Overview of Ch 11-13

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• Overview of Ch 11-13

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Properties of SolutionsChapter 11
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Solution Composition

• 1. Molarity (M)
• 2. Mole fraction (?A)
• 3. Molality (m)

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Henrys Law
The amount of a gas dissolved in a solution is
directly proportional to the pressure of the gas
above the solution.

• P kC
• P partial pressure of gaseous solute above the
  solution
• C concentration of dissolved gas
• k the Henrys Law constant

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Temperature Effects

• Solubility of gases generally decreases with
  temperature.
• Solubility of solids generally increases with
  temperature.

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Colligative Properties

• Depend only on the number, not on the identity,
  of the solute particles in an ideal solution.
• Vapor pressure depression
• Boiling point elevation
• Freezing point depression
• Osmotic pressure increase

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Raoults Law
The presence of a nonvolatile solute lowers the
vapor pressure of a solvent.

• Psoln ?solvent P?solvent
• Psoln vapor pressure of the solution
• ?solvent mole fraction of the solvent
• P?solvent vapor pressure of the pure solvent

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Boiling Point Elevation

• A nonvolatile solute elevates the boiling point
  of the solvent.
• ?T Kbmsolute
• Kb molal boiling point elevation constant
• m molality of the solute

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Freezing Point Depression

• A nonvolatile solute depresses the freezing point
  of the solvent.
• ?T Kfmsolute
• Kf molal freezing point depression constant
• m molality of the solute

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Osmotic Pressure

• Osmosis The flow of solvent into the solution
  through a semipermeable membrane.
• Osmotic Pressure A nonvolatile solute increases
  the osmotic pressure of the solvent.

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Chemical EquilibriumChapter 13

• The state where the concentrations of all
  reactants and products remain constant with time.

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Equilibrium Constant

• jA kB ? lC mD
• The equilibrium expression

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• 4NH3(g) 7O2(g) ? 4NO2(g) 6H2O(g)

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Manipulations of K

• The equilibrium constant for a reaction is the
  reciprocal of that for the reaction written in
  reverse.
• When the equation for a reaction is multiplied by
  n, Knew (Koriginal)n

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K v. Kp

• For
• jA kB ? lC mD
• Kp K(RT)?n
• ?n sum of coefficients of gaseous products
  minus sum of coefficients of gaseous reactants.

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Heterogeneous Equilibria

• . . . are equilibria that involve more than one
  phase.
• CaCO3(s) ? CaO(s) CO2(g)
• K CO2
• The position of a heterogeneous equilibrium does
  not depend on the amounts of pure solids or
  liquids present.

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Reaction Quotient

• . . . helps to determine the direction of the
  move toward equilibrium.
• The law of mass action is applied with initial
  concentrations.

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• H2(g) F2(g) ? 2HF(g)

• Q lt K, shift right
• Q gt K, shift left

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Solving Equilibrium Problems

• 1. Write the equilibrium expression.
• 2. Set up an ICE box with relevant
  concentrations.
• 3. Use the stoichiometry of the reaction to
  determine changes in products and reactants,
  solving for unknowns.

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Le Châteliers Principle

• . . . if a change is imposed on a system at
  equilibrium, the position of the equilibrium will
  shift in a direction that tends to reduce that
  change.

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Effects of Changes on the System

• 1. Concentration The system will shift away
  from the added component.
• 2. Temperature treat the energy change as a
  reactant (endothermic) or product exothermic).

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Effects of Changes on the System (continued)

• 3. Pressure
• a. Addition of inert gas does not affect the
  equilibrium position.
• b. Decreasing the volume shifts the
  equilibrium toward the side with fewer moles.

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Chemical KineticsChapter 12

• The area of chemistry that concerns reaction
  rates.

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Reaction Rate

• Change in concentration (conc) of a reactant or
  product per unit time.

Reaction rates are positive by convention.
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(Differential) Rate Laws

• Rate kNO2n
• k rate constant
• n rate order

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Types of Rate Laws

• Differential Rate Law expresses how rate
  depends on concentration.
• Integrated Rate Law expresses how
  concentration depends on time.

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Method of Initial Rates

• Initial Rate the instantaneous rate just
  after the reaction begins.
• The initial rate is determined in several
  experiments using different initial
  concentrations.

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Overall Reaction Order

• Sum of the order of each component in the rate
  law.
• rate kH2SeO3H2I?3
• The overall reaction order is 1 2 3 6.

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First-Order Rate Law
For aA ? Products in a 1st-order reaction,

• Integrated first-order rate law is
• lnA ?kt lnAo

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Half-Life of a 1st-Order Rxn

• t1/2 half-life of the reaction
• k rate constant
• For a first-order reaction, the half-life does
  not depend on concentration.

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Second-Order Rate Law

• For aA ? products in a second-order reaction,

• Integrated rate law is

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Half-Life of a 2nd-Order Rxn

• t1/2 half-life of the reaction
• k rate constant
• Ao initial concentration of A
• The half-life is dependent upon the initial
  concentration.

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Zero-Order Rate Law

• For aA ? products in a zero-order reaction,

Rate k

• Integrated rate law is

A -kt Ao
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Half-Life of a Zero-Order Rxn
t1/2 Ao
2k

• t1/2 half-life of the reaction
• k rate constant
• Ao initial concentration of A
• The half-life is dependent upon the initial
  concentration.

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Reaction Mechanism

• The series of steps by which a reaction occurs.
• A chemical equation does not tell us how
  reactants become products - it is a summary of
  the overall process.

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Reaction Mechanism (continued)

• The reaction
• has many steps in the reaction mechanism.

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• Intermediate formed in one step and used up in
  a subsequent step and so is never seen as a
  product.
• Molecularity the number of species that must
  collide to produce the reaction indicated by that
  step.
• Elementary Step A reaction for which a rate law
  can be written from its molecularity.
• uni, bi and termolecular

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Rate-Determining Step

• In a multistep reaction, it is the slowest step.
  It therefore determines the rate of reaction.

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Arrhenius Equation

• Collisions must have enough energy to produce the
  reaction (must equal or exceed the activation
  energy).
• Orientation of reactants must allow formation of
  new bonds.

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• k rate constant
• A frequency factor
• Ea activation energy
• T temperature (in K)
• R gas constant

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lnk -Ea 1 ln A R T
slope -Ea /R

• catalysts decrease Ea.