Equilibrium Notes
The Concept of Equilibrium
• Ex) elevator, football game, moving walkway
1. Reversible Reactions: Chemical reaction in which the
products can regenerate the original reactants. A double
arrow is used to express a reversible reaction.
Ex)
2NO2(g) ↔ N2O4(g)
reactants yield products
↔
or

 forward reaction




or

 reverse reaction
Forward Rxn: 2NO2 → N2O4 Reverse Rxn: N2O4 → 2NO2
2. Some reactions can reverse on their own, some under certain
conditions, (temp, pressure or a catalyst) and some
will not reverse.
Ex) Single Replacement Rxns:
3CuCl2(aq) + 2Al(s) → 3Cu(s) + 2AlCl3(aq)
3. Chemical Equilibrium:
State of a product’s and
reactant’s concentrations
remaining constant, not
equal. This is because
the rate of the forward
and reverse reactions are
equal. The symbol “[ ]”
denotes concentration.
A. Reaction rates are affected by concentration.
If the concentration of the reactant increases,
the rate of the forward reaction will increase.
B. As the reaction proceeds, the concentration
of the reactants decrease and the products
increases. This will cause the forward
reaction to decrease.
4. Reaching equilibrium does not mean the
reaction has stopped. It only means the rate of
the forward and reverse reactions are equal.
This is not a static equilibrium, but a dynamic
equilibrium that stays constant over time.
Concentration becomes
Constant NOT Equal
SPEED
AMOUNT
A↔B
Rates become Equal
Reaction Rates & Equilibrium:
A. Collision Theory: (Theoretical Model)
1. In order for reactions to occur between
substances, the particles must collide.
2. According to the Collision Theory, a
successful collision occurs when
A. the collision is energetic enough, and
B. the particles collide with the correct
orientation.
3. Effective collisions lead to the formation of
products; ineffective collisions do not lead to
the formation of products.
Ineffective Collision – Insufficient Energy - No Products
Effective and Ineffective Collisions
Effective Collision – Sufficient Energy – Forms Products
Effective collisions have enough energy, and
the correct orientation to form products.
Ineffective collisions revert to the original
reactants.
4. A successful collision results in bond
breaking (endothermic) and bond forming
(exothermic).
5. The minimum energy needed to produce an
effective collision is called the activation
energy for the reaction. It’s abbreviation is Ea.
Activation Energy
6. A transitional structure results from a successful collision.
The structure is present while old bonds are breaking and
new bonds are forming. It is called an activated complex
and is unstable and short-lived. It is neither reactant nor
product.
7. A substance that increases the rate of a chemical reaction
by providing a mechanism with a lower energy of
activation is called a catalyst.
8. A system that has just one phase is called a homogeneous
system, (g→g→g) while a system that has more than one
phase is called a heterogeneous (s→l→g) system.
Endothermic
9. Energy
Diagram: shows
changes in energy
during a reaction.
Exothermic
Energy Diagram for a Chemical Rxn.
Reaction
Pathway
Activated
Complex
(kJ)
Ea
forward
reaction
Ea
Added
Catalyst
(80 kJ)
reverse
reaction
(140 kJ)
Energy of
Reactants
Potential
∆H = -60 kJ (for.)
+60 kJ (rev.)
Energy of
Products
Forward Reaction
(Exo)
Reverse Reaction
(Endo)
Factors Affecting Reaction Rates
1. Chemical Kinetics is concerned with the rate at
which a reaction occurs.
2. Reaction rate is a measure of the rate or speed of
a chemical reaction. Reaction rate is determined
by measuring the change in concentration of
reactants and products over a certain amount of
time. These reaction rates are determined
experimentally.
3. Rate-influencing factors are those factors that
affect rate of reactions by altering the frequency,
orientation, or energy at which particles collide.
4. According to the Collision Theory, the following
5 factors affect the rate of a reaction:
1. Nature of Reactants:
A. Structure:
complexity of
bonds broken and
formed & the
orientation.
B. State: Homogeneous systems have reactants and
products in the same state. Heterogeneous systems
have reactants and products in different states.
Homogeneous systems usually react faster and
heterogeneous systems react slower.
2 liquids would usually react quickly
Iron rusts in air very slowly
Reaction Rate
2. Temperature: (Average kinetic energy!) Rule
of thumb is that every 10°C increase in
temperature doubles the reaction rate.
3. Concentration: increased concentration, increases
# collisions. Aqueous solutions can change [conc]
or molarity, gases can change pressure. Solids
and pure liquids like water cannot change
concentration.
4. Surface Area: Increased surface area, increases
frequency of collisions. This is especially true
for heterogeneous systems (s→l→g).
5. Catalyst: Increases the rate of reaction without being
consumed in the reaction. Catalysts speed up
the reaction rate by lowering the activation
energy needed for the reaction to occur.
Uncatalyzed reaction (slow)
Catalyzed reaction (fast)
5. Inhibitors: Decrease the rate of the reaction by
taking the place of a reactant and stopping the
reaction (opposite of a catalyst.)
Energy diagram involving a catalyst:
11. What is in the reaction vessel at time = 0?
H2 + N2
12. Write the forward reaction:
3H2 + N2 → 2NH3
13. What kind of reaction is the forward reaction?
synthesis
it slows down
14. Write the reverse reaction:
2NH3 → 3H2 + N2
15. What kind of reaction is the reverse reaction?
decomposition
it speeds up
16. Over time the concentration of which substance(s)
decreases?
H2 + N2
17. Over time the concentration of which substance(s)
increases?
NH3
18. Mark on the graph with a dashed line when
equilibrium is reached.
19. At equilibrium which substance(s) is(are) present
in the greater concentration?
H2 + N2
20. Is the forward or reverse reaction favored?
reverse
(min)
Equilibrium Systems and Stress – Le Chatlier’s Principle
1. Le Chatlier’s Principle states that if a stress is imposed
on a system at equilibrium, the equilibrium position will
shift toward the direction that tends to minimize the stress.
(Pure liquids and solids are not affected by changes in
equilibrium.)
2. In chemistry, Le Chatlier’s Principle is used to
manipulate the outcome of reversible reactions to maximize
the amount of product produced by altering the
temperature, pressure, or removing a reactant or product
from a reaction.
Crash Course in Chemistry – Equilibrium Video
Khan Academy – LeChatlier’s Principle
The Haber Process
3. An application of Le
Chatlier’s Principle is The
Haber Process. This was used
by Germany in WWI to
produce ammonia (NH3) for
nitrogen containing explosives.
Today this process is used to
create ammonia (NH3) for
household cleaners and
especially fertilizers allowing
for a four-fold increase in food
production from 1900 to 2000.
4. There are 4 factors which determine
whether a reaction favors making reactants
“lies to the left”
or favors making
products “lies to the right”.
A. Changes in Concentration
-Adding/Removing Reactants
-Adding/Removing Products
B. Changes in Pressure
C. Changes in Temperature
Changes in Concentration - (Reactants and Products)
1. Adding a substance to a system at equilibrium drives the system
AWAY from that substance and makes more of the substance(s)
on the opposite side.
2. Removing a substance from a system at equilibrium drives the
system toward REPLACING that substance.
Ex) N2O4(g) ↔ 2NO2(g)
↑ N2O4(g)
↑ NO2(g)
↓ N2O4(g)
↓ NO2(g)
shift right making more NO2
shift left making more N2O4
shift left making more N2O4
shift right making more NO2
B. Changes in Pressure
1. Avogadro’s Law – Equal volumes of gases at the same
temperature and pressure have the same number of
molecules. (1 mole of any gas at STP = 22.4L = 6.02 x1023
particles)
2. According to Boyle’s Law an increase in pressure means
a decrease in volume. So if the pressure is increased on a
system at equilibrium, the side which occupies the lower
volume will be favored. (inverse relationship)
3. If pressure is increased, the reaction will shift in the
direction that produces fewer moles.
Ex)
2NO2(g) ↔
N2O4(g)
↑ pressure shift to right making more N2O4
2H2O2(g) ↔ 2H2O(g) + O2(g)
↑ pressure shift to left making more H2O2
H2(g) + Cl2(g) ↔ 2HCl(g)
↑ pressure shift to neither! making more same # of
moles each side
4.Pressure changes will only affect gases!.
C. Changes in Temperature
1. Exothermic Reactions release heat. If temperature is increased
on this system then the reaction which absorbs or uses heat will
increase. If temperature is decreased on an exothermic reaction,
then the reaction which releases heat will decrease.
Ex) H2(g) + I2(g) ↔ 2HI(g) + Heat
Increasing the temperature will produce a higher yield of H2 + I2.
Lowering the temperature will produce a higher yield of HI.
C. Changes in Temperature
2. Endothermic Reactions absorb heat. If temperature is increased
on this system then the reaction which absorbs or uses heat will
increase. If temperature is decreased on an endothermic reaction,
then the reaction which releases heat will decrease.
Ex) HEAT + NH4Cl(s)
↔ NH3(g) + HCl(g)
Increasing the temperature will produce a higher yield of NH3 + HCl.
Lowering the temperature will produce a higher yield of NH4Cl.
Conclusion: These examples have been
illustrations of Le Chatlier’s Principle
which states that a system at equilibrium,
when subjected to a stress, will temporarily
adjust itself to relieve the stress. This means
that the shift to the right or left, or the
increased forward or reverse reaction, will
be temporary and a new equilibrium will be
reestablished.
Practice:
1. 2SO2(g) +
O2(g)
 2SO3(g) + heat
What conditions of temperature and pressure favor high
equilibrium concentrations of SO3?
(high / low) pressure; (high / low) temperature
1.
3H2(g) + N2(g)

2NH3(g) + heat
The commercial production of ammonia uses the Haber
Process which is expressed by the above equation.
What condition of temperature and pressure will
provide a maximum yield of NH3?
(high / low) pressure: (high / low) temperature
1. 4HCl(g) + O2(g)  2H2O(g) + 2Cl2(g) + heat
Increasing the temperature of the reaction will
(increase / decrease) the forward reaction.
Decreasing the pressure on the system will
(increase / decrease) the forward reaction.