Kinetics%20,%20Thermodynamics%20and%20Equilibrium

1
Kinetics , Thermodynamics and Equilibrium

• Regents Chemistry

2
Kinetics and Thermodynamics

• Kinetics deals with rates of reactions
• Thermodynamics involves changes in energy that
  occur in reactions

3
Kinetics Collision Theory

• Measured in
• moles of reactant used per unit time
• Or
• moles of product formed per unit time
• Frequency of collisions more collisions faster
  rate
• Effective collisions must have proper
  orientation and enough energy

4
Factors Affecting Rate

• 1. Type of substance
• Ionic substances react faster bonds require less
  energy to break
• Covalent react more slowly bonds require more
  energy to break

5
Factors Affecting Rate

• 2. Temperature increase
• Kinetic energy increases and the number of
  collisions increases.
• Soreactants have more energy when colliding.
  This increases rate.

6
Factors Affecting Rate

• 3. Concentration increase
• Increases rate due to the fact that more
  particles in a container, which creates more
  collisions.

7
Factors Affecting Rate

• 4. Surface Area Increase
• Increases rate by increasing reactant interaction
  or collisions

8
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
9
Factors Affecting Rate

• 6. Catalyst substance that increases rate of
  reaction, lowers 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

10
Potential Energy Diagrams

• Graphs the Change in heat during the course of a
  reaction.

11
Exothermic PE of products is less because energy
was lost.
PE of reactants (ER)
Activation Energy (Ea)
PE of Activated Complex
PE of products (EP)
Activation Energy (Ea) reverse reaction
Heat of reaction (?H) Ep - ER
12
Endothermic PE of products is more because
energy was gained.
PE of reactants (ER)
PE of products (EP)
Heat of reaction (?H)
Activation Energy (Ea)
Activation Energy (Ea) reverse reaction
PE of Activated Complex
13
Catalysts
14
Thermodynamics

• Heat content (Enthalpy) amount of heat absorbed
  or released in a chemical reaction
• Enthalpy (?H Hproducts Hreactants)

15
?H Hproducts Hreactants

• ?H is positive when the reaction is endothermic.
  Heat of products are greater than reactants
• ?H is negative when the reaction is exothermic.
  Heat of reactants were greater than the products

16
Table I

• 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

17
Table I- Practice

1. Which reaction gives off the most energy?
2. Which reaction gives off the least energy?
3. Which reaction requires the most energy to occur?

18
Entropy (?S)

• Definition randomness, disorder in a sample of
  matter
• Gases have high entropy
• Solids have low entropy

19
Increasing ?S

• Phase change from s ? l ? g
• Mixing gases
• Dissolving a substance

20
Spontaneous Reactions

• Nature favors low energy and high entropy
• Reactions are spontaneous when heat (?H)
  decreases and entropy (?S) increases
• ?H (-)
• ?S ()

21
Chemical Equilibrium

• Regents Chemistry

22
Reversible Reactions

• 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)

23
Reversible Reactions II

• 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.

24
Equilibrium

• Equilibrium is dynamic condition where rates of
  opposing processes are equal.
• Types of Equilibrium
• Phase equilibrium
• Solution Equilibrium
• Chemical Equilibrium

25
Phase Equilibrium

• 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)

26
Solution Equilibrium

• Rate of dissolving rate of crystallization
• Occurs in saturated solutions

27
Chemical Equilibrium

• Rateforward reaction Ratereverse reaction
• Concentration of reactants and products are
  constant NOT necessarily equal.

28
The Concept of Equilibrium

• 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.

29
Le Chateliers Principle

• Whenever stress put on a reaction, the reaction
  will shift its point of equilibrium
• Stresses include
• Temperature, pressure, changes in reactant or
  product concentrations

30
Example The Haber Process

• N2 (g) 3 H2 (g) ? 2 NH3 (g) heat
• ? N2
• ? H2
• ? NH3
• ? NH3
• ? pressure
• ? pressure
• ? temperature
• ? temperature

31
Example The Haber Process

• N2 (g) 3 H2 (g) ? 2 NH3 (g) heat
• ? N2 shift towards products (right)
• ? H2 shift towards reactants (left)
• ? NH3 shift towards reactants (left)
• ? NH3 shift towards products (right)
• ? pressure shift towards products (right)
• ? pressure shift towards reactants (left)
• ? temperature shift towards reactants (left)
• ? temperature shift towards products (right)

32
Equilibrium shifts due to stresses

• 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
• Shift away from heat
• ? temperature favors exothermic reaction
• Shift towards heat

33
Effect of Catalyst

• 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.

34
Reactions that go to completion

• 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
• Gases in liquid solution
• Insoluble products (precipitate)

35
The Haber Process

• Application of LeChateliers 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