The document discusses reaction rates and factors that affect them. It covers collision theory and the concept of an activated complex. Reaction rates can be determined experimentally and expressed in terms of changes in concentration over time. Factors like temperature, concentration, surface area, and catalysts influence reaction rates. The rate law expresses the relationship between reaction rate and reactant concentrations. Determining reaction order involves measuring initial rates with varying concentrations. Complex reactions occur in a series of steps with an intermediate rate determining step.

1. CHEMICAL
REACTION

2. Reaction Rates
Section 16.1 A Model for Reaction
Rates
Section 16.2 Factors Affecting
Reaction Rates
Section 16.3 Reaction Rate Laws
Section 16.4 Instantaneous
Reaction Rates and
Reaction Mechanisms
Exit

3. Section 16.1 A Model for Reaction Rates
• Calculate average rates of chemical reactions from
experimental data.
energy: the ability to do work or produce heat; it
exists in two basic forms: potential energy and kinetic
energy
• Relate rates of chemical reactions to collisions
between reacting particles.

4. Section 16.1 A Model for Reaction Rates (cont.)
reaction rate
collision theory
activated complex
activation energy
Collision theory is the key to
understanding why some reactions are
faster than others.

5. Expressing Reaction Rates
• The reaction rate of a chemical reaction is
stated as the change in concentration of a
reactant or product per unit of time.

6. Expressing Reaction Rates (cont.)
• Reaction rates are determined
experimentally.

7. Collision Theory
• Collision theory states that atoms, ions,
and molecules must collide in order to
react.

8. Collision Theory (cont.)
• An activated complex is a temporary,
unstable arrangement of atoms in which
old bonds are breaking and new bonds are
forming.

9. Collision Theory (cont.)
• The minimum amount of energy that
reacting particles must have to form the
activated complex and lead to a reaction is
called the activation energy.
• High activation energy means that few
collisions have the required energy and the
reaction rate is slow.

10. Collision Theory (cont.)

11. Collision Theory (cont.)

12. Collision Theory (cont.)

13. Spontaneity and Reaction Rate
• Are more spontaneous reactions faster
than less spontaneous reactions?
• ΔG indicates only the natural tendency for a
reaction to proceed—it does not affect the
rate of a chemical reaction.

14. A. A
B. B
C. C
D. D
A
B
C
D
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Section 16.1 Assessment
Which of the following is NOT a requirement
for a reaction to occur, according to the
collision theory?
A. Reacting substances must collide.
B. Reacting substances must be
in an exothermic reaction.
C. Reacting substances must
collide in the correct orientation.
D. Reacting substances must collide
with sufficient energy to form an
activated complex.

15. A. A
B. B
C. C
D. D
A
B
C
D
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Section 16.1 Assessment
A temporary, unstable arrangement of
atoms in which old bonds are breaking
and new bonds are forming is called ____.
A. reaction complex
B. reaction substrate
C. activated complex
D. activated molecule

16. Section 16.2 Factors Affecting Reaction Rates
• Identify factors that
affect the rates of
chemical reactions.
concentration: a
quantitative measure of
the amount of solute in a
given amount of solvent
or solution
catalyst
inhibitor
heterogeneous catalyst
homogeneous catalyst
• Explain the role of a
catalyst.
Factors such as
reactivity, concentration,
temperature, surface
area, and catalysts affect
the rate of a chemical
reaction.

17. The Nature of Reactants
• Some substances react more readily than
others.

18. Concentration
• Chemists change reaction rates by
changing concentrations of reactants.
• When concentrations are increased, more
molecules are available to collide, and
therefore collisions occur more frequently.

19. Surface Area
• Greater surface area allows particles to
collide with many more particles per unit of
time.
• For the same mass, many small particles
have more surface area than one large
particle.
• Reaction rate increases with increasing
surface area.

20. Temperature
• Increasing temperature generally increases
reaction rate.
• Increasing temperature increases the kinetic
energy of the particles.
• Reacting particles collide more frequently at
higher temperatures.

21. Temperature (cont.)
• High-energy collisions are more frequent at
a higher temperature.
• As temperature increases, reaction rate
increases.

22. Temperature (cont.)

23. Catalysts and Inhibitors
• A catalyst is a substance that increases
the rate of a chemical reaction without
being consumed in the reaction.
• An inhibitor is a substance that slows or
prevents a reaction.

24. Catalysts and Inhibitors (cont.)
• Catalysts lower the activation energy.
• Low activation energy means more collisions
between particles have sufficient energy to
react.

25. Catalysts and Inhibitors (cont.)
• A heterogeneous catalyst exists in a
physical state different than that of the
reaction it catalyzes.
• A homogeneous catalyst exists in the same
physical state as the reaction it catalyzes.

26. A. A
B. B
C. C
D. D
A
B
C
D
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Section 16.2 Assessment
Which of the following generally does not
increase the rate of a chemical reaction?
A. increasing concentration
B. adding a catalyst
C. adding an inhibitor
D. increasing temperature

27. A. A
B. B
C. C
D. D
Section 16.2 Assessment
A
B
C
D
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High-energy particle collisions are more
frequent:
A. when an inhibitor is present
B. when temperature is decreased
C. when activation energy is higher
D. when temperature is increased

29. Section 16.3 Reaction Rate Laws
• Express the relationship
between reaction rate and
concentration.
reactant: the starting
substance in a chemical
reaction
rate law
specific rate constant
reaction order
method of initial rates
• Determine reaction orders
using the method of initial
rates.
The reaction rate law is an experimentally
determined mathematical relationship that relates
the speed of a reaction to the concentrations of
the reactants.

30. Writing Reaction Rate Laws
• A rate law expresses the relationship
between the rate of a chemical reaction
and the concentration of the reactants.
• Rate = k[A] where [A] is the concentration
and k is a constant.

31. Writing Reaction Rate Laws (cont.)
• The symbol k is the specific rate
constant, a numerical value that relates
the reaction rate and the concentrations of
reactants at a given temperature.
• The specific rate constant is unique for every
reaction.

32. Writing Reaction Rate Laws (cont.)
• The reaction order for a reactant defines
how the rate is affected by the
concentration of that reactant.
• Rate = k[H2O2]
• The reaction is first order,
so the rate changes in the
same proportion the
concentration of H2O2
changes.

33. Writing Reaction Rate Laws (cont.)
• The General Rate Law
– Rate = k[A]m[B]n
– Rate = k[NO]2[H2]
– If H2 is doubled, the rate doubles.
– If NO is doubled, the rate quadruples because
22 = 4.
– First-order H2, second-order NO, third-order
overall

34. Determining Reaction Order
• The method of initial rates determines
reaction order by comparing the initial rates
of a reaction carried out with varying
reactant concentrations.
• Initial rate measures how fast the reaction
proceeds at the moment when reactants are
mixed.

35. Determining Reaction Order (cont.)

36. Determining Reaction Order (cont.)
• Doubling [A] doubles the reaction rate, so
[A] is first order.
• Doubling [B] quadruples the reaction rate, so
[B] is second order.
• Rate = k[A][B]2

37. A. A
B. B
C. C
D. D
A
B
C
D
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Section 16.3 Assessment
What is the overall reaction order of the
following reaction?
Rate = k[A]2[B]2
A. 1st
B. 2nd
C. 3rd
D. 4th

38. A. A
B. B
C. C
D. D
Section 16.3 Assessment
A
B
C
D
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In the following reaction, what is the
overall reaction order if doubling [A]
results in quadrupling the reaction rate
and doubling [B] results in a reaction rate
eight times faster?
Rate = k[A]m[B]n
A. 12
B. 5
C. 6
D. 10

40. Section 16.4 Instantaneous Reaction Rates
and Reaction Mechanisms
• Calculate instantaneous rates of chemical reactions.
decomposition reaction: a chemical reaction that
occurs when a single compound breaks down into two
or more elements or new compounds
• Understand that many chemical reactions occur in
steps.
• Relate the instantaneous rate of a complex reaction
to its reaction mechanism.

41. Section 16.4 Instantaneous Reaction Rates
and Reaction Mechanisms (cont.)
instantaneous rate
complex reaction
reaction mechanism
intermediate
rate-determining step
The slowest step in a sequence of
steps determines the rate of the overall
chemical reaction.

42. Instantaneous Reaction Rates
• This figure shows the concentration of
H2O2 over time during the decomposition
reaction 2H2O2(aq) → 2H2O(l) + O2(g).
• The instantaneous
rate is the slope of
the straight line
tangent to the curve
at the specific time.

43. Instantaneous Reaction Rates (cont.)
• Instantaneous rate can be calculated if the
concentrations are known, the temperature
is known, and the experimentally
determined rate law and specific rate
constant at that temperature are known.
• 2N2O5(g) → 4NO2(g) + O2(g)
• Rate = k[N2O5]
• k = 1.0 × 10–5 s–1 and [N2O5] = 0.350 mol/L
• Rate = (1.0 × 10–5 s–1)(0.350 mol/L) =
3.5 × 10–6 mol/(L•s)

44. Reaction Mechanisms
• Most chemical reactions consist of
sequences of two or more simpler
reactions.
• Each step is called an elementary step.
• A complex reaction contains two or more
elementary steps.

45. Reaction Mechanisms (cont.)
• A reaction mechanism is the complete
sequence of elementary steps that makes
up a complex reaction.
• An intermediate is a substance produced in
one of the elementary steps and consumed in
a subsequent elementary step.
• Intermediates do not appear in the net
chemical equation.

46. Reaction Mechanisms (cont.)

47. Reaction Mechanisms (cont.)
• Every complex reaction has one
elementary step that is slower than the
others.
• The slowest elementary step in a complex
reaction is called the rate-determining step.

48. Reaction Mechanisms (cont.)

49. A. A
B. B
C. C
D. D
A
B
C
D
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Section 16.4 Assessment
What is a reaction with two or more
elementary steps called?
A. compound reaction
B. complex reaction
C. multi-step reaction
D. combined reaction

50. A. A
B. B
C. C
D. D
Section 16.4 Assessment
A
B
C
D
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What is the slowest step in a complex
reaction called?
A. elementary step
B. reducing step
C. rate-determining step
D. intermediate step

52. Chemistry Online
Study Guide
Chapter Assessment
Standardized Test Practice
Image Bank
Concepts in Motion

53. Section 16.1 A Model for
Reaction Rates
Key Concepts
• The rate of a chemical reaction is expressed as the
rate at which a reactant is consumed or the rate at
which a product is formed.
• Reaction rates are generally calculated and expressed
in moles per liter per second (mol/(L ● s)).
• In order to react, the particles in a chemical reaction
must collide.
• The rate of a chemical reaction is unrelated to the
spontaneity of the reaction.

54. Section 16.2 Factors Affecting
Reaction Rates
Key Concepts
• Key factors that influence the rate of chemical reactions
include reactivity, concentration, surface area,
temperature, and catalysts.
• Raising the temperature of a reaction generally
increases the rate of the reaction by increasing the
collision frequency and the number of collisions that
form an activated complex.
• Catalysts increase the rates of chemical reactions by
lowering activation energies.

55. Section 16.3 Reaction Rate Laws
Key Concepts
• The mathematical relationship between the rate of a
chemical reaction at a given temperature and the
concentrations of reactants is called the rate law.
rate = k[A]
rate = k[A]m[B]n
• The rate law for a chemical reaction is determined
experimentally using the method of initial rates.

56. Section 16.4 Instantaneous Reaction
Rates and Reaction
Mechanisms
Key Concepts
• The reaction mechanism of a chemical reaction must
be determined experimentally.
• For a complex reaction, the rate-determining step limits
the instantaneous rate of the overall reaction.

57. A. A
B. B
C. C
D. D
A
B
C
D
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The energy required to initiate a reaction
is called ____.
A. initiation energy
B. activation energy
C. complex energy
D. catalyst energy

58. A. A
B. B
C. C
D. D
A
B
C
D
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0%
0%
In general, which of the following does not
cause a reaction rate to increase?
A. increasing surface area
B. increasing temperature
C. increasing volume
D. adding a catalyst

59. A. A
B. B
C. C
D. D
A
B
C
D
0% 0%
0%
0%
What is the overall reaction order of the
following reaction?
Rate = k[A][B]2[C]
A. 1st order
B. 2nd order
C. 3rd order
D. 4th order

60. A. A
B. B
C. C
D. D
A
B
C
D
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0%
0%
A substance produced by an elementary
step in a complex reaction that is
consumed later and does not show up in
the net reaction is called a(n) ____.
A. activated complex
B. catalyst
C. enzyme
D. intermediate

61. A. A
B. B
C. C
D. D
A
B
C
D
0% 0%
0%
0%
Increasing the temperature of a reaction
increases the rate of reaction by:
A. increasing the collision frequency
B. increasing the number of
high-energy collisions
C. both a and b
D. none of the above

62. A. A
B. B
C. C
D. D
A
B
C
D
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0%
0%
Which of the following is an acceptable
unit for expressing a rate?
A. mol/L ● s
B. L/s
C. M
D. mL/h

63. A. A
B. B
C. C
D. D
A
B
C
D
0% 0%
0%
0%
How many moles are in 4.03 × 102 g of
calcium phosphate (Ca3(PO4)2)?
A. 0.721 moles
B. 1.39 moles
C. 1.54 moles
D. 3.18 moles

64. A. A
B. B
C. C
D. D
A
B
C
D
0% 0%
0%
0%
Doubling the concentration of one
reactant in a reaction causes the reaction
rate to double. What is the order of that
reactant?
A. 1st
B. 2nd
C. unable to determine
D. none of the above

65. A. A
B. B
C. C
D. D
A
B
C
D
0% 0%
0%
0%
The rate law for the reaction A + B + C →
Product is rate = k[A]2[B][C]. If [A] = 0.350M,
[B] = .500M, [C] = .125M, and k = 6.50 × 10–5
L3/(mol3 ● s), what is the instantaneous rate
of reaction?
A. 2.84 × 10–6 mol/L ● s
B. 4.98 × 10–7 mol/L ● s
C. 5.84 × 10–6 mol/L ● s
D. 2.84 × 10–7 mol/L ● s

66. A. A
B. B
C. C
D. D
A
B
C
D
0% 0%
0%
0%
H2O2 breaks down to form hydrogen and
oxygen gas in what type of reaction?
A. synthesis
B. double replacement
C. decomposition
D. single replacement

67. Click on an image to enlarge.

81. Figure 16.4 Effect of Molecular Orientation
on Collision Effectiveness

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