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Chapter 12:. Chemical Equilibrium. The Dynamic Nature of Equilibrium. A. What is equilibrium? 1. Definition a state of balance; no net change in a dynamic process. 2. Definition of chemical equilibrium.

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chapter 12

Chapter 12:

Chemical Equilibrium

the dynamic nature of equilibrium
The Dynamic Nature of Equilibrium
  • A. What is equilibrium?
  • 1. Definition
      • a state of balance; no net change in a dynamic process
2 definition of chemical equilibrium
2. Definition of chemical equilibrium

No net change with the total amount of reactants and products remaining constant, while the reaction continues

Eek = equilibrium

14 1 the dynamic nature of equilibrium
14.1 The Dynamic Nature of Equilibrium
  • 3. General Characteristics
      • Double Headed Arrow
      • Beginning of Rxn  Form lots of products (before eek is established)
      • Moments Later  Forming both products and reactants (eek)
        • **Not Necessarily equal proportions of both sides of the reaction!
14 1 the dynamic nature of equilibrium1
14.1 The Dynamic Nature of Equilibrium
  • B. Dynamic Equilibrium - Characteristics
      • two opposing processes occur at exactly the same rate
      • Rate of:

Reactants  Products = Products  Reactants

      • Dynamic Eek will not occur as soon as the reversible reaction begins

Like all good things, it takes time!

c graph
C. Graph
  • The N2O4 2NO2 Equilibrium System
    • Initially there is a large amount of N2O4 and no NO2
    • Equilibrium is established when the amounts plateau
d the 2no 2 n 2 o 4 equilibrium system
D. The 2NO2N2O4 Equilibrium System
  • 1. Description
    • Dynamic, reversible, no net change
  • 2. Equilibrium Conditions
      • Specific for a Reaction (will be unique)
        • MUST have the Balanced Written Equation
        • Coefficients will matter!
      • Temperature Dependent
      • CONSTANT regardless of concentration
the equilibrium expression k eq equilibrium law or law of mass action
The Equilibrium Expression, Keq(Equilibrium Law or Law of Mass-Action)
  • A. Writing Expressions for Keq
      • aA + bB <-----> cC + dD
      • Keq = (PC)c (PD)d_ for gases use pressures

(PA)a (PB)b

      • Keq = [C]c [D]d _ for aqueous solutions use

[A]a [B]b Molarities

the equilibrium expression
The Equilibrium Expression
  • Example: Write the Keq for:
      • 4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g)
      • N2(g) + 3H2(g)  2NH3(g)
b characteristics of k c
B. Characteristics of Kc
  • 1. Independent of :
    • Pure solids and pure liquids as long as some of the substance is present
  • 2. Dependent on:
    • Gaseous substances and solutions
c the meaning of k c
C. The meaning of Kc
  • Kc = 1; reactants and products are present in equal amounts at eek
  • Kc > 1; products are present in greater proportion at eek
  • Kc < 1; reactants are present in greater proportion at eek
slide12
4. Examples: Who will be favored, in other words, be in greater proportion? Answer in a complete sentence.

N2(g) + O2(g)  2 NO(g) Keq=55

N2(g) + 3H2(g)  2NH3(g) Keq=5x10-6

iii modifying equilibrium constant expressions
III. Modifying Equilibrium Constant Expressions
  • A. Reversible reactions
  • Rule:
    • The Keq value for the reverse of a reaction will be the reciprocal of the forward reaction.
iii modifying equilibrium constant expressions1
III. Modifying Equilibrium Constant Expressions
  • Example: Write the Keq for the following equation and its reverse.

A(g) + B(g) C(g)

iii modifying equilibrium constant expressions2
III. Modifying Equilibrium Constant Expressions
  • B. Summation of Reactions
  • Rule:

The Keq of summed reactions will be the product of the Keq’s

Rule:

If you multiply the coefficients in an equation by a factor you must raise Keq to the power of that factor

example
Example:

A(g) + B(g)  C(g) Kp=3

A(g) + B(g)  D(g) Kp=6

What is the Kp for C(g)  D(g)?

example1
Example:

½ N2(g)+ ½ O2(g)NO(g) Keq=6.9x10-16

NO2(g) NO(g) + ½ O2(g) Keq=6.7x10-7

N2O4(g) 2NO2(g) Keq=0.15

What is the Keq for N2(g)+2O2(g)N2O4(g)

c heterogeneous systems
C. Heterogeneous Systems

Rule: Systems where all the substances are not in the same phase of matter. Specifically, the reaction includes at least one pure solid or liquid which would not be included in the Keq expression.

slide19
Example: Write the Keq for:

a. Zn(s) + 2H+(aq) Zn2+(aq) + H2(g)

b. CaCO3(s) CaO(s) + CO2(g)

iv determination of k
IV. Determination of K
  • A. Calculating Keq from experimental values (Intro to IRE problems).
  • Example: Write the expression for the equilibrium reaction between solid ammonium chloride and gaseous products, hydrogen chloride and ammonia (NH3). At equilibrium in a 1 liter container, the following amounts are present: 12.0 mol ammonium chloride, 3.0 mol of ammonia and 5.0 mol of hydrogen chloride. Determine the Keq.
slide22
Example: When 4.29 moles of PCl3(g) and 4.29 moles of Cl2(g) are placed in a 1.00 Liter container at 250oC, the following equilibrium is established:

PCl3(g) + Cl2(g)  PCl5(g)

The equilibrium concentration of phosphorus pentachloride is 2.59 mole/L. What are the equilibrium concentrations of the other two gases? Calculate Keq for the above reaction system.

slide23
Example: Consider the equilibrium system:

2NO(g) + Br2(g)  2NOBr(g)

At a given temperature, 1.6 mol of NO and 1.6 mol of Br2 are added to a 1.00 Liter flask and the equilibrium concentration of NOBr is found to be 0.53 mol/L. Calculate the equilibrium concentrations of the other 2 gases and the value of Keq.

b ire problems with unknown equilbrium concentrations
B. IRE Problems with unknown equilbrium concentrations

Description of problems:

set one of the “R” values as X and determine the others by their mole (coefficient) relationships

14 5 some illustrative equilibrium calculations
14.5 Some Illustrative Equilibrium Calculations
  • 2. Example Carbon dioxide and hydrogen gas at concentrations of 1.00 M each are introduced into a container and the following system is established: CO2(g) + H2(g)  CO(g) + H2O(g) Kc = 0.64 What are the equilibrium concentrations of all the species?
14 5 some illustrative equilibrium calculations1
14.5 Some Illustrative Equilibrium Calculations
  • 3. Example: For the system: H2(g) + I2(g)  2HI(g) at 425 oC, K = 55.5. If 2.5 mol of hydrogen and 2.5 mol of iodine are placed in a 1-liter vessel and heated to 425oC and the system reaches equilibrium, what are the concentrations of all species?
c applications of the equilibrium constant
C. Applications of the Equilibrium Constant
  • The Reaction quotient = Q

(Direction of Shift to Reach Equilibrium)

a. Q = the same expression as for Keq

The values for each substance are not equilibrium values but the actual concentrations at any moment in time.

slide28
b. If Q < Kc , then the reaction will proceed to the right

c. If Q = Kc , then the system is at equilibrium

d. If Q > Kc , then the reaction will proceed to the left

slide29
Example: For N2O4(g)2NO2(g) Kc=11.0 at 100°C. Is the system at equilibrium if the concentration of both gases is 0.20M? In which direction will the system shift to reach equilibrium?
slide30
Example: Answer the same question for the system above if the gas concentrations are 0.07M for N2O4 and 1.00M for NO2.
slide31
If the system is started with 0.50 mol of each of the reactants and 9.30 mole of the HI in a 5.0L container, predict the direction the system will shift to reach equilibrium.
14 4 qualitative treatment of equilibrium le chatelier
14.4 Qualitative Treatment of Equilibrium: Le Chatelier
  • A. Statement of Le Chatelier’s Principle

“When a system is stressed (changes in concentration, temp, gas pressure, or volume of container), the system will respond by attaining new equilibrium conditions that counteract the change”

14 4 qualitative treatment of equilibrium le chatelier1
14.4 Qualitative Treatment of Equilibrium: Le Chatelier
  • B. For the system, N2O4(g)  2NO2(g) H = 57.2 kJ
  • 1. Changes in amounts of species
  • a. Adding or removing reactant
      • Add Reactant  Eek shifts toward products (use it up)
      • Remove Reactant  Eek shifts toward reactants (make more)
14 4 qualitative treatment of equilibrium le chatelier2
14.4 Qualitative Treatment of Equilibrium: Le Chatelier
  • b. Adding or removing product
      • Add Product  Eek shifts toward reactants (use it up)
      • Remove Product  Eek shifts toward products (make more)
  • c. Adding or removing pure solid or solvent
      • Has NO effect on equilibrium (no shift)
  • d. Adding inert substance (inert = non-reactive)
      • Has NO effect on equilibrium (no shift)
14 4 qualitative treatment of equilibrium le chatelier3
14.4 Qualitative Treatment of Equilibrium: Le Chatelier
  • 2. Changes in pressure
      • Pressure increased  Eek shifts in the direction producing the smaller number of moles of gas
      • Pressure decreased  Eek shifts in the direction producing the larger number of moles of gas
      • Moles of gas equal on both sides = pressure does not effect Eek.
14 4 qualitative treatment of equilibrium le chatelier4
14.4 Qualitative Treatment of Equilibrium: Le Chatelier
  • 3. Changes in volume
      • Decreases in volume = increase in pressure
      • Increases in volume = decrease in pressure
14 4 qualitative treatment of equilibrium le chatelier5
14.4 Qualitative Treatment of Equilibrium: Le Chatelier
  • 4. Changes in temperature
      • Increase Temp  Eek shifts in the direction of the endothermic rxn; away from heat
      • Decrease Temp  Eek shifts in the direction of the exothermic rxn; toward heat
14 4 qualitative treatment of equilibrium le chatelier6
14.4 Qualitative Treatment of Equilibrium: Le Chatelier
  • Example: List all of the ways to shift the equilibrium of the following systems to the right:
  • a. N2(g) + 3H2(g)  2NH3(g)
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