This document discusses the concept of chemical equilibrium. Some key points: 1. Chemical equilibrium is a dynamic state where the rates of the forward and reverse reactions are equal, so concentrations of reactants and products no longer change. 2. The equilibrium constant, Kc, relates concentrations of products and reactants at equilibrium. When the reaction quotient, Q, equals Kc the system is at equilibrium. 3. Le Chatelier's principle states that if a system at equilibrium experiences a change, it will shift in a way to counteract the change and reestablish equilibrium. Changes in concentration, pressure, and temperature can disturb equilibrium. 3. Increasing temperature favors the endothermic reaction, while

1. CHEMICAL EQUILIBRIUM
It is a dynamic state of things.
Chapters 4.7 & 17 - Silberberg

2. Goals & Objectives
 See the following Learning
Objectives on page 773.
 Understand These Concepts:
 17.1-7, 9, 11-16.
 Master These Skills:
 17.1-4, 6-17

3. The Equilibrium State
All reactions are reversible and under suitable conditions will reach a
state of equilibrium.
At equilibrium, the concentrations of products and reactants no longer
change because the rates of the forward and reverse reactions are
equal.
At equilibrium: rateforward = ratereverse
Chemical equilibrium is a dynamic state because reactions continue to
occur, but because they occur at the same rate, no net change is
observed on the macroscopic level.

4. Basic Concepts
 Most chemical reactions do not go
to completion.
 Reversible reactions do not go to
completion and can occur in either
direction.

5. Basic Concepts
 Reversible reactions may be
represented in general terms as the
following:
 aA(g) + bB(g) = cC(g) + dD(g)
 Chemical equilibrium exists when
the forward and reverse reactions
are occuring at exactly the same
rate and therefore there is no net
change in the concentration of
reactants and products.

6. Basic Concepts
 For the generalized reaction
 A(g) + B(g) = C(g) + D(g)
 the reaction may be represented
graphically

7. Figure 17.3 The change in Q during the N2O4-NO2 reaction.

8. Basic Concepts
 It makes no difference if we start
with the “reactants” or the
“products” of a reversible reaction,
because equilibrium can be
established from either direction.
 Once established, the equilibrium is
“dynamic”--it continues with no net
change.

9. The Equilibrium Constant
 The general form for the equilibrium
constant for the reaction
 aA(g) + bB(g) = cC(g) + dD(g)
 Kc = [C]c[D]d
 [A]a[B]b

10. The Equilibrium Constant
 Write equilibrium constant
expressions for the following
reactions at 500oC.
 1. PCl5(g) = PCl3(g) + Cl2(g)
 2. H2(g) + I2(g) = 2HI(g)
 3. 4NH3(g) + 5O2(g) = 4NO(g)
+ 6H2O(g)

12. The Equilibrium Constant
 The thermodynamic definition of the
equilibrium constant involves activities
rather than concentrations.
 For pure liquids and solids, the activity is
taken to be 1.
 For components of ideal solutions, the
activity of each component is taken to be
its molar concentration.
 For mixtures of ideal gases, the activity of
each component is taken to be its partial
pressure in atmospheres.

13. The Equilibrium Constant
 Because of the use of activities,
equilibrium constants have no units.

14. Exercises
 One liter of the following reaction
system at a high temperature was
found to contain 0.172 moles of
phosporous trichloride, 0.086 moles
of chlorine and 0.028 moles of
phosporous pentachloride at
equilibrium. Determine the value of
Kc for the reaction.
 PCl5 (g) = PCl3 (g) + Cl2(g)

16. Exercises
 The decomposition of PCl5 was
studied at another temperature.
One mole of PCl5 was introduced
into an evacuated 1.00L container
at the new temperature. At
equilibrium, 0.60 moles of PCl3 were
present in the container. Calculate
the equilibrium constant at this
temperature.
 PCl5 (g) = PCl3 (g) + Cl2(g)

18. Exercises
 At a given temperature, 0.80 moles
of N2 and 0.90 moles of H2 were
placed in an evacuated 1.00L
container. At equilibrium, 0.20
moles of NH3 were present.
Determine the value for Kc for the
reaction.
 N2(g) + 3H2(g) = 2NH3(g)

20. The Reaction Quotient, Q
 The reaction quotient, Q, has the
same form as the equilibrium
constant, Kc, except the
concentrations are not necessarily
equilibrium concentrations.
 Comparison of Q with Kc enables
the prediction of the direction the
reaction will occur to the greater
extent when a system is not at
equilibrium.

21. The Reaction Quotient, Q
 The relationship between Q and Kc
 When
 Q = Kc the system is at equilibrium
 Q > Kc reaction occurs to the left to
the greater extent
 Q < Kc reaction occurs to the right to
the greater extent

22. Determining the Direction of
Reaction
The value of Q indicates the direction in which a reaction must
proceed to reach equilibrium.
If Q < K, the reactants must increase and the products decrease;
reactants → products until equilibrium is reached.
If Q > K, the reactants must decrease and the products increase;
products → reactants until equilibrium is reached.
If Q = K, the system is at equilibrium and no further net change
takes place.

23. Figure 17.5 Reaction direction and the relative sizes of Q and
K.
Q < K
Q > K
Q = K

24. Solving Equilibrium Problems
If equilibrium quantities are given, we simply substitute these into
the expression for Kc to calculate its value.
If only some equilibrium quantities are given, we use a reaction table
to calculate them and find Kc.
A reaction table shows
• the balanced equation,
• the initial quantities of reactants and products,
• the changes in these quantities during the reaction, and
• the equilibrium quantities.

25. Figure 17.6 Steps in solving equilibrium problems.
PRELIMINARY SETTING UP
1. Write the balanced
equation.
2. Write the reaction quotient,
Q.
3. Convert all amounts into
the correct units (M or atm).
WORKING ON THE REACTION
TABLE
4. When reaction direction is not
known, compare Q with K.
5. Construct a reaction table.
 Check the sign of x, the
change in the concentration
(or pressure).

26. Figure 17.6 continued
SOLVING FOR x AND EQUILIBRIUM
QUANTITIES
6. Substitute the quantities into Q.
7. To simplify the math, assume that x
is negligible:
([A]init – x = [A]eq ≈ [A]init)
8. Solve for x.
9. Find the equilibrium quantities.
 Check to see that calculated
values give the known K.
 Check that assumption is
justified (<5% error). If not,
solve quadratic equation for x.

27. Exercises
 The equilibrium constant for the
following reaction is 49 at 450oC. If
0.22 moles of I2, 0.22 moles of H2
and 0.66 moles of HI were placed in
an evacuated 1.00L container,
determine if the system is at
equilibrium. If not, in which
direction will the reaction occur to
the greater extent to achieve
equlibrium.

28. Exercises(cont)
 H2(g) + I2(g) = 2HI(g)

31. Uses of the Equilibrium Constant
 The equilibrium constant, Kc, is 3.00
for the following reaction at a given
temperature. If 1.00 moles of SO2
and 1.00 moles of NO2 are put into
an evacuated 2.00L container and
allowed to reach equilibrium,
determine the concentration of each
compound at equilibrium.

32. Uses of the Equilibrium
Constant(cont.)
 SO2(g) + NO2(g) = SO3(g) + NO(g)

34. Uses of the Equilibrium
Constant(cont.)
 The equilibrium constant is 49 for
the following reaction at 450oC. If
1.00 mole of HI is placed in an
evacuated 1.00L container and
allowed to reach equilibrium,
determine the equilibrium
concentration of all species.
 H2(g) + I2(g) = 2HI(g)

35. Le Châtelier’s Principle
When a chemical system at equilibrium is disturbed, it reattains
equilibrium by undergoing a net reaction that reduces the effect of
the disturbance.
A system is disturbed when a change in conditions forces it
temporarily out of equilibrium.
A shift to the left is a net reaction from product to reactant.
The system responds to a disturbance by a shift in the equilibrium
position.
A shift to the right is a net reaction from reactant to product.

36. Le Châtelier’s Principle
When a chemical system at equilibrium is disturbed, it reattains
equilibrium by undergoing a net reaction that reduces the effect of
the disturbance.
A system is disturbed when a change in conditions forces it
temporarily out of equilibrium.
A shift to the left is a net reaction from product to reactant.
The system responds to a disturbance by a shift in the equilibrium
position.
A shift to the right is a net reaction from reactant to product.

37. Factors That Affect Equilibrium
 LeChatelier’s Principle
 If a change of conditions(stress) is
applied to a system in equilibrium, the
system responds in the way that best
tends to reduce the stress by reaching
a new state of equilibrium.
 Changes in concentrations,
pressure, and temperature are
considered stresses to the system.

38. Figure
17.7
The effect of a change in
concentration.

39. 1. Changes in concentration
 Consider the following reaction at
equilibrium
 H2(g) + I2(g) = 2HI(g)
 Determine the equilibrium shift if
 H2 is added
 H2 is removed
 HI is removed
 I2 is added

40. 2. Changes in volume
 If the volume of the container is
decreased for a system at
equilibrium, the concentrations of
all gases(but not liquids or solids)
will increase. If the balanced
equation has more moles of gas on
the reactant side than on the
product side, the forward reaction is
favored. An increase in volume has
the opposite effect.

41. 2. Changes in volume
 Consider the following reaction at
equilibrium at a constant
temperature:
 2 NO2(g) = N2O4(g)
 Determine the equilibrium shift if
 the volume is decreased(pressure
increased)
 the volume is increased(pressure
decreased)
 the pressure is increased at constant
volume by the addition of an inert gas

42. The Effect of a Change in Temperature
To determine the effect of a change in temperature on equilibrium,
heat is considered a component of the system.
Heat is a product in an exothermic reaction (DH°rxn < 0).
Heat is a reactant in an endothermic reaction (DH°rxn > 0).
An increase in temperature adds heat, which favors the endothermic
reaction.
A decrease in temperature removes heat, which favors the
exothermic reaction.

43. 3. Changes in temperature
 Increasing the temperature always
favors the reaction that consumes
heat, and vice-versa.
 Consider the following reaction at
equilibrium at a given temperature
 2SO2(g) + O2(g) = 2SO3(g) + 198kJ
 Determine the equilibrium shift if
 T is increased
 T is decreased

44. 4. Introduction of a catalyst
 Addition of a catalyst increases the
rate at which equilibrium is
achieved but has no effect on the
final equilibrium. The same
equilibrium is achieved but in a
shorter time.

45. Changes in the Value of Kc
 Changes in concentration, volume,
and the addition of a catalyst do not
change the value of Kc.
 Changes in temperature do effect
the value of Kc.
 Kc will increase if the forward
reaction is endothermic and
decrease if the forward reaction
is exothermic. See table 17.4.

46. Table 17.4 Effects of Various Disturbances on a System at
Equilibrium

47. Exercises
 Given the following reaction at
equilibrium in a closed container at
500oC. Indicate the direction the
equilibrium would shift(left, right,
no shift) and the effect on the value
of Kc(increase, decrease, no
change) for each of the following
changes.

48. Exercises(cont.)
 N2(g) + 3H2(g) =2NH3(g) DH=-92kJ
 Position Kc
 Increase T
 Decrease V
 Add N2
 Remove NH3
 Add a catalyst

49. Exercises
 Given the following reaction at
equilibrium in a closed container at
500oC. Indicate the direction the
equilibrium would shift(left, right,
no shift) and the effect on the value
of Kc(increase, decrease, no
change) for each of the following
changes.

50. Exercises
 H2(g) + I2(g) = 2HI (g) DH= +25kJ
 Position Kc
 Increase T
 Increase V
 Add HI
 Remove H2
 Add a catalyst

52. K and Q for hetereogeneous equilibrium
A hetereogeneous equilibrium involves reactants and/or
products in different phases.
A pure solid or liquid always has the same “concentration”, i.e., the
same number of moles per liter of solid or liquid.
The expressions for Q and K include only species whose
concentrations change as the reaction approaches equilbrium.
Pure solids and liquids are omitted from the expression for
Q or K..
For the above reaction, Q = [CO2]

53. Heterogeneous Equilibria
 involve two or more phases or states
of matter. The activities of solids and
liquids are unity.
 Write Kc expressions for each of the
lowing equilibria.
 CaCO3(s) = CaO(s) + CO2(g)
 SO2(g) + H2O(l) = H2SO3(aq)
 CaF2(s) = Ca2+(aq) + 2F-1(aq)