Unit 14 Chemical Kinetics and Equilibrium
Kinetics (or Reaction Rates): How fast or slow a reaction takes place, reaction speed
Chemical Equilibrium ↔ or
:

State at which [reactant] and [products] have no further tendency to change during time

Forward and reverse reaction proceed at same rate
[brackets denote concentration]
How Chemical Reactions occur
Collision model: In order for molecules to react with each other they must first collide

Collide with sufficient energy

Collide in the correct orientation

Minimum energy = activation energy (Ea)

Energy diagram: activation energy is the energy to “get over the hill” on reaction
coordinate
Conditions that affect Reaction Rate
1. Surface Area – (particle size) bigger particles, more collisions
2. [Concentration] – more molecules, more collisions (see above)
3. Temperature - T↑ molecules move faster, hit each other more frequently and with more
energy
4. Catalyst – increases reaction rate by lowering activation energy (at same concentration and
temperature)
In the space provided draw 3 energy diagrams. Label ΔH on all 3:
1.
Exothermic
2. Endothermic
3. Reaction and catalyzed reaction
Chemical Equilibrium
Reversible reaction – reactants form products (forward reaction) and products form reactants
(reverse reaction)
Equilibrium – forward and reverse reactions occur at same rate, concentrations don’t change
Equilibrium Constant
Law of Chemical Equilibrium – general description of equilibrium condition
aA + bB ↔ cC + dD
Equilibrium Constant (K) – ratio of [products] to [reactants] at equilibrium (at constant temp)
Equilibrium Expression – mathematical expression of equilibrium

Expressed only for gases and aqueous solutions

Solids and liquids are not included because they are pure (elements or compounds), not
mixtures
K=
[𝒑𝒓𝒐𝒅𝒖𝒄𝒕𝒔]
[𝒓𝒆𝒂𝒄𝒕𝒂𝒏𝒕𝒔]
K=
[𝑪]𝒄 [𝑫]𝒅
[𝑨]𝒂 [𝑩]𝒃
Example: Balance equations and write equilibrium constants
2 O3 (g) ↔ 3 O2 (g )
K=
H2 (g) + F2 (g) ↔ 2 HF (g)
K=
[𝑶𝟐 ]𝟑
[𝑶𝟑 ]𝟐
[𝑯𝑭]𝟐
[𝑯𝟐 ][𝑭𝟐 ]
At constant temperature, the system always comes to an equilibrium position regardless of
amounts of reactants and products mixed together initially.
Equilibrium concentrations are not always the same.
Example: Calculate K given the concentrations
Experiment 1:
2 SO2 (g)
1.50 M
Experiment 2:
0.590 M
+
O2 (g)
1.25 M
.045 M
↔
2 SO3 (g)
3.50 M
.260 M
Experiment 1
[𝑺𝑶𝟑 ]𝟐
K =
[𝑺𝑶𝟐 ]𝟐 [𝑶𝟐 ]
=
[𝟑.𝟓𝟎]𝟐
[𝟏.𝟓𝟎]𝟐 [𝟏.𝟐𝟓]
= 4.36
Experiment 2
[𝑺𝑶𝟑 ]𝟐
K =
[𝑺𝑶𝟐 ]𝟐 [𝑶𝟐 ]
=
[𝟎.𝟐𝟔𝟎]𝟐
[𝟎.𝟓𝟗𝟎]𝟐 [𝟎.𝟎𝟒𝟓]
= 4.3
Heterogeneous Equilibria

Homogeneous Equilibria – systems in which the substances are all in the same state (g or aq)

Heterogeneous Equilibria – a system involving more than one state

Liquid water and all solids have a [1] which is why they are not included in K
Examples: Balance the equations and write the equilibrium expression
CaCO3 (s)
CaO (s)
↔
+
CO2 (g)
K = [CO2]
2 H2 (g)
+
K=
O2 (g)
↔
2 H2O (l)
𝟏
[𝑯𝟐 ]𝟐 [𝑶𝟐 ]
LeChâtelier’s Principle
LeChâtelier’s Principle – predicts what happen to systems at equilibrium when condition are changed

When a system at equilibrium is changed, the system will shift its equilibrium position
in order to reduce the change
Most common changes include concentration, volume, pressure and temperature.
1.
Concentration:

Add reactant or product: shift away from added component (“use it up”)

Remove reactant or product: shift toward removed component (“make more”)
2. Volume/Pressure:

V↓ or P↑ shift to decrease pressure (less moles of gas)

V↑ or P↓ shift to increase pressure (more moles of gas)
3. Temperature

Exothermic: produces heat -ΔH (heat is a product)

Endothermic: absorbs heat +ΔH (heat is a reactant)

↑T shift toward side without heat (use heat – cool it down)

↓T shift toward side with heat (produce heat – warm it up)
4. If an inert (noble) gas is inserted into the container, the equilibrium will not shift because it
will not participate in the reaction.
5. If a catalyst is added to the system, the reaction will go faster but the equilibrium will not
shift.
Applications of Equilibrium Constant

Magnitude or size of K tells the tendency of the reaction to occur

If the value of K >> 1 means at equilibrium the system favors forward reaction and will
consist mostly of products

If the value of K is small means at equilibrium the system favors reverse reaction and will
consist mostly of reactants
Solubility Equilibria
Equilibrium exists between a soluble ionic compound dissolving into individual ions and reforming the
solid compound.
The solubility product constant (Ksp) (equilibrium constant for dissolution) common ion effect
Write the dissociation reaction and Ksp for the dissolution shown in the illustration