Chapter 15. Chemical Equilibrium
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 15.2 Achieving Chemical Equilibrium in the 15.1 The Concept of Equilibrium
Activities: Section:
Chemical Equilibrium 15.1 The Concept of Equilibrium
Animations: Section:
NO2-N2O4 Equilibrium 15.1 The Concept of Equilibrium
Temperature Dependence of Equilibrium 15.7 Le Châtelier’s Principle
Movies: Section:
Nitrogen Dioxide and Dinitrogen Tetroxide 15.1 The Concept of Equilibrium
Formation of Water 15.7 Le Châtelier’s Principle
3-D Models: Section:
Dinitrogen Tetroxide 15.1 The Concept of Equilibrium
Other Resources
Further Readings: Section:
Fritz Haber 15.2 The Equilibrium Constant
Chemical Equilibrium
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Introducing Dynamic Equilibrium as an 15.2 The Equilibrium Constant
Explanatory Model
Live Demonstrations: Section:
Equilibrium and Le Châtelier’s Principle 15.2 The Equilibrium Constant
Effect of Concentration on Equilibrium: Cobalt 15.7 Le Châtelier’s Principle
Complex
Effect of Temperature Change on Equilibrium: 15.7 Le Châtelier’s Principle
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Chapter 15. Chemical Equilibrium
Common Student Misconceptions
• Students confuse the arrows used for resonance and equilibrium.
Teaching Tips
• Many students need to see how the numerical problems in this chapter are solved.
• Students who have difficulty with some of the mathematical manipulations in this chapter should be
directed to Appendix A of the text.
Lecture Outline
15.1 The Concept of Equilibrium
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• Consider colorless frozen N2O4.
• At room temperature, it decomposes to brown NO2.
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“Chemical Equilibrium” Activity from Instructor’s Resource CD/DVD
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Figure 15.2 from Transparency Pack
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“NO2-N2O4 Equilibrium” Activity from Instructor’s Resource CD/DVD
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“NO2-N2O4 Equilibrium” Animation from Instructor’s Resource CD/DVD
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“Nitrogen Dioxide and Dinitrogen Tetroxide” Movie from Instructor’s Resource CD/DVD
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“Dinitrogen Tetroxide” 3–D Model from Instructor’s Resource CD/DVD
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“Nitrogen Dioxide” 3–D Model from Instructor’s Resource CD/DVD
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“Oxygen” 3–D Model from Instructor’s Resource CD/DVD
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• At equilibrium the concentrations of N2O4 and NO2 do not change.
• This mixture is called an equilibrium mixture.
• The equilibrium mixture results because the reaction is reversible.
• This is an example of a dynamic equilibrium.
FORWARD REFERENCES
• Equilibria involving acids and bases will be the subject of Chapters 16 and 17.
• Solubility equilibria will be discussed in Chapter 17 (section 17.4).
• The concept of equilibrium will be important throughout Chapter 19.
15.2 The Equilibrium Constant
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• Consider the Haber process:
N2(g) + 3H2(g) 2NH3(g)
• It is used for the preparation of ammonia from nitrogen and hydrogen.
• The process is carried out at high temperature and pressure.
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“Fritz Haber” from Further Readings
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“Equilibrium Constant” Activity from Instructor’s Resource CD/DVD
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“An Elementary Discussion of Chemical Equilibrium” from Further Readings
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• We can write an expression for the relationship between the concentration of the reactants and
products at equilibrium.
• This expression is based on the law of mass action.
• For a general reaction, aA + bB dD + eE
Evaluating Kc
• The value of Kc does not depend on initial concentrations of products or reactants.
• Consider the reaction:
N2O4(g) 2NO2(g)
• We generally omit the units of the equilibrium constant.
Equilibrium Constants in Terms of Pressure, Kp
• When the reactants and products are gases, we can write an equilibrium expression using partial
pressures rather than molar concentrations.
• The equilibrium constant is Kp where “p” stands for pressure.
• For the reaction: aA + bB dD + eE
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FORWARD REFERENCES
• Various equilibrium constants: Ka, Kb, Kw, Ksp will be used throughout Chapters 16 and 17,
and later in select end-of-chapter problems in Chapters 19 and 20.
• Kp and Kc will be revisited in Chapter 19 (section 19.7).
• Haber process will be further discussed in Chapter 22 (sections 22.2 and 22.7).
Equilibrium Constants and Units
• Equilibrium constants are reported without units.
• The equilibrium constant may be derived from thermodynamic measurements.
• The constants are defined in terms of activities rather than concentrations or partial pressures.
15.3 Understanding and Working with Equilibrium Constants
The Magnitude of Equilibrium Constants
• The equilibrium constant, K, is the ratio of products to reactants.
• Therefore, the larger K the more products are present at equilibrium.
• Conversely, the smaller K the more reactants are present at equilibrium.
• If K >> 1, then products dominate at equilibrium and equilibrium lies to the right.
• If K << 1, then reactants dominate at equilibrium and the equilibrium lies to the left.
The Direction of the Chemical Equation and K
• An equilibrium can be approached from either direction.
• Consider the reaction:
Relating Chemical Equation Stoichiometry and Equilibrium Constants
• It is possible to calculate the equilibrium constant for a reaction if we know the equilibrium constants
for other reactions that add up to give us the one we want.
• This is similar to using Hess’s law.
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• The equilibrium constant of a reaction in the reverse direction is the inverse of the
equilibrium constant of the reaction in the forward direction.
15.4 Heterogeneous Equilibria
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• Equilibria in which all reactants and products are present in the same phase are called homogeneous
equilibria.
• Equilibria in which one or more reactants or products are present in a different phase are called
heterogeneous equilibria.
• Consider the equilibrium established when solid lead(II) chloride dissolves in water to form a
• If a pure solid or pure liquid is involved in a heterogeneous equilibrium, its concentration is not
included in the equilibrium constant expression.
• Note: Although the concentrations of these species are not included in the equilibrium expression,
they do participate in the reaction and must be present for an equilibrium to be established!
• Other common examples of heterogeneous equilibria include:
• systems involving solids and gases.
liquid.
FORWARD REFERENCES
• Solubility equilibria will be discussed in detail in Chapter 17 (section 17.4).
• Exclusion of pure liquids and solids from equilibrium expressions will be revisited in Chapter
19.
15.5 Calculating Equilibrium Constants
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• Proceed as follows:
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• Deduce the equilibrium concentrations of all species.
• Use these to calculate the value of the equilibrium constant.
FORWARD REFERENCES
• Methods of finding equilibrium constants from thermodynamic or electrochemical will be
discussed in Chapters 19 and 20, respectively.
15.6 Applications of Equilibrium Constants
Predicting the Direction of Reaction
• For a general reaction: aA + bB dD + eE
• We define Q, the reaction quotient, as:
Calculating Equilibrium Concentrations
• The same steps used to calculate equilibrium constants are used to calculate equilibrium
concentrations.
• Generally, we do not have a number for the change in concentration.
• Therefore, we need to assume that x mol/L of a species is produced (or used).
• The equilibrium concentrations are given as algebraic expressions.
FORWARD REFERENCES
• Q will be used in Chapter 19 (section 19.7) to determine Gibb’s free energy change at
nonstandard conditions.
• Q will be used in Chapter 20 (section 20.6) to determine cell potentials at nonstandard
conditions (Nernst equation).
15.7 Le Châtelier’s Principle
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• Consider the Haber process:
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Change in Reactant or Product Concentration
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• If a chemical system is at equilibrium and we add or remove a product or reactant, the reaction will
shift so as to reestablish equilibrium.
• For example, consider the Haber process again:
N2(g) + 3H2(g) 2NH3(g)
• If H2 is added while the system is at equilibrium, Q < K.
• The system must respond to counteract the added H2 (Le Châtelier’s principle).
• That is, the system must consume the H2 and produce products until a new equilibrium is
established.
Effects of Volume and Pressure Changes
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• Consider a system at equilibrium.
• If the equilibrium involves gaseous products or reactants, the concentration of these species will be
changed if we change the volume of the container.
• For example, if we decrease the volume of the container, the partial pressures of each gaseous
species will increase.
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“Amounts Tables as a Diagnostic Tool for Flawed Stoichiometric Reasoning” from Further Readings
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“Calculating Equilibrium Concentrations by Iteration: Recycle Your Approximations” from Further
Readings
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• The instant the pressure increases, the concentration of both gases increases and the system is not
at equilibrium.
Effect of Temperature Changes
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• The equilibrium constant is temperature dependent.
• How will a change in temperature alter a system at equilibrium?
• It depends on the particular reaction.
• For example, consider the endothermic reaction:
Co(H2O)62+(aq) + 4Cl–(aq) CoCl42–(aq) + 6H2O(l) ∆H > 0
• Co(H2O)62+ is pale pink and CoCl42– is a deep blue.
• At room temperature, an equilibrium mixture (light purple) is placed in a beaker of warm water.
The Effect of Catalysts
• A catalyst lowers the activation energy barrier for the reaction.
• Therefore, a catalyst will decrease the amount of time needed to reach equilibrium.
• A catalyst does not affect the composition of the equilibrium mixture.
FORWARD REFERENCES
• Le Châtelier’s principle versus percent ionization will be mentioned in Chapter 16 (section
16.6).
• Le Châtelier’s principle will be brought up in the common ion effect in Chapter 17 (sections
17.1 and 17.3).
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“Effect of Temperature Change on Equilibrium: Cobalt Complexes” from Live Demonstrations
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“Equilibrium in the Gas Phase” from Live Demonstrations
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Further Readings:
1. Martin R. Feldman and Monica L. Tarver, “Fritz Haber,” J. Chem. Educ., Vol. 60, 1983, 463–464.
2. Carl W. David, “An Elementary Discussion of Chemical Equilibrium,” J. Chem. Educ., Vol. 65, 1988,
407–409.
6. Jan H. Van Driel, Wobbe de Vos, and Nico Verloop, “Introducing Dynamic Equilibrium as an
Explanatory Model,” J. Chem. Educ., Vol. 76, 1999, 559–561.
7. Penelope A. Huddle, Margie W. White, and Fiona Rogers, “Simulations for Teaching Chemical
Equilibrium,” J. Chem. Educ., Vol. 77, 2000, 920–926.
11. William R. Smith and Ronald W. Missen, “Chemical Equilibrium and Polynomial Equations: Beware
of Roots,” J. Chem. Educ., Vol. 66, 1989, 489–490.
12. E. Weltin, “Calculating Equilibrium Concentrations by Iteration: Recycle Your Approximations,” J.
Chem. Educ., Vol. 72, 1995, 36–38.
Live Demonstrations:
1. Lee R. Summerlin and James L. Ealy, Jr., “Equilibrium in the Gas Phase,” Chemical Demonstrations,
A Sourcebook for Teachers (Washington: American Chemical Society, 1988), pp. 85–86. Color changes
in a mixture of NO2 and N2O4 as a sealed tube of gas is heated or cooled are used to demonstrate Le
Châtelier’s principle.
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Chemical Society, 1988), pp. 79–80. An equilibrium system containing the dehydrated-hydrated cobalt
complex is shifted in response to changes in temperature.
3. Lee. R. Summerlin, and James. L. Ealy, Jr., “Equilibrium and Le Châtelier’s Principle,” Chemical
Demonstrations, A Sourcebook for Teachers, Volume 1 (Washington: American Chemical Society, 1988),
pp. 77–78. An overhead projector demonstration of the effect of reactant concentration on equilibrium.