Kinetics Web Review Answers

1. Define the following:

Activation energy                    The minimum energy required in a successful collision

Mechanism                              A series of steps that determines the overall equation

Activated Complex                 An unstable reaction intermediate with high potential and low kinetic energy

Successful collision                One that has favourable geometry and sufficient energy

Catalyst                                  Speeds up a chemical reaction by providing an alternate mechanism with lower activation energy-it is not consumed

Reaction rate                           The change in the amount of reactant or product per unit of time

Enthalpy                                  Potential or bond energy

Intermediate                            A species that is produced and then consumed in a reaction mechanism

Homogeneous reaction                        Reactants are in the same phase

Rate determining step              The slowest step or the one with highest activation energy in a mechanism that determines the overall rate

Heterogenous reaction             Reactants have different phases

2. Na2CO3(aq) + 2HCl(aq) → CO2(g) + 2NaCl(aq) +H2O(l)

a) Give four ways to increase the rate of the reaction.

Increase the temperature

Increase[Na2CO3] or [HCl]

Add a catalyst)

b) Give three properties that you could measure in order to determine the rate of the rection. For each one describe how they would change during the experiment.

The volume of CO2(g)                     in a closed System                                                                                         Increase.

The pressure of CO2(g)  in a closed System                                                  Increase.

The mass of an open system where CO2(g) is allowed to escape                 Decreases.

3. Zn(s) + 2HCl(aq) → H2(g) + ZnCl2(aq)

a) Give five specific ways to increase the rate.

Increase Zn(s) surface area

Increase[HCl]

Increase Temperature

Add a catalyst

Agitate.

4. In low light, H2 and Cl2 do not react at all. When exposed to UV light, they react explosively! Explain using the collision theory. The activation energy is too high for successful collisions.

UV light provides the activation energy to begin the reaction and because it is exothermic, it continues and explodes.

5. A mixture of KClO3 and C12H12O11 do not react at all at room temperature. A drop of H2SO4 starts the reaction, slowly at first, then it quickens into a flaming inferno. Explain using the collision theory.

H2SO4 is a catalyst, which lowers the activation energy, allowing lower energy collisions to be successful.

6. Water puts out a fire. Explain using the collision theory.

Water lowers the temperature of the reaction so that collisions no longer have the required activation energy.

7. A glowing splint reignits in pure O2. Explain using the collision theory.

Pure O2 has a higher concentration and therefore more collision, which increases the reaction rate.

8. Enzymes in the human body allow the oxidation of carbohydrates at 37 degrees celcius.

Enzymes are catalysts, which lower the activation energy allowing low temperature collisions to be successful.

9. Draw an exothermic PE diagram. Include a catalyst. Label the change in enthalpy, the forward and reverse activation energies and the activated complex.

10. Repeat the above for an endothermic reaction.

11.       A reaction generally starts fast because reactant concentrations are highest at the start of a reaction. As the reaction proceeds the reactants concentrations decrease and the rate will get slower and slower as time goes on.

12.       A + WY                   AWY                          FAST

AWY + HA            A2WY + H                   SLOW

A2WY + HA            A3 + WY + H FAST

I. List the following:

a) the overall equation                        2HA + A                 2H + A3

b) a catalyst                             WY

c) Intermediates                       AWY and A2WY

d) Reactants                            2HA + A

e) Products                              2H + A3

II. Describe how each change affects the rate.

a) increasing the concentration of A                            No change, A is not in rate the determining step or the slowest

b) increasing the concentration of H                            No change, H is not a reactant

c) increasing the concentration of HA                         Increases, HA is a reactant in the rate determining step

d) removing WY completely                                       Decreases, WY is a catalyst as it crosses out from left to right

e) Decreasing the temperature                                     Decreases

13. Describe the KE and PE changes as two molecules:

a) approach to collide,                                                            KE decreases and PE increases

b) form an activated complex, and                              KE is a minimum and PE is a maximum

c) form products in an exothermic reaction.                KE increases to a maximum and PE decreases to a minimum

14. Draw the PE diagram for a mechanism with three steps.

How many activated complexes are there?     Three

How many intermediates are there?                Two

15. Describe as endo or exothermic.

a) 2H2 + O2              2H2O + 300 KJ                                               Exothermic

b) NH4NO3(s)          NH4NO3(aq)       ΔH = +150 KJ                                     Endothermic

16. Change each equation in 15. from standard to ΔH notation or vice-versa.

2H2 + O2                2H2O                          ΔH = -300 KJ

NH4NO3(s)       +          150 kJ                     NH4NO3(aq)

17. Which reaction at room temperature is faster and why?

a) Pb+2(aq)       +            2Cl-1(aq)                    PbCl2               Fast simple ionic or double replacement- there are no bonds to break and low Ea

b) 2H2(g)           + O2(g)                      2H2O                          Slow as there are gaseous reactants are slower than aqueous and higher Ea

18. List three commercial catalysts.    Pt        Os       H2SO4             Ni

19. Calculate the rate in mols/s.

 Moles H2 10 15 21 24 Time (seconds) 200 300 400 500

Rate    =          (24.0  -  10.0) moles               =          0.0467 mol/ s

(500.0 – 200.0) s

20. Indicate how each change will effect the rate of the reaction and the PE diagram and explain with the collision theory.

a) increasing the temperature

Increase the rate as the average collision will have more Ea and there will be more collisions. PE diagram will not change.

b) increasing the concentration of a reactant

Increase the rate as there will be more collisions. PE diagram will not change.

c) increasing the concentration of a product

The rate will not increase. PE diagram will not change.

b) addition of a catalyst

Increase the rate as the the Ea will be lower and more low energy collisions will be successful. The activated complex is lower.

21. X + Y                XY                  slow

XY + Z                   XYZ                fast

XYZ + W               XYW + Z        fast

a) Identify intermediates

XY and XYZ               Cross out from right to left

the catalyst

Z                                  Crosses out from left to right

and the rate determining step

1 which is the slow step

b) If the concentration of X was increased will the rate increase?

Yes

22. You can only increase the surface area of a substance if it is in certain physical state. What is it?

Solid

23. Two solid reactants react. Is it a homogeneous or heterogeneous reaction?

Heterogeneous

24. What is the mathematical relationship between reaction rate and activation energy

Inverse

reaction rate and reactant concentration

Direct

reaction rate and temperature?

Direct

25. Draw a collision energy distribution diagram for a reaction where the y axis is fraction of collisions and the x axis is collision energy. Draw the Ea line showing about 10% of the collisions having sufficient energy. Draw the Ea line for the catalyzed reaction where 20% have sufficient energy.

26.

 Fraction of collisions

 Collision Energy