
Convenient to set up experiments so
The concentration of one species is doubled or tripled
And the concentration of all other species are held constant
Tells us effect of [varied species] on initial rate
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

37. Determining Rate Law Exponents
Rate Laws
Determining Rate Law Exponents
If reaction is 1st order in [X],
Doubling [X]1  21
Doubles the rate
If reaction is 2nd order in [X],
Doubling [X]2  22
Quadruples the rate
If reaction is 0th order in [X],
Doubling [X]0  20
Rate doesn't change
If reaction is nth order in [X]
Doubling [X]n  2n times the initial rate
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

38. Comparing Expt. 1 and 2 2m = 4 or m = 2 Rate Laws
Determining Rate Law Exponents: Example
Expt. #
[A]0, M
[B]0, M
Initial Rate, M/s 1
0.10
1.2  104 2
0.20
4.8  104 3
Comparing Expt. 1 and 2
Doubling [A]
Quadruples rate
Reaction 2nd order in A = [A]2
2m = 4 or m = 2
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

39. Comparing Expt. 2 and 3 2n = 1 or n = 0 Rate Laws
Determining Rate Law Exponents: Example
Expt. #
[A]0, M
[B]0, M
Initial Rate, M/s 1
0.10
1.2  104 2
0.20
4.8  104 3
Comparing Expt. 2 and 3
Doubling [B]
Rate does not change
Reaction 0th order in B = [B]0 = 1
2n = 1 or n = 0
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

40. Expt. # [A]0, M [B]0, M Initial Rate, M/s
Rate Laws
Determining Rate Law Exponents: Example
Expt. #
[A]0, M
[B]0, M
Initial Rate, M/s 1
0.10
1.2  10–4 2
0.20
4.8  10–4 3
Conclusion: rate = k[A]2
Can use data from any experiment to determine k
Let’s choose first experiment
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

41. Initial Rate of SO3 Formation, M s–1
Rate Laws
Determining Rate Law Exponents: Ex 2
2 SO O2  2 SO3
Rate = k[SO2]m[O2]n
Expt #
[SO2] M
[O2] M
Initial Rate of SO3 Formation, M s–1 1
0.25
0.30
2.5  103 2
0.50
1.0  102 3
0.75
0.60
4.5  102 4
0.90
3.0  102
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

42. 4 = 2m or m = 2 Rate Laws Determining Rate Law Exponents: Ex 2
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

43. 3 = 3n or n = 1 Rate Laws Determining Rate Law Exponents: Ex 2
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

44. Determining Rate Law Exponents: Ex 2
Rate Laws
Determining Rate Law Exponents: Ex 2
Rate = k[SO2]2[O2]1
1st order in [O2]
2nd order in [SO2]
3rd order overall
Can use any experiment to find k
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

45. Expt # [NO] M [O2] M Initial Rate M s–1
Group
Problem
Using the following experimental data, determine the order with respect to NO and O2 .
A. 2, 0
B. 3,1
C. 2, 1
D. 1, 1
Expt #
[NO] M
[O2] M
Initial Rate
M s–1 1
0.12
0.25
1.5  10–3 2
0.24
6.0  10–3 3
0.50
1.2  10–2
Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E