Equilibrium Part I Notes

Equilibrium Part I 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) can also be used
reactants yield products
 forward reaction
 reverse reaction
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
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
3. Effective collisions lead to the formation of
products; ineffective collisions do not lead to
the formation of products.
4. During a collision, kinetic energy (mass &
velocity) from the reactants motion is
converted to potential energy within the
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
5. A successful collision results in bond
breaking (endothermic) and bond forming
6. The minimum energy needed to produce an
effective collision is called the activation
energy for the reaction. It’s abbreviation is Ea.
Activation Energy
7. 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.
8. Most reactions occur in a series of steps called a
reaction mechanism.
B Intermediates (found in one
C reaction & used in a later step)
9. Often, one step is slower than the others. This is
called the rate-determining step.
Ex) relay race
10. A substance that increases the rate of a
chemical reaction by providing a mechanism
with a lower energy of activation is called a
11. 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.
12. Energy
Diagram: shows
changes in energy
during a reaction.
Energy Diagram for a Chemical Rxn.
(80 kJ)
(140 kJ)
Energy of
∆H = -60 kJ (for.)
+60 kJ (rev.)
Energy of
Forward Reaction
Reverse Reaction
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
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:
A. Nature of Reactants:
1. Structure:
complexity of
bonds broken and
formed & the
2. State: homogeneous (faster)
heterogeneous (slower)
2 liquids would usually react quickly
A gas and a solid might
react slowly
Reaction Rate
B. Temperature: Rule of thumb – every 10°C
increase doubles the reaction rate.
C. Concentration: increase concentration, increase
# collisions. Aqueous solutions change [conc],
gases change pressure. Solids and pure liquids
(like water) cannot change concentration.
D. Surface Area: Increase surface area, increases
frequency of collisions. This is especially true
for heterogeneous systems (s→l→g).
E. 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:
1. What is in the reaction vessel at time = 0?
H2 + N2
2. Write the forward reaction:
3H2 + N2 → 2NH3
3. What kind of reaction is the forward reaction?
4. Write the reverse reaction:
2NH3 → 3H2 + N2
5. What kind of reaction is the reverse reaction?
6. Over time the concentration of which substance(s)
H2 + N2
7. Over time the concentration of which substance(s)
8. Mark on the graph with a dashed line when
equilibrium is reached.
9. At equilibrium which substance(s) is(are) present in
the greater concentration?
H2 + N2
10. Is the forward or reverse reaction favored?