Kinetics , Thermodynamics
and Equilibrium
Regents Chemistry
Kinetics and Thermodynamics
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Kinetics: deals with rates of reactions
(how quickly a reaction occurs)
Thermodynamics: involves changes in
energy that occur in reactions
Kinetics: Collision Theory
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

Measured in:
#moles of reactant used/unit time
Or
# moles of product formed/unit time
Frequency of collisions: more collisions = faster
rate
Effective collisions: must have 1) proper
orientation and 2) enough energy
Factors Affecting Rate
1. Type of substance:
 Ionic substances react faster: bonds require less energy
to break
AgNO3 (aq)+NaCl(aq)AgCl(s)+NaNO3 (aq)
In solution ionic solids dissociate into ions:
Ag+ NO3Na+
Cl
Covalent react more slowly: bonds require more energy
to break
H2 (g)+I2 (g)2 HI
(g)
Bonds must be broken then be reformed. (takes more
time)
Factors Affecting Rate
2. Temperature increase
 Average kinetic energy increases and the
number of collisions increases. Reactants
have more energy when colliding. This
increases rate.
Factors Affecting Rate
3. Concentration increase
 Increases rate due to the fact that more
particles are in a given volume, which
creates more collisions.
Factors Affecting Rate
4. Surface Area Increase
 Increases rate due to increased reactant
interaction or collisions (powder vs. lump)
Factors Affecting Rate
5. Pressure Increases
 Increases the rate of reactions involving
gases only
As pressure  Volume  so:
spaces between molecules 
 frequency of effective collisions
Factors Affecting Rate
6. Catalyst: substance that increases rate of
reaction, provides a shorter or alternate
pathway by lowering the activation energy
of the reaction.
 Catalysts remain unchanged during the
reaction and can be reused.

Activation energy: amount of energy
required to “start” a reaction
Quick Review –
Factors that affect reactions
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Ionic solutions have faster reactions than
molecule compounds. (bonding)
Temp.  Rate
 conc. rate
 surface area  rate
 Pressure  rate,  P  rate
Catalysts speed up reactions.
Potential Energy Diagrams

Graphs heat during the course of a
reaction.
Exothermic: PE of products is less because energy was lost.
PE of reactants (ER)
PE of Activated Complex
Heat of reaction
(ΔH) = Ep - ER
Activation Energy (Ea)
PE of products (EP)
Activation Energy (Ea)*
reverse reaction
Endothermic: PE of products is more because energy was gained.
PE of reactants (ER)
Activation Energy (Ea)
PE of Activated Complex
PE of products (EP)
Heat of reaction (ΔH)
Activation Energy (Ea)*
reverse reaction
Catalysts
Thermodynamics


Heat content (Enthalpy): amount of heat
absorbed or released in a chemical
reaction
Enthalpy (ΔH = Hproducts – Hreactants)
ΔH = Hproducts – Hreactants
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ΔH is positive when the reaction is
endothermic. Heat of products are greater
than reactants
ΔH is negative when the reaction is
exothermic. Heat of reactants were
greater than the products
Table I
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Includes heats of reaction for combustion,
synthesis (formation) and solution
reactions.
You must remember equation
stoichiometry (balanced equations).
Endothermic: heat is a reactant
Exothermic: heat is a product
Table I- Practice
1.
2.
3.
Which reaction gives off the most
energy?
Which reaction gives off the least
energy?
Which reaction requires the most energy
to occur?
Entropy (ΔS)
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Definition: randomness, disorder in a
sample of matter
Gases have high entropy
Solids have low entropy
Increasing ΔS
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Phase change from s  l  g
Mixing gases
Dissolving a substance
Spontaneous Reactions
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Nature favors low energy (more stable)
and high entropy
Reactions are spontaneous when heat
(ΔH) decreases and entropy (ΔS)
increases
ΔH = (-)
ΔS= (+)
Analogy: Your Bedroom
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You like to have low enthalpy
(low energy) when it comes to
household chores.
As a result, your room tends to
have high entropy (very messy,
disorderly).
This is what nature prefers: low
enthalpy and high entropy.
Stability of Products and H
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Help determine if a reaction is spontaneous
Products tend toward Lower energy (-ΔH)
Products tend toward more randomness (+ΔS)
Products of exothermic reactions are usually
more stable. Result in lower amounts of heat.
The more negative the H, the more stable the
product is.
Gas products result in increased Entropy.
Chemical Equilibrium
Regents Chemistry
Reversible Reactions

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Most chemical reactions are able to
proceed in both directions under the
appropriate conditions.
Example:
Fe3O4 (s) + 4 H2 (g) ↔ 3 Fe(s) + 4 H2O(g)
Reversible Reactions cont.
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In a closed system, as products are
produced they will react in the reverse
reaction until the rates of the forward and
reverse reactions are equal.
Ratefwd = Raterev
This is called chemical equilibrium.
Equilibrium

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Equilibrium is dynamic condition where
rates of opposing processes are equal.
Types of Equilibrium:

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Phase equilibrium
Solution Equilibrium
Chemical Equilibrium
Phase Equilibrium
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Rate of one phase change is equal to the
rate of the opposing phase change.
Occurs when two phases exist at the same
temperature.
Example: Ratemelting = Ratefreezing
H2O (s)  H2O (l)
Solution Equilibrium

Rate of dissolving = rate of crystallization

Occurs in saturated solutions
Chemical Equilibrium
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Rateforward reaction = Ratereverse reaction
Concentration of reactants and products
are constant NOT necessarily equal.
[reactants] and [products] is constant.
The Concept of Equilibrium
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As a system approaches equilibrium, both the
forward and reverse reactions are occurring.
At equilibrium, the forward and reverse reactions
are proceeding at the same rate.
Le Chatelier’s Principle
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Whenever stress is applied to a reaction at
equilibrium, the reaction will shift its point
of equilibrium to offset the stress.
Stresses include:

Temperature, pressure, changes in reactant
or product concentrations
Example: The Haber Process
N2 (g) + 3 H2 (g)  2 NH3 (g) + heat
a)
b)
c)
d)
e)
f)
g)
h)








[N2]
[H2]
[NH3]
[NH3]
pressure
pressure
temperature
temperature
Example: The Haber Process
N2 (g) + 3 H2 (g)  2 NH3 (g) + heat
a)
b)
c)
d)
e)
f)
g)
h)








[N2]
[H2]
[NH3]
[NH3]
pressure
pressure
temperature
temperature
shift
shift
shift
shift
shift
shift
shift
shift
towards
towards
towards
towards
towards
towards
towards
towards
products (right)
reactants (left)
reactants (left)
products (right)
products (right)
reactants (left)
reactants (left)
products (right)
Equilibrium shifts due to stresses:
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Concentration increase shift away from
increase
Concentration decrease shift toward decrease
 pressure shifts in direction of fewer gas
molecules.
 pressure shifts in direction of more gas
molecules
 temperature favors endothermic reaction
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Shift away from heat
 temperature favors exothermic reaction
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Shift towards heat
Effect of Catalyst:
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Addition of catalysts changes the rate of
both the forward and reverse reactions.
There is no change in concentrations but
equilibrium is reached more rapidly.
Reactions that go to completion:
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Equilibrium is not reached if one of the products
is withdrawn as quickly as it is produced and no
new reactants are added.
Reaction continues until reactants are used up.
Products are removed if:
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Gases in liquid solution
Insoluble products (precipitate)
The Haber Process

Application of LeChatelier’s Principle
N2 (g) + 3 H2 (g)  2 NH3 (g) + 92 kJ
increase pressure
Shift 
decrease Temp
Shift 
remove NH3 add N2 and H2
Shift 
****Maximum yields of NH3 occurs under high
pressures, low temperatures and by constantly removing
NH3 and adding N2 & H2