1
• Energy Changes
• Reaction Rates
• Equilibrium
2
Ch 9.4 Collision Theory and Chemical Reactions
For actual damage to occur, they must collide and
they must do so with sufficient ___________.
3
For a chemical reaction to
occur, reactant particles must
collide with sufficient energy
(activation energy).
Activation energy (Ea) is the
minimum combined kinetic
energy that colliding reactants
must possess in order for their
collision to result in a chemical
reaction.
Higher activation energy →
____________ reaction
Lower activation energy →
____________ reaction
4
Collision Orientation
Reaction rates are sometimes very slow because reactant
molecules must be oriented in a certain way.
Figure 9.6
NO2(g) + CO(g) →
___(g) + ___(g)
5
Ch 9.5 Exothermic Chemical Reactions
When the energy required to break bonds in the reactants
is less than the energy released by bond formation in the
products, the reaction is Exothermic. ∆H = __________
6
Figure 9.7 Energy _______
for an exothermic reaction.
Sometimes an initial input of
energy may be needed but once
it has started, an exothermic
reaction is self-sustaining.
7
Ch 9.5 Endothermic Chemical Reactions
When there is more energy stored in the product molecules than in
reactant molecules, the reaction is Endothermic. ∆H = ________
8
Figure 9.7 Energy diagram
for an endothermic reaction.
A continuous input of ______
is needed for endothermic
reactions to occur.
9
Figure 9.8 Natural processes occur at a wide range of reaction rates.
Ch 9.6 Factors That Influence Chemical Reaction ______
• Physical nature of the reactants
• Reactant concentrations
• Reaction temperature
• Reaction pressure in the case of gases
• Presence of catalysts
10
Particle size of solid reactants greatly affects reaction rates.
Wood fire.
Saw dust explosion.
Coal fire.
Coal mine explosion.
smaller particles = more surface area → ________ reaction rates
11
Reaction rates _________ as
the concentrations of reactants
increase.
Demo: Iodine clock
12
Reaction Temperature greatly affects reaction _________.
Higher temperature
→
→
→
→
Higher average kinetic energy
More frequent collision
More forceful collision
Reactions are faster
Generally, the rate of a chemical reaction doubles for
every 10oC increase in temperature.
Milk spoils faster on the counter top
than in the refrigerator.
Drill question: If milk stays fresh
for 8 days in the refrigerator at 4oC,
in how many days will it spoil on
the counter top at 24oC?
13
Catalysts increase reaction rates without being consumed.
A catalyst provides an alternative route for the reaction
with a _________ activation energy (Ea).
Catalysts DO NOT influence the amount of product
formed, they only speed up the process.
14
Uncatalyzed reaction: X + Y → Z
Steps of the catalyzed reaction:
X + C → XC
Y + XC → XYC
XYC → CZ
CZ → C + Z
15
Solid heterogeneous catalysts
such as in automobile catalytic
converters are plated on
structures designed to ________
their surface area.
Enzymes increase reaction rates in multiple ways, one of which
is to provide a pocket (active site) where reactants (substrates)
can come together with the required orientation. Enzymes are
extremely reactant-specific and catalyze one _________ reaction.
16
Drill Problem: Will each of the changes listed increase or
decrease the rate of the following reaction?
N2 + 3 H2 → 2 NH3
• Adding some N2
• Raising T
• Removing a catalyst
• Removing some H2
Drill Problem:
Reaction A ∆H = -20 kcal/mole Ea = 25 kcal/mole
Reaction B ∆H = -15 kcal/mole Ea = 20 kcal/mole
Which reaction occurs faster at the same temperature?
Which reaction releases more heat energy?
17
Ch 9.7 Chemical Equilibrium
Many chemical
reactions do not go to
completion.
Figure 9.10 Reaction
progress after equal
molar amounts of H2(g)
and I2(g) are mixed.
H2(g) + I2(g) → 2 HI(g)
As the concentration of
the product HI increases,
collisions between HI
molecules increase,
causing the _________
reaction to occur:
2 HI(g) → H2(g) + I2(g)
18
When the forward and reverse chemical reactions occur at the
same rates, a Chemical Equilibrium has been reached.
H2(g) + I2(g)
2 HI(g)
Half-headed arrows denote a chemical system at equilibrium.
At equilibrium, concentrations of
reactants and products are constant.
There is no net change at the
macroscopic level while at the
submicroscopic level opposing
processes continue to occur. This is
called a ____________ Equilibrium.
Figure depicting a dynamic equilibrium: If the liquid flows into the
container at the same rate as it flows out, the level of the liquid in the
container remains constant.
19
Ch 9.8 An Equilibrium Constant is a numerical value that
characterizes the relationship between the concentrations of
reactants and products in a system at chemical equilibrium.
wA + xB
yC + zD
[_______]
Keq =
[C] [D]
y
z
w
x
=
[reactants]
[A] [B]
[ ] = molar concentrations (M = moles/L)
4 NH3(g) + 7 O2(g)
4
[NO2] [H2O]
Keq =
4
[NH3] [O2]
7
4 NO2(g) + 6 H2O(g)
6
20
Pure solids and pure liquids have constant concentrations,
which are incorporated into the equilibrium constant itself.
2 KClO3(s)
2 KCl(s) + 3 O2(g)
Keq = _______
21
Fritz Haber 1918
The Haber-Bosch process is the nitrogen fixation
reaction of N2 and H2 over an enriched Fe or Ru catalyst.
It is one of the most important
industrial reactions and produces
NH3, which is used to synthesize
the huge amounts of fertilizers
that are needed to feed much of the world today.
catalyst
N2(g) + 3 H2(g)
2 NH3(g)
Sample calculation: Keq = _____ at 350oC
[N2] = 0.100 M [H2] = 0.300 M
[NH3] = ?
Keq =
[NH3]2
[NH3]2
√[NH3]2
[NH3]
[NH3]2
[N2][H2]
=
3
[NH3]2
[0.100][0.300]
= _____ x 0.100 x (0.300)3
= 0.189
= √0.189
= 0.435 M
= _____
3
22
Temperature Dependence of Equilibrium Constants
catalyst
N2(g) + 3 H2(g)
2 NH3(g)
Keq = _____ at 350oC
When the reaction temperature changes, Keq also changes!
If forward rxn is exothermic - Keq decreases with _________ T
If forward rxn is endothermic - Keq increases with increasing T
23
Equilibrium Constant Values and Reaction Completeness
wA + xB
yC + zD
[products]
[C] [D]
y
z
[reactants]
[A] [B]
w
x
Keq =
At equilibrium:
When [products] >> [reactants]
Keq = _____
Equilibrium = to the right
large number = large #
small number
When [reactants] >> [products]
Keq = _____
Equilibrium = to the left
small number = small #
large number
24
Table 9.2 Equilibrium Constant Values and Position of Equilibrium
At values of Keq > 1000, [products] >> [reactants] and the reaction is
usually considered _____________.
Unequal arrows are sometimes used to indicate the
direction of the predominant position of the equilibrium.
Caution is advised as this may be misleading. Remember that the rates
for the forward and the reverse reactions at equilibrium are _________.
25
Ch 9.9 Altering Equilibrium Conditions
Le Châtelier’s Principle states that if a stress
(change of conditions) is applied to a system in
equilibrium, the system will readjust (change the
equilibrium position) in the direction that best
reduces the stress imposed on the system.
Henri Le Châtelier
1850-1936
•
•
•
•
Concentration changes (Keq is unchanged)
Pressure changes (Keq is unchanged)
Temperature changes (Keq changes)
Addition of catalyst has NO ____________ on the
equilibrium position, it merely allows the equilibrium
to be reached more quickly. (Keqis unchanged)
26
Concentration changes that result when H2 is added to
an equilibrium mixture involving the following reaction:
Figure 9.12
N2(g) + 3 H2(g)
2 NH3(g)
Keq = [NH3]2 .
[N2][H2]3
• Stress imposed on system: addition of H2 → [H2] too high
• Response: system uses “extra” H2 to make more NH3
• Consequences: some N2 used up = [N2] decreased
more NH3 formed = [NH3] increased
• Conclusion: the stress of “added H2” caused the
equilibrium to “shift to the _________”
27
Drill Problem: Calculate the Keq for the conditions in (a) and (c)
N2(g) + 3 H2(g)
2 NH3(g)
28
Le Châtelier’s principle applies in the same way to removing
reactant or product from the equilibrium mixture as it does to
adding reactant or product at equilibrium.
For example: If we remove some NH3, the equilibrium position
shifts to the right to replenish the NH3.
N2(g) + 3 H2(g)
Stress imposed
2 NH3(g)
Shift observed
Add N2
Remove N2
right
left
Add H2
Remove H2
right
left
Add NH3
Remove NH3
_____
_____
29
Thousands of chemical
equilibria simultaneously
exist in biochemical
systems. Many of them
are ______________.
glucose in blood
stored glucose + H2O
30
Demo: Le Châtelier’s Principle
If SCN− is added to a solution
containing Fe3+, a red solution is
formed due to the formation of
[Fe(SCN)(H2O)5]2+. Lab Exp 18
Fe3+ + SCN-
Applied stress
Add more Fe3+
Add more SCNRemove SCN- (with Ag+)
Remove Fe3+ (with F-)
Remove Fe3+
(by reduction to Fe2+)
Equilibrium shifted
right
right
left
left
left
FeSCN2+
red color
Effect on color
increases
increases
decreases
________
________
31
The equilibrium for the formation of NH3 can also be shifted by
a change in pressure (Keq remains unchanged).
N2(g) + 3 H2(g)
4 moles of gas
2 NH3(g)
2 moles of gas
According to Le Châtelier’s Principle, the stress of increased
pressure is relieved by ____________ the number of moles of
gaseous substances in the system. As there are fewer moles of
product than moles of reactant in the above reaction, the stress
of an increase in pressure can be overcome by formation of
more NH3. Again, the equilibrium shifts to the __________.
32
Pressure changes in the system are brought about through
________ changes. A pressure increase results from a volume
decrease as shown in this example.
17 molecules
initial conditions
17 molecules
stress applied
11 molecules
system adjusted
33
Pressure changes affect systems at equilibrium only when gases
are involved –and then only in cases where the chemical
reaction is such that a change in the total number of moles in the
gaseous state occurs.
Drill Problem:
Rxn A: 2 NO2(g) + 7 H2(g)
Rxn B:
H2(g) + I2(g)
2 NH3(g) + 4 H2O(g)
2 HI(g)
What will be the effects of increasing pressure?
Rxn A:
Rxn B:
What will be the effects of decreasing pressure?
Rxn A:
Rxn B:
34
Temperature-Changes affect equilibrium mixtures.
Effect of temperature
on the equilibrium mixture of
CoCl42- (blue) and Co(H2O)62+
(pink):
Figure 9.13
CoCl42- + 6 H2O
Co(H2O)62+ + 4 Cl- + ____
Le Châtelier’s principle can be used to predict the influence of
temperature-changes on an equilibrium.
35
When the reaction temperature changes, Keq also changes!
H2(g) + F2(g)
2 HF(g) + Heat
For exothermic reactions heat can be treated like a _________.
Increase T, shift left - [HF] decreases, Keq decreases
Decrease T, shift right - [HF] increases, Keq increases
2 CO2(g) + Heat
2 CO(g) + O2(g)
For endothermic reactions heat can be treated like a __________.
Increase T, shift right - [CO] + [O2] increase, Keq increases
Decrease T, shift left - [CO] + [O2] decrease, Keq decreases
36
Supplemental material: Spontaneity and Chemical Reactions
Many processes,
including chemical
reactions, can happen
all by themselves –
they are spontaneous.
Most spontaneous
chemical reactions
are exothermic.
∆H = ________
Demo: Spontaneous oxidation of glycerol by permanganate
3 C3H5(OH)3 + 14 KMnO4
→ 14 MnO2 + 14 KOH + 9 CO2 + 5 H2O + _____
37
Explosions are exothermic
reactions that occur rapidly.
The rapid release of heat energy
causes some of the destruction,
but the entropy change (∆S) for
the reaction is more important
and more devastating.
trinitroglycerin is the
explosive in __________
Example:
2 C3H6(NO2)3(s) →
6 CO(g) + 3 N2(g) + 6 H2O(g)
15 molecules of gas are produced
from 2 molecules of a solid
positive ∆S = very large
Changes in entropy
with changes in state
of matter. Disorder
increases in the
direction shown and
decreases in the
opposite direction.
solid
→ liquid
→
∆S > 0
∆S > 0
gas
When entropy increases, ∆S is __________.
38
Many endothermic reactions are nonspontaneous.
Photosynthesis is nonspontaneous; take a way the light source
and photosynthesis stops.
Endothermic. ∆H = __________
39
Some endothermic reactions are spontaneous because they are
driven by entropy (disorder) = ∆S
Recall Ch 8.3
Instant Cold Pack
NH4NO3 + H2O
Endothermic & Spontaneous
Endothermic: ∆H = ________
Entropy increases: ∆S = positive
40
Conclusion: Two quantities determine whether a reaction is
spontaneous or nonspontaneous: Enthaply (∆Ho) and Entropy
(∆So). The relationship between these is given by:
∆Go = ∆Ho - T∆So
∆Go = the free-energy change for a reaction occurring under a
set of standard conditions: 298 K, 1 atm, and all solutes at 1 M.
When ∆Go = negative, the reaction occurs spontaneously.
41
Example: for an exothermic reaction which shows an
increase in disorder, ∆Ho = (-x) and ∆So = (+y)
∆Go = ∆Ho - T∆So
∆Go = (-x) - T(+y) = negative at all ______
∆Ho
∆So
∆Go = ∆Ho - T∆So
Result
negative positive
negative at all T
spontaneous rxn
positive
negative
positive at all T
nonspontaneous rxn
negative negative
negative if T is
sufficiently low
might be spontaneous
at sufficiently low T
positive
negative if T is
sufficiently high
might be spontaneous
at sufficiently high T
positive
42
Drill problem.
When sodium metal is brought into
contact with chlorine gas, a vigorous
reaction ensues:
2 Na(s) + Cl2(g) → 2 NaCl(s)
Based on this information, is the reaction endothermic or
exothermic?
What are the signs for ∆H _______ (exothermic reaction)
∆S _______ (products are more ordered
∆G
than reactants)
_______ (the reaction is spontaneous)
43
Drill Problem:
PCl5(g)
PCl3(g) + Cl2(g) ∆H = (+)
1. What changes would produce more PCl5(g)?
[ ]?
P?
T?
Catalyst?
2. What is the sign of ∆S in the forward direction?
3. A 4.0 L flask at equilibrium contains
0.60 mole PCl5
0.20 mole PCl3
1.0 mole Cl2
Calculate the Keq for the reaction.