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Ch. 13: Chemical Equilibrium

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Title: Ch. 13: Chemical Equilibrium


1
Ch. 13 Chemical Equilibrium
  • 13.1 The Equilibrium Condition

2
Equilibrium
  • dynamic equilibrium
  • may seem like no changes are occurring but there
    are changes
  • no NET changes
  • When did this reaction reach it?

3
reactants
products
H2O(g) CO(g) ? H2(g) CO2(g)
4
  • At equilibrium, forward and reverse reaction
    rates are ____________

5
Equilibrium
  • equilibrium position of a reaction is determined
    by
  • initial _________________
  • ____________ of reactants and products
  • degree of __________________ of reactants and
    products
  • GOAL
  • _
  • _

6
Equilibrium Position
  • lies to the left
  • more _________
  • less ___________
  • lies to the right
  • less __________
  • more _________
  • If reactants are mixed and concentrations do not
    change
  • could already be at equilibrium
  • reaction rates are so ________ that change is too
    difficult to detect

7
Ch. 13 Chemical Equilibrium
  • 13.2 Equilibrium Constant

8
Law of Mass Action
  • created in 1864 by Guldberg and Waage (Norweigen)
  • For a reaction jA kB ? lC mD
  • equilibrium constant K

9
Law of Mass Action
  • because
  • Rateforward
  • Ratereverse
  • so if Ratef Rater
  • kfAjBk krClDm

10
Example
  • Write the equilibrium expression for
  • 4NH3(g) 7O2(g) ? 4NO2(g) 6H2O(g)
  • What would it be for the reverse reaction?

11
Equilibrium Constant
  • will always have the same value at a certain
    temperature
  • no matter what amounts are added
  • ratio at equilibrium will always be same

12
Equilibrium Position
  • each set of equilibrium concentrations
  • depends on initial concentrations

13
Ch. 13 Chemical Equilibrium
  • 13.3 Equilibrium Expressions with Pressure

14
Equilibrium with Gases
  • equilibria involving gases can be described using
    __________ instead of _______________
  • N2(g) 3H2(g) ? 2NH3(g)

15
Equilibrium with Gases
  • N2(g) 3H2(g) ? 2NH3(g)

16
Calculating K from KP
  • KP KC RT can cancel out if total of
    coefficients are same on each side
  • where ?n is the difference in moles of gas on
    either side of the equation
  • ?n (lm) (jk)
  • N2(g) 3H2(g) ? 2NH3(g)
  • ?n

17
Example
  • Setup the expression for KP in terms of KC, R and
    T
  • 2NO(g) Cl2(g) ? 2NOCl(g)

18
Ch. 13 Chemical Equilibrium
  • 13.4 Heterogeneous Equilibria

19
Heterogeneous Equilibria
  • involve more than one phase
  • position of heterogeneous equilibria does NOT
    depend on amounts of
  • _
  • _
  • because their concentrations stay constant (since
    they are PURE)

20
Heterogeneous Equilibria
  • do not include liquids or solids in equilibrium
    expression
  • only include ________ and ______________

21
Example 1
  • 2H2O(l) ? 2H2(g) O2(g)
  • 2H2O(g) ? 2H2(g) O2(g)

22
Ch. 13 Chemical Equilibrium
  • 13.5/6 Applications of Equilibrium Constant (K)

23
Equilibrium Constant
  • if we know the value of K, we can predict
  • tendency of a reaction to occur
  • if a set of concentrations could be at
    equilibrium
  • equilibrium position, given initial
    concentrations

24
Equilibrium Constant
  • If you start a reaction with only reactants
  • concentration of reactants will decrease by a
    certain amount
  • concentration of products will increase by a same
    amount

25
Example 2
  • The following reaction has a K of 16. You are
    starting reaction with 9 O3 molecules and 12 CO
    molecules.
  • Find the amount of each species at equilibrium.
  • O3(g) CO(g) ? CO2(g) O2(g)

26
Example 2
O3(g) CO(g) ? O2(g) CO2(g) O3(g) CO(g) ? O2(g) CO2(g) O3(g) CO(g) ? O2(g) CO2(g) O3(g) CO(g) ? O2(g) CO2(g) O3(g) CO(g) ? O2(g) CO2(g) O3(g) CO(g) ? O2(g) CO2(g)
Initial I
Change C
Equilibrium E
27
Example 2
28
Example 2
O3(g) CO(g) ? O2(g) CO2(g) O3(g) CO(g) ? O2(g) CO2(g) O3(g) CO(g) ? O2(g) CO2(g) O3(g) CO(g) ? O2(g) CO2(g) O3(g) CO(g) ? O2(g) CO2(g)
I 9 12 0 0
C -x -x x x
E
29
Extent of a Reaction
  • If _________
  • mostly products
  • goes essentially to completion
  • lies far to right
  • If _________
  • mostly reactants
  • reaction is negligible
  • lies far to left
  • size of K and time needed to reach equilibrium
    are NOT related
  • time required is determined by reaction rate (Ea)

30
Reaction Quotient
  • Q equal to equilibrium expression but
    ___________ have to be at equilibrium
  • used to tell if a reaction is at equilibrium or
    not
  • relationship between Q and K tells which way the
    reaction will shift
  • _______ at equilibrium, no shift
  • _______ too large, forms reactants, shift to
    left
  • ______ too small, forms products, shift to right

31
Example 3
  • For the synthesis of ammonia at 500C, the
    equilibrium constant is 6.0 x 10-2. Predict the
    direction the system will shift to reach
    equilibrium in the following case
  • N2(g) 3H2(g) ? 2NH3(g)

32
Example 3
  • NH30 1.0x10-3 M,
  • N201.0x10-5 M
  • H202.0x10-3 M
  • Q __ K so forms _________, shifts to _____

33
Example 4
  • In the gas phase, dinitrogen tetroxide decomposes
    to gaseous nitrogen dioxide
  • N2O4(g) ? 2NO2(g)
  • Consider an experiment in which gaseous N2O4 was
    placed in a flask and allowed to reach
    equilibrium at a T where KP 0.133. At
    equilibrium, the pressure of N2O4 was found to be
    2.71 atm.
  • Calculate the equilibrium pressure of NO2.

34
Example 4
35
Example 5
  • At a certain temperature a 1.00 L flask initially
    contained 0.298 mol PCl3(g) and 8.70x10-3 mol
    PCl5(g). After the system had reached
    equilibrium, 2.00x10-3 mol Cl2(g) was found in
    the flask.
  • PCl5(g) ? PCl3(g) Cl2(g)
  • Calculate the equilibrium concentrations of all
    the species and the value of K.

36
Example 5
PCl5(g) ? PCl3(g) Cl2(g) PCl5(g) ? PCl3(g) Cl2(g) PCl5(g) ? PCl3(g) Cl2(g) PCl5(g) ? PCl3(g) Cl2(g)
I
C
E
37
Approximations
  • If K is very small, we can assume that the change
    (x) is going to be negligible
  • can be used to cancel out when adding or
    subtracting from a normal sized number
  • to simplify algebra

0
38
Example 6
  • At 35C, K1.6x10-5 for the reaction
  • 2NOCl(g) ? 2NO(g) Cl2(g)
  • Calculate the concentration of all species at
    equilibrium for the following mixtures
  • 2.0 mol NOCl in 2.0 L flask
  • 1.0 mol NOCl and 1.0 mol NO in 1.0 L flask
  • 2.0 mol NOCl and 1.0 mol Cl2 in 1.0 L flask

39
Example 6
  • 2.0 mol NOCl in 2.0 L flask
  • NOCl
  • NO Cl2

2NOCl(g) ? 2NO(g) Cl2(g) 2NOCl(g) ? 2NO(g) Cl2(g) 2NOCl(g) ? 2NO(g) Cl2(g) 2NOCl(g) ? 2NO(g) Cl2(g)
I
C
E
40
Example 6
  • 1.0 mol NOCl and 1.0 mol NO in 1.0 L flask
  • NOCl
  • NO Cl2

2NOCl(g) ? 2NO(g) Cl2(g) 2NOCl(g) ? 2NO(g) Cl2(g) 2NOCl(g) ? 2NO(g) Cl2(g) 2NOCl(g) ? 2NO(g) Cl2(g)
I
C
E
41
Example 6
  • 2.0 mol NOCl and 1.0 mol Cl2 in 1.0 L flask
  • NOCl
  • Cl2 NO

2NOCl(g) ? 2NO(g) Cl2(g) 2NOCl(g) ? 2NO(g) Cl2(g) 2NOCl(g) ? 2NO(g) Cl2(g) 2NOCl(g) ? 2NO(g) Cl2(g)
I
C
E
42
Ch. 13 Chemical Equilibrium
  • 13.7 Le Chatliers Principle

43
Le Châtliers Principle
  • can predict how certain changes in a reaction
    will affect the position of equilibrium

44
Changing Concentration
  • system will shift away from the added component
    or towards a removed component
  • Ex N2 3H2 ? 2NH3
  • if more N2 is added, then equilibrium position
    shifts to right
  • if some NH3 is removed, then equilibrium position
    shifts to right

45
Change in Pressure
  • adding or removing gaseous reactant or product is
    same as changing conc.
  • adding inert or uninvolved gas
  • increase the ___________________
  • ___________effect the equilibrium position

46
Change in Pressure
  • changing the volume
  • decrease V
  • decrease in gas molecules
  • shifts towards the side of the reaction with
    _____ gas molecules
  • increase V
  • increase in of gas molecules
  • shifts towards the side of the reaction with
    _____ gas molecules

47
Change in Temperature
  • all other changes alter the concentration at
    equilibrium position but dont actually change
    value of K
  • value of K does change with temperature

48
Change in Temperature
  • if energy is added, the reaction will shift in
    direction that consumes energy
  • treat energy as a
  • __________ for endothermic reactions
  • __________ for exothermic reactions

49
As4O6(s) 6C(s) ? As4(g) 6CO(g)
  • add CO
  • add C
  • remove C
  • add As4O6
  • remove As4O6
  • remove As4
  • decrease volume
  • add Ne gas

50
P4(s) 6Cl2(g) ? 4PCl3(l)
  • decrease volume
  • increase volume
  • add P4
  • remove Cl2
  • add Kr gas
  • add PCl3

51
energy N2(g) O2(g) ? 2NO(g)
  • endo or exo?
  • increase temp
  • increase volume
  • decrease temp

52
N2(g) 3H2(g) ? 2NH3(g)
53
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