Chapter 14

Chemical Kinetics 


Please watch the videos and
take good notes


Follow the blue Zumdahl l Zumdahl Chemistry book (starts on page 555)

           12.1 Intro Reaction Rates
             Method of Initial Rates

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     12.2 Integrated Rate Law

Click on the link "12.2 Integrated Rate Law" to watch 12.2

    12.3 Catalysts and Mechanisms

Click on the link "12.3 Catalysts and Mechanisms" to watch 12.3


Topic Concepts Experiment

       Reaction Rates

Click on the "Reaction Rates"
The stopped-flow technique and its use in studying chemical reactions is described. The speed or rate of a reaction is discussed. Concentration vs time plots are recorded for reactants and product in a chemical reaction. Characteristic features for rates of change in the concentrations of reactants and product are explored.







Kinetics: the branch of chemistry that deals with rates of reactions-chemistry & engineering.

Usually expressed in terms of reactant:


I. Reaction Rate


rate =
Usually expressed in terms of reactant :



Because it is a reactant a minus sign is used to make rate a positive quantity.
Concentration will be expressed in moles
/liter; time may be in seconds, minutes, years. . .



2N2O5 4 NO2 O2

Rate = -conc./ time = - [N2O5] / t



 Bellow you will find 3 sets of practice questions!!!!

Kinetics Practice Questions

Answer each question on your own paper. The answers have been provided below. Remember, you can only learn from your mistakes! Better to make those mistakes here than on your tests.

Use these two graphs to answer question 1.

1.) Find the reaction rate of the decomposition of methane between 3 and 7 seconds.

1b.) Find the reaction rate of the production of oxygen gas between 2 and 6 seconds.

1c.) Identify the reactant and product based on the graph.

2.) How does temperature affect reaction rate?

3.) What is activation energy?

4.) Explain how a catalyst works relative to activation energy.

5.) With the given information, determine the rate law, the rate constant, and the overall reaction order.

2Mg + O2 Þ 2MgO Rate = k[Mg]n[O2]m
Trial Initial [Mg] (mol/L) Initial [O2] (mol/L) Measured Rate
1 .10 .10 2.0x10^-3
2 .20 .10 4.0x10^-3
3 .10 .20 8.0x10^-3

6a.) Find the half-life of a first-order reaction if the reaction constant, k , is 2.0x10^-3 s-1.

6b.) Find the time when only 1% of reactant remains.

7.) Find  Ea using the following information:

ClO3- + H2O Þ ClO4- + H2

Reaction constant k, (s-1) Temperature(C° )
2.0x10^-3 25
4.0x10^-3 35
8.0x10^-3 45
1.6x10^-2 55

8.) Find the 2nd order reaction’s activation energy with the given information.

K1=4.0L/mol*s at 37C° K2=8.0L/mol*s at 87C°




  • The reaction rate is simply: the change in amount of something / D T.
  • I know the rate law is: D [A] / D T, where [A] is the molarity.

17-35= -18

-18/4= -4.5 L/t


  • The reaction rate is simply: the change in amount of something / D T.
  • I know the rate law is: D [A] / D T, where [A] is the molarity.

36ml-13ml / 6-3

23ml/3s = 7.7 ml/s


  • Oxygen would be the product because its reaction rate graph has a positive slope; the concentration of oxygen is increasing.
  • Methane would be the reactant because its reaction rate graph has a negative slope; the concentration of methane is decreasing.

2.) Temperature increased the reaction rate. Remember for a reaction to take place, the molecules must move faster. An increase in temperature speeds up the molecules.

3.) Activation energy is the minimum amount of energy the reaction must overcome in order for the reaction to take place.

4.) A catalyst lowers the activation energy; therefore, increases the reaction rate.


  • To find the order for Mg, setup the equation like above so the concentration of oxygen cancels out.
  • To find the order for O2, simply try to cancel the concentration of Mg.
  • For the rate constant, once you have the rate law, plug in any row of data from the chart and solve for k. Make sure the concentration corresponds to the measured rate.



Final rate law = k[Mg]1[O2]2

Over all reaction order = n + m= 3

6a). Because it is a first-order reaction, the half-life equation is ln2 / k.

ln2/ 2.0x10^-3 = 350 seconds.

I know the unit is seconds because the unit of my reaction constant is s s-1.

*If the unit is minute-1, the half life would be 350 minutes.

6b) The integrated rate law of a first order-reaction is: ln[A]=-kt+ln[A]0


Another way of looking at this equation is: y=mx+b. Looks familiar? It’s the equation of a line.

y=lnk; m=-Ea/R; x=1/T; b= ln(A).

slope =-Ea/R

Find the slope of the line when graphing lnk vs. 1/T.

(.00325-.00336)/(-5.53+ 6.21)= -6270

Remember R= 8.3145 J and convert T to kelvin.


-6270*-8.3145= 5.21x10^4 J/mol


Make sure you change the temperature to kelvin.

37C° =300k

87C° =360k

R= 8.3145 J

The answer is 1.1x10^4 J/mol.






Answers to REVIEW problems 1-14 are at the end of this document.

Equations to know:

Ln[A[/[A]0 = -kt

T1/2 = ln2/k


1. The reaction of hydrogen and iodine monochloride is represented by this equation:

                                    H2(g) + 2ICl(g) ---> 2HCl(g) + I2(g)


This reaction is first order in H2(g) and also first order in ICl(g). Which of these proposed mechanisms can be consistent with the given information about this reaction?


            Mechanism 1               H2(g) + 2ICl(g) ---> 2HCl(g) + I2(g)


            Mechanism 2               H2(g) + ICl(g) ---> HCl(g) + HI(g)   (slow)

                                                HI(g) + ICl(g) ---> HCl(g) + I2(g)     (fast)


            (A) 1 only                     (B) 2 only                  (C) Both 1 and 2        (D) Neither 1 nor 2



2. Nitramide, N2H2O2, decomposes slowly in aqueous solution. This decomposition is believed to occur according to the reaction mechanism shown below. Which of the rate laws for the decomposition of nitramide is consistent with this mechanism.


            Step 1) N2H2O2 <-> N2HO2- + H+                 (fast equilibrium)

            Step 2) N2HO2- ---> N2O + OH-                   (slow)

            Step 3) H+ + OH- ---> H2O                            (fast)


(A) Rate = k [N2H2O2]

(B) Rate = k [N2H2O2] [H+]

(C) Rate = (k [N2H2O2]) / [H+])

(D) Rate = (k [N2H2O2]) / [N2HO2-]

(E) Rate = k [N2H2O2] [OH-]



3. CH3Cl(g) + H2O(g) ---> CH3OH(g) + HCl(g)

When this reaction was studied, these data were obtained.


                        Initial Concentration, M           Initial Rate, M/s                     

                        (CH3Cl)           (H2O)

Exp. 1              0.100               0.100                           0.182

Exp.2               0.200               0.200                           1.45

Exp.3               0.200               0.400                           5.81


a) Based on these data, what are the orders of the two reactants in this reaction?

b) Write the rate law for the reaction

c) Calculate the value for the rate constant for the reaction



4. 2N2O5(g) ---> 4NO2(g) + O2(g)

What is the ratio of the rate of decomposition of N2O5 to the rate of the formation of NO2?

                        (A) 1:2                        (B) 2:1             (C) 1:4             (D) 4:1



5. The reaction between chloroform, CHCl3(g), and chlorine, Cl2(g), to form carbon tetrachloride CCl4(g) and HCl(g) is believed to occur by this series of steps.


            (Step 1)   Cl2 <---> Cl(g) + Cl(g)

            (Step 2)   CHCl3(g) + Cl(g) ---> CCl3(g) + HCl(g)

            (Step 3)   CCl3(g) + Cl(g) ---> CCl4(g)


a) If this reaction is first order in CHCl3 and half order in Cl2, which statement about the relative rates of steps 1, 2, and 3 is correct?


            (I) Step 1 is the slowest.

            (II) Steps 1 and 2 must both be slow.

            (III) Step 2 must be the slower than step 1.

            (IV) Step 3 must be the slowest


b) Identify the intermediate and the catalyst in the reaction (if they exist).



6. Draw a graph that corresponds to the change in concentration of a reactant that is a first order reaction?



7. How does an increase in temperature affect the rates of the forward and reverse reactions for an

exothermic and an endothermic reaction?



8. The reaction between KMnO4 and H2C2O4 can be followed by monitoring the disappearance of the

purple color of the MnO4- ion. These data were obtained for the reaction carried out at a constant

temperature of 25C.


            [MnO4-]          [H2C2O4]         Rate (M/min)

 0.015              0.020               0.60

 0.015              0.040               1.20

 0.030              0.020               2.40


a) Specify the order of this reaction with respect to [MnO4-] and [H2C2O4]


 b) Write the rate law expression for the reaction and calculate the value for the rate constant.



9.   The reaction 2 N2O5 --> 4 NO2 + O2 is first order with respect to N2O5.

            (i) Draw a graph that represents the change in [N2O5] over time as the reaction proceeds.

            (ii) Describe how the graph in (i) could be used to find the reaction rate at a given time, t.

            (iii) Considering the rate law and the graph in (i), describe how the value of the rate constant, k, could be determined.

            (iv) If more N2O5 were added to the reaction mixture at constant temperature, what would be the effect on the rate constant, k? Explain.




10.       Cl2(aq) + H2S(aq) ---> S(s) + 2 H+(aq) + 2 Cl- (aq)

The rate equation for this reaction is: rate = k[Cl2][H2S] Which of these mechanisms is (are) consistent with this rate equation?


I.          Cl2 + H2S ---> H+ + Cl- + Cl+ + HS- (slow)            

            Cl+ + HS- ---> H+ + Cl- + S                (fast)


II.        H2S  <-> H+ + HS-                 (fast equilibrium)             

            Cl2 + HS- ---> 2Cl- + H+ + S              (slow)


            (A) I only        (B) II only       (C) Both I and II        (D) Neither I or II



11.  2 A(g) + B(g) <===> 2 C(g)


When the concentration of substance B in the reaction above is doubled, all other factors being held constant, it is found that the rate of the reaction remains unchanged. The most probable explanation for this observation is that

(A) the order of the reaction with respect to substance B is 1

(B) substance B is not involved in any of the steps in the mechanism of the reaction

(C) substance B is not involved in the rate-determining step of the mechanism, but is involved in subsequent steps

(D) substance B is probably a catalyst, and as such, its effect on the rate of the reaction does not depend on its concentration

(E) the reactant with the smallest coefficient in the balanced equation generally has little or no effecton the rate of the reaction




12.       H2(g) + I2(g) ---> 2 HI(g)


For the exothermic reaction represented above, carried out at 298 K, the rate law is as follows.

                                 Rate = k [H2] [I2]


Predict the effect of each of the following changes on the initial rate of the reaction and explain your


(a) Addition of hydrogen gas at constant temperature and volume.

(b) Increase in volume of the reaction vessel at constant temperature.

(c) Addition of a catalyst. In your explanation, include a diagram of potential energy versus reaction


(d) Increase in temperature. In your explanation, include a diagram showing the number of

molecules as a function of energy.



13.  2 ClO2(g) + F2(g) ---> 2 ClO2F(g)


The following results were obtained when the reaction represented above was studied at 25C.


Experiment                  [ClO2]                         [F2]                  Initial Rate M/sec

1                                  0.010                           0.010               2.4E-3

2                                  0.010                           0.040               9.6E-3

3                                  0.020                           0.020               9.6E-3


(a) Write the rate law expression for the reaction above.

(b) Calculate the numerical value of the rate constant and specify the units.

(c) In experiment 2, what is the initial rate of decrease of [F2]?

(d) Which of the following reaction mechanisms is consistent with the rate law developed in (a)?

Justify your choice.



       ClO2 + F2 <---> ClO2F2 (fast)

       ClO2F2 ---> ClO2F + F (slow)

       ClO2 + F ---> ClO2F (fast)



       F2 ---> 2 F- (slow)

       2 (ClO2 + F- ---> ClO2F) (fast)



14. What is the activation energy for the reverse of this reaction? N2O4(g) ---> 2NO2(g)  if the enthalpy change for the reverse reaction is DH = - 58.04 kJ/mole and the activation energy for the forward reaction is Ea = +61.24 kJ/mole.

     (A) -3.2 kJ

     (B) +3.2 kJ

     (C) -119.28 kJ

     (D) +119.28 kJ



15. Draw the potential energy profile for the reaction described in question #14, label the reactants, products, DH and Ea.



16. Sketch a graph that shows the pathway of reaction that is endothermic and has a high activation energy.




Miscellaneous Questions


17. Ozone decomposes according to the following balanced equation: 2 O3 -> 3 O2 


The following mechanism was proposed

O3 ->  O2 + O (fast)

O3 + O ->  2 O2 (slow)


Select the rate law rate expression for this mechanism.

(a) rate = k [O3]2 / [O2]

(b) rate = k [O3]2 [O]

(c) rate = k [O3]2

(d) rate = k [O2]3


18. Concentration of reactant [A] versus time that produces a straight line for a first order reaction when plotted is:

(a) 1/[A]

(b) [A]

(c) ln[A]

(d) [A]2



19. Which statement about second order reactions is correct?

(a) Second order reactions require different reactants.

(b) Second order reactions are faster than first order reactions.

(c) Second order reactions are unaffected by changes in temperature.

(d) The half-life of a second order reaction depends on the initial reactant concentration.



20. A first order reaction has a rate constant of 0.0541 s-1 at 25C. What is the half-life for this reaction?

(a) 18.5 s

(b) 12.8 s

(c) 0.0781 s

(d) 0.0375 s


21. The reaction between NO and I2 is second-order in NO and first-order in I2. What change occurs in the rate of the reaction if the concentration of each reactant is tripled?

(a) 3-fold increase

(b) 6-fold increase

(c) 18-fold increase

(d) 27-fold increase



22. For the rate equation, Rate = k[A][B]2, what are the units for the rate constant, k, if the rate is given in mol*L-1 sec-1?

(a) L*mol*sec

(b) L*mol-1*sec-1

(c) L2*mol-2*sec-1

(d) L3*mol-3*sec-2


23. For the reaction: 2A  +  2B -> Product, the rate law is Rate = k[A][B]2. Which mechanism is consistent with this information?


(A)    B + B <->  C

         C + A ->  Product (slow)


(B)    A + B -> C (slow)

         C + B -> product


(C)    A + A <-> C

         B + B -> D

         C + D -> Product (slow)


(D)    A + B <-> C

         B + C -> D (slow)

         D + A -> product



24.  Which straight line gives the activation energy for a reaction?

(a) rate constant vs T

(b) ln (rate constant) vs T

(c) rate constant vs T-1

(d) ln (rate constant) vs T-1




25.  The rate data given were obtained for the reaction, 2NO(g)  +  2H2(g)  ->  N2(g)  +  2H2O(g)

What is the rate law for this reaction?


NO pressure (atm)        H2 pressure (atm)                   Rate (atm*sec-1)

0.375                            0.500                                      6.43 E-4

0.375                            0.250                                      3.15 E-4

0.188                            0.500                                      1.56 E-4


(A) Rate = kPNO

(B) Rate = kP2NO

(C) Rate = k(PNO)(P2H2)

(D) Rate = k(P2NO)(PH2)



26.  What is the rate law for the hypothetical reaction with the

mechanism shown?

                                                         2A     <->   intermediate 1              fast equilibrium

                                     intermediate 1 + B  ->   intermediate 2               slow

                                     intermediate 2 + B  ->   A2B2                             fast


(A) Rate = k[A]2

(B) Rate = [B] 2  

(C) Rate = k[A][B]

(D) Rate = k[A] 2 [B]


27.  According to the Arrhenius equation: k = Ae-Ea/RT, a plot of lnk against 1/T yields

(A) Ea as the slope and A as the intercept

(B) Ea/R as the slope and A as the intercept

(C) Ea/R as the slope and ln A as the intercept

(D) -Ea/R as the slope and ln A as the intercept



Answers to REVIEW problems 1-14:

1. B

2. C

3. a) H2O order = 2, CH3Cl order = 1

c) k = 181.56 M-2-sec-1

4. A

5. Answer is (III)

step (I) is supposed to be an equilibrium, rate of step 2 = k[CHCl3][Cl], from step 1 you get that

[Cl]2 = k[Cl2] so

[Cl] = sqrt k [Cl2]1/2


6. D

7. POORLY WORDED Question, should say what effect does increase in temp have on exo and endo thermic directions of reaction? If assume forward is exothermic rxn, then increase in temp increases the rate forward, for endothermic (reverse rxn), increase in temp increases the rate reverse more than for forward because majority of molecules have Ea in exo direction so change in number of molecules with Ea greater in endo direction. Bottom line increasing temp increases rate in both directions.

8. a) MnO4- order = 2, H2C2O4 order = 1

b) 1.33E5 M-2-sec-1

9. (i) exp decrease

(ii) take dc/dt at any time t

(iii) plot ln[] vs time, calculate slope

(iv) none

10. A

11. C

12. increase, decrease, increase, increase

13. first order in both

b) k = 2.4 M-1-sec-1

c) initial rate divided by 2

d) I

14. B


Chapters by books:

Brown LeMay Bursten
 "Chemistry The Central Science" 
Eighth Edition
Chapter 14
"Chemical Kinetics"
Starts on page 509

FOR CH 14 "Chemical Kinetics"

Zumdahl Zumdahl
"Chemistry" For AP
Sixth Edition
Chapter 12
"Chemical Kinetics"
 Starts on page 555

FOR CH 12 "Chemical Kinetics"


Rate of reaction
Rate of disappearance of reactants
Rate of appearance of products
Nature of reactants
Theoretical rate law
Experimental rate law and determination
Zero, first, second order reactions
Plots of data
Half life
Integrated rate expression

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Chemistry SBHS,
Feb 10, 2010, 4:29 PM