Chapter 16: Chemical Equilibrium General Concepts

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Chapter 16 Chemical Equilibrium- General Concepts

• When a system is at equilibrium, the forward and
  reverse reaction are proceeding at the same rate
• The concentrations of all species remains
  constant over time, but both the forward and
  reverse reaction never cease
• Equilibrium is signified with double arrows or
  the equal sign

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The decomposition of N2O4(g) into NO2(g). The
concentrations of N2O4 and NO2 change relatively
quickly at first, but eventually stop changing
with time when equilibrium is reached.
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• The equilibrium mixture is independent of whether
  we start on the reactant side or the product
  side

The equilibrium between N2O4 and NO2.
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The same equilibrium composition is reached from
either the forward or reverse direction, provided
the overall system composition is the same. Pure
NO2 is brown and pure N2O4 is colorless. The
amber color of the equilibrium mixture indicates
that both species are present at equilibrium.
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• There is a simple relationship among the
  concentrations of the reactants and products for
  any chemical system at equilibrium
• It is called the mass action expression, and is
  derived from thermodynamics (discussed in Chapter
  20)
• Consider the equilibrium

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Four experiments to study the equilibrium among
H2, I2, and HI gases. Different amounts of the
reactants and products are placed in a 10.0 L
reaction vessel at 440oC where the gases
establish equilibrium. When equilibrium is
reached, different amounts of reactants and
products remain.

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• The numerical value of the mass action expression
  is called the reaction quotient, Q

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• The reaction can be evaluated at any
  concentrations
• At equilibrium (and 440oC) for this reaction the
  reaction quotient has the value 49.5 (a unitless
  number)
• This relationship is called the equilibrium law
  for the system

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• The value 49.5 is called the equilibrium
  constant, Kc, and characterizes the system
• For chemical equilibrium to exist, the reaction
  quotient Q must be equal to the equilibrium
  constant Kc
• Consider the general chemical equation

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• The exponents in the mass action expression are
  the same as the stoichiometric coefficients
• At equilibrium
• The form is always products over reactants
  raised to the appropriate powers

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• Various operations can be performed on
  equilibrium expressions
• Changing the direction of equilibrium when the
  direction of an equilibrium is reversed, the new
  equilibrium constant is the reciprocal of the
  original

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• Multiplying the coefficients by a factor when
  the coefficients in an equation are multiplied by
  a factor, the equilibrium constant is raised to a
  power equal to that factor

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• Adding chemical equilibria when chemical
  equilibria are added, their equilibrium constants
  are multiplied

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• The gas law can be used to write the equilibrium
  constant in terms of partial pressures
• Equilibrium constants written in terms of partial
  pressures are given the symbol Kp

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• The size of the equilibrium constant gives a
  measure of how the reaction proceeds
• General statements can be made about the
  equilibrium constant (either Kc or KP)

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The magnitude of K and the position of
equilibrium. A large amount of product and very
little reactant at equilibrium gives Kgtgt1 (large
K). When , approximately equal amounts
of reactant and product are present at
equilibrium. When Kltlt1, mostly reactant and very
little product are present at equilibrium.
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• The two different forms of the equilibrium
  constants can be related

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• In a homogeneous reactions, all the reactants and
  products are in the same phase
• Heterogeneous reactions involve more than one
  phase
• For example the thermal decomposition of sodium
  bicarbonate (baking soda)
• Heterogeneous reactions can come to equilibrium
  just like homogeneous systems

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• If NaHCO3 is placed in a sealed container,
  homogeneous equilibrium is established
• The equilibrium law involving pure liquids and
  pure solids can be simplified

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• For a pure liquid or solid, the ratio of amount
  of substance to volume of substance is constant

The concentration of a substance in a solid is
constant. Doubling the number of moles doubles
the volume, but the ratio of moles to volume
remains the same.
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• The equilibrium law for a heterogeneous reaction
  is written without concentrations terms for pure
  solids or pure liquids.
• The equilibrium constants found in tables
  represent all the constants combined

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• According to Le Châteliers principle
• If an outside influence upsets an equilibrium,
  the system undergoes a change in the direction
  that counteracts the disturbing influence and, if
  possible, returns the system to equilibrium
• We can consider some common stresses
• Adding or removing a product or reactant
• The equilibrium shifts to remove reactants or
  products that have been added
• The equilibrium shifts to replace reactants or
  products that have been removed

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• Changing the volume
• Reducing the volume of a gaseous reaction causes
  the reaction to decreases the number of molecules
  of gas, if it can
• Moderate pressure changes have a negligible
  effect on reactions involving only liquids or
  solids
• Changing the temperature
• Increasing the temperature shifts a reaction in a
  direction that produces an endothermic
  (heat-absorbing) change
• Decreasing the temperature shifts a reaction in a
  direction that produces an exothermic
  (heat-releasing) change

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• Catalysts have no effect on the position of
  equilibrium
• Catalysts change how fast a system achieves
  equilibrium, not the relative distribution of
  reactants and products
• Adding an inert gas at constant volume
• If the added gas cannot react with any reactants
  or products it is inert towards the substances in
  the equilibrium
• No concentration changes occur, so Q still equals
  K and no shift in equilibrium occurs

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• Equilibrium calculations can be divided into two
  main categories
• Calculating equilibrium constants from known
  equilibrium concentrations or partial pressures
• Calculating one or more equilibrium
  concentrations or partial pressures using the
  known value of Kc or KP
• Consider the decomposition of N2O4

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• Calculating the equilibrium constant this way is
  easy

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• More commonly, you will have a set of initial
  conditions and an equilibrium constant
• If a KP describes the system, equilibrium will
  usually be described in terms of partial
  pressures
• If a Kc describes the system, equilibrium will
  usually be describe in terms of concentration
  (molarity, mol/L)
• The Initial, Change, Equilibrium or ICE table
  is a useful way to summarize the problem

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• Example Ethyl acetate, CH3CO2C2H5, is produced
  from acetic acid and ethanol by the reaction
• At 25oC, Kc4.10 for this reaction. Suppose
  0.100 mol of ethyl acetate and 0.150 mol of water
  are placed in a 1.00 L reaction vessel. What are
  the concentrations of all species at equilibrium?
• ANALYSIS Use an ICE table and the equilibrium
  constant to find the concentrations.

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• This can be solved by putting it in quadratic
  form

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• Negative concentrations are not allowed, so
• A similar procedure can be used to calculate
  partial pressures using KP

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• Sometime simplifications can be made
• Example Nitrogen and oxygen react to form
  nitrogen monoxide
• with Kc4.8x10-31. In air at 25oC and 1 atm, the
  N2 concentrations and O2 are initially 0.033 M
  and 0.00810 M. What are the equilibrium
  concentrations?
• ANALYSIS The equilibrium constant is very small,
  very little of the reactants will be converted
  into products

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• Substituting
• N20.033-x0.033 M
• O20.00810-x0.030810 M
• NO2x1.60x10-17 M