Kinetics , Thermodynamics and Equilibrium

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Kinetics , Thermodynamics and Equilibrium

• Regents Chemistry

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Kinetics and Thermodynamics

• Kinetics deals with rates of reactions (how
  quickly a reaction occurs)
• Thermodynamics involves changes in energy that
  occur in reactions

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Kinetics Collision Theory

• 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

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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 NO3- Na 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)

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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.

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Factors Affecting Rate

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

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Factors Affecting Rate

• 4. Surface Area Increase
• Increases rate due to increased reactant
  interaction or collisions (powder vs. lump)

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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
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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

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Quick Review Factors that affect reactions

• Ionic solutions have faster reactions than
  molecule compounds. (bonding)
• ?Temp. ? Rate
• ? conc. ?rate
• ? surface area ? rate
• ? Pressure ? rate, ? P ? rate
• Catalysts speed up reactions.

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Potential Energy Diagrams

• Graphs heat during the course of a reaction.

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

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

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?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

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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

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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?

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Entropy (?S)

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

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Increasing ?S

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

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Spontaneous Reactions

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

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Analogy Your Bedroom

• 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.

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Stability of Products and ?H

• Help determine if a reaction is spontaneous
• Products tend toward Lower energy (-?H)
• Products tend toward more randomness (?S)
• 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.

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Chemical Equilibrium

• Regents Chemistry

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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)

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Reversible Reactions cont.

• 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.

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Equilibrium

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

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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)

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Solution Equilibrium

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

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Chemical Equilibrium

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

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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.

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Le Chateliers Principle

• 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

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Example The Haber Process

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

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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)

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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

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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.

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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)

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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