Chemical Equilibrium

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Chemical Equilibrium
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The Equilibrium State
Chemical Equilibrium The state reached when the
concentrations of reactants and products remain
constant over time.
Brown
Colorless
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• Chemical equilibrium is achieved when
• the rates of the forward and reverse reactions
  are equal and
• the concentrations of the reactants and products
  remain constant

Physical equilibrium
Chemical equilibrium
brown
colorless
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Demo
Start with NO2
Start with N2O4
Start with NO2 N2O4
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(No Transcript)
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Using Equilibrium Constants 01

• We can make the following generalizations
  concerning the composition of equilibrium
  mixtures
• If Kc gt 103, products predominate over reactants.
  If Kc is very large, the reaction is said to
  proceed to completion.
• If Kc is in the range 103 to 103, appreciable
  concentrations of both reactants and products are
  present.
• If Kc lt 103, reactants predominate over
  products. If Kc is very small, the reaction
  proceeds hardly at all.

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Using the Equilibrium Constant
Kc 4.2 x 10-48
(at 500 K)
Kc 57.0
(at 500 K)
Kc 2.4 x 1047
(at 500 K)
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4.63 x 10-3
Law of Mass Action
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216
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Write the Equilibrium Constant Kc
If you multiply both side of an equilibrium
reaction by n the equilibrium constant should be
raised to the power of n.
Kc

Kc


Kc
Kc
2
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Homogenous equilibrium applies to reactions in
which all reacting species are in the same phase.
Kp Kc(RT)Dn
Dn moles of gaseous products moles of gaseous
reactants
(c d) (a b)
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Homogeneous Equilibrium
H2O constant
Kc
The concentration of pure liquids are not
included in the expression for the equilibrium
constant.
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The equilibrium concentrations for the reaction
between carbon monoxide and molecular chlorine to
form COCl2 (g) at 740C are CO 0.012 M, Cl2
0.054 M, and COCl2 0.14 M. Calculate the
equilibrium constants Kc and Kp.
Kc
220
Kp Kc(RT)Dn
Dn 1 2 -1
R 0.0821
T 273 74 347 K
Kp 220 x (0.0821 x 347)-1 7.7
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The equilibrium constant Kp for the reaction is
158 at 1000K. What is the equilibrium pressure
of O2 if the PNO 0.400 atm and PNO 0.270 atm?
2
347 atm
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Heterogenous equilibrium applies to reactions in
which reactants and products are in different
phases.
CaCO3 constant CaO constant
Kc CO2
The concentration of pure solids and pure liquids
are not included in the expression for the
equilibrium constant.
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Kp
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Consider the following equilibrium at 295
K The partial pressure of each gas is 0.265
atm. Calculate Kp and Kc for the reaction?
0.265 x 0.265 0.0702
Kp Kc(RT)Dn
Kc Kp(RT)-Dn
Dn 2 0 2
T 295 K
Kc 0.0702 x (0.0821 x 295)-2 1.20 x 10-4
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Writing Equilibrium Constant Expressions

• The concentrations of the reacting species in the
  liquid phase are expressed in M. In the gaseous
  phase, the concentrations can be expressed in M
  or in atm.
• The concentrations of pure solids and pure
  liquids, do not appear in the equilibrium
  constant expressions.
• The equilibrium constant is a dimensionless
  quantity.
• In quoting a value for the equilibrium constant,
  you must specify the balanced equation and the
  temperature.
• If a reaction can be expressed as a sum of two or
  more reactions, the equilibrium constant for the
  overall reaction is given by the product of the
  equilibrium constants of the individual reactions.

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Demo
Review
Start with NO2
Start with N2O4
Start with NO2 N2O4
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The reaction quotient (Qc) is calculated by
substituting the initial concentrations of the
reactants and products into the equilibrium
constant (Kc) expression.

• IF
• Qc gt Kc system proceeds from right to left to
  reach equilibrium
• Qc Kc the system is at equilibrium
• Qc lt Kc system proceeds from left to right to
  reach equilibrium

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Use reaction quotient to predict the direction of
shift when the volume is halved in the following
equilibrium N2O4(g) æ 2 NO2(g),

• Consider the reaction N2O4(g) æ 2 NO2(g), taking
  place in a cylinder with a volume 1 unit.

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Le Châteliers Principle 10

• The Volume is then halved, which is equivalent to
  doubling the pressure.
• Since Q gt K, the product is too high and the
  reaction progresses in the reverse direction.

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At 12800C the equilibrium constant (Kc) for the
reaction Is 1.1 x 10-3. If the initial
concentrations are Br2 0.063 M and Br
0.012 M, calculate the concentrations of these
species at equilibrium.
Let (x) be the change in concentration of Br2
Initial (M)
0.063
0.012
ICE
Change (M)
-x
2x
Equilibrium (M)
0.063 - x
0.012 2x
Solve for x
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4x2 0.048x 0.000144 0.0000693 0.0011x
4x2 0.0491x 0.0000747 0
ax2 bx c 0
x -0.00178
x -0.0105
At equilibrium, Br 0.012 2x -0.009 M
or 0.00844 M
At equilibrium, Br2 0.063 x 0.0648 M
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Calculating Equilibrium Concentrations

1. Express the equilibrium concentrations of all
   species in terms of the initial concentrations
   and a single unknown x, which represents the
   change in concentration.
2. Write the equilibrium constant expression in
   terms of the equilibrium concentrations. Knowing
   the value of the equilibrium constant, solve for
   x.
3. Having solved for x, calculate the equilibrium
   concentrations of all species.

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Predicting the direction of a Reaction

• The reaction quotient (Qc) is obtained by
  substituting initial concentrations into the
  equilibrium constant. Predicts reaction
  direction.Qc gt Kc System proceeds to form
  reactants.Qc Kc System is at equilibrium.Qc
  lt Kc System proceeds to form products.

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Le Châteliers Principle 01

• Le Châteliers principle If an external stress
  is applied to a system at equilibrium, the
  system adjusts in such a way that the stress
  is partially offset.

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Le Châteliers Principle 02

• Concentration Changes
• The concentration stress of an added reactant or
  product is relieved by reaction in the direction
  that consumes the added substance.
• The concentration stress of a removed reactant or
  product is relieved by reaction in the direction
  that replenishes the removed substance.

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Le Châteliers Principle Haber process
N2(g) 3 H2(g) æ 2 NH3(g)
Exothermic
Cat  iron or  ruthenium 
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Le Châteliers Principle 06

• The reaction of iron(III) oxide with carbon
  monoxide occurs in a blast furnace when iron ore
  is reduced to iron metal
• Fe2O3(s) 3 CO(g) æ 2 Fe(l) 3 CO2(g)
• Use Le Châteliers principle to predict the
  direction of reaction when an equilibrium mixture
  is disturbed by
• (a) Adding Fe2O3 (b) Removing CO2 (c) Removing CO

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Le Châteliers Principle 07

• Volume and Pressure Changes Only reactions
  containing gases are affected by changes in
  volume and pressure.
• Increasing pressure Decreasing volume
• PV nRT tells us that increasing pressure or
  decreasing volume increases concentration.

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Le Châteliers Principle 08

• N2(g) 3 H2(g) æ 2 NH3(g) Kc 0.291 at
  700 K

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Le Châteliers Principle 11

• Does the number of moles of reaction products
  increase, decrease, or remain the same when each
  of the following equilibria is subjected to a
  decrease in pressure by increasing the volume.
• PCl5(g) æ PCl3(g) Cl2(g)
• CaO(s) CO2(g) æ CaCO3(s)
• 3 Fe(s) 4 H2O(g) æ Fe3O4(s) 4 H2(g)

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Le Châteliers Principle 13

• Temperature Changes Changes in temperature can
  change the equilibrium constant.
• Endothermic processes are favored when
  temperature increases.
• Exothermic processes are favored when
  temperature decreases.

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Le Châteliers Principle 14

• Example
• The reaction N2(g) 3 H2(g) æ 2 NH3(g) which is
  exothermic by 92.2 kJ.

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Le Châteliers Principle 15

• In the first step of the Ostwald process for
  synthesis of nitric acid, ammonia is oxidized to
  nitric oxide by the reaction
• 4 NH3(g) 5 O2(g) æ 4 NO(g) 6 H2O(g) ?H
  905.6 kJ
• How does the equilibrium amount vary with an
  increase in temperature?

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Le Châteliers Principle 16

• The following pictures represent the composition
  of the equilibrium mixture at 400 K and 500 K for
  the reaction A(g) B(g) æ AB(g).
• Is the reaction endothermic or exothermic?

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Effect of Catalysis, Reduction of Activation
Energy
No effect