Chemistry 12. Unit 1 Review
A: Reaction Kinetics (Introduction)

1. Give examples of reactions proceeding at
different rates.

2. Describe rate in terms of some quantity
(produced or consumed) per unit of time.

3. Experimentally determine rate of a
reaction. Calculate rate of reaction from a
given graph (slope w/ units).

4. Identify properties that could be monitored
in order to determine a reaction rate.

5. Recognize some of the factors that
control reaction rates.

6. Compare and contrast factors affecting
the rates of both homogeneous and
heterogeneous reactions.

7. Discuss situations in which the rate of
reaction must be controlled.

8. Predict if a reaction rate is spontaneous
or not in terms of enthalpy and entropy.
B: Reaction Kinetics (Collision
Theory)

1. Know the following:
- reactions are the result of collisions between
reactant particles.
- all collisions are successful when there is
sufficient kinetic energy (KE) and collide with
favourable geometry.
- to increase the rate of a reaction one must
increase the frequency of successful collision.
Understand the KE vs. Particle distribution graph.
- energy changes are involved in reactions as
bonds are broken and formed (exothermic and
endothermic considerations).


2. Describe the activated complex in terms of
its potential energy (PE), stability, and
structure.
3. Define activation energy.

4. Describe the relationship between
activation energy and rate of reaction.

5. Describe the changes in KE and PE as
reactant molecules approach each other.

6. Draw and label PE diagrams for both
exothermic and endothermic reactions,
including, activation energy, and the energy of
the activated complex. Know how to
determine the activation energy and enthalpy
for both forward and reverse reaction

7. Relate the sign of H to whether the
reaction is exothermic or endothermic

8. Write a chemical equation including the
energy term (given a value) and vice versa.

9. Describe the role of the following factors in
reaction rate (use collision theory):
- nature of reactants
- concentration
- temperature
- surface area
C: Reaction Kinetics (Reaction
Mechanisms and Catalysts)

1. Use examples to demonstrate that most
reactions involve more than one step.
Understand why some reactions occur with
more than one step.

2. Describe a reaction mechanism as the
series of steps (collisions) that result in the
overall reaction. Each step is known as an
elementary step. Provide examples of
several reaction mechanisms.

3. Define catalyst.

4. Compare and contrast the PE diagrams
for a catalyzed and uncatalyzed reaction in
terms of reaction mechanism and activation
energy.

5. Identify reactant, product, reaction
intermediate, and catalyst from a given
reaction mechanism.

6. Describe the uses of specific catalysts in a
variety of situations.
A. Reaction Kinetics
1. Consider the following reaction:
Na2CO3(s) + 2 HCl(aq) ------ > 2NaCl(aq) + H2O(l) + CO2(g)
A 45.0g sample of sodium carbonate is placed in a flask and HCl is added. The reaction consumes
2.15 g of sodium carbonate in 1.0 minute.
a) Calculate the average rate of reaction in moles per second.
b) How many moles of sodium carbonate is left after 6.0 minutes?
2. One mechanism for the destruction of ozone in the upper atmosphere is
O3(g) + NO(g)
--------- > NO2(g) + O2(g)
NO2(g) + _____ -------- > NO(g) + ______ fast
--------------------------------------------------------Overall:
O3(g) + O(g)
slow
------- > 2O2(g)
a) What are the missing substancess?
b) Which specie is an intermediate? Define the term intermediate.
c) Which specie is a catalyst? Define the term catalyst.
d) Make a rough sketch of the potential energy diagram for the reaction. Assume this is an
exothermic reaction.
e) Define “rate determining step”. For the reaction above, which one is the rate determining step?
f) Define reaction mechanism:
3. Air is mainly a mixture of nitrogen and oxygen molecules. These molecules can react to produce
nitrogen dioxide, a red-brown gas.
N2(g) + 2O2(g) ------ > 2NO2(g)
Even though there are more than 4 billion collisions per second between Nitrogen and oxygen, the
amount of nitrogen dioxide present after a year is too small to be detected.
a) Using collision theory, why is this reaction slow?
b) Describe how the rate of reaction can be increased.