Chapter 15
1 / 38

Chapter 15 Principles of Chemical Equilibrium - PowerPoint PPT Presentation

  • Uploaded on

Chapter 15 Principles of Chemical Equilibrium. Contents in Chapter 15. 15-1 Dynamic Equilibrium 15-2 The Equilibrium Constant Expression 15-3 Relationships Involving Equilibrium Constants 15-4 The Magnitude of an Equilibrium Constant

I am the owner, or an agent authorized to act on behalf of the owner, of the copyrighted work described.
Download Presentation

PowerPoint Slideshow about 'Chapter 15 Principles of Chemical Equilibrium' - rigg

An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.

- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

Chapter 15Principles of Chemical Equilibrium

Contents in Chapter 15

15-1 Dynamic Equilibrium

15-2 The Equilibrium Constant Expression

15-3 Relationships Involving Equilibrium Constants

15-4 The Magnitude of an Equilibrium Constant

15-5 The Reaction Quotient, Q: Predicting the Direction of Net Change

15-6 Altering Equilibrium Conditions: Le Châtelier’s Principle

15-7 Equilibrium Calculations: Some Illustrative Examples

15-1 Dynamic Equilibrium

  • Dynamic equilibrium: A condition where forward and reverse processes proceed at equal rates and no further net change occurs.

    • Chemical equilibrium: At the condition that the concentrations of reactants and products no longer changes with time.

    • Physical equilibrium: The equilibrium between two phases, that is according to physical processes.

AgI(s) Ag+(aq) + I–(aq)

  • Example of Chemical equilibrium







I2(H2O) I2(CCl4)

  • Example of physical equilibrium

H2O top layer

CCl4 bottom layer

15-2 The Equilibrium Constant Expression

  • Established empirically (trial and error)

15-2 (Continuous)

  • Derived theoretically (kinetic view of equilibrium)

    • In chemical kinetics view, the equilibrium constant of a reaction can be expressed as a ratio of the rate constants of “forward” and “reverse” reactions:

      Kc = (forward rate)/(reverse rate) = kf/kr

    • The expression of chemical equilibrium according to the coefficients of the balance equation, independent to the kinetic mechanism.

Example: Proof a net reaction A + 2B  AB2 has same equilibrium constant expression although the reaction mechanisms are different:


  • If the mechanism with single elementary step:

  • A + 2B AB2K (equilibrium constant)


(B) If the mechanism with multi-elementary step, such as:

Step 1 2B B2K' (equilibrium constant)

Step 2 A + B2 AB2K'' (equilibrium constant)

Net A + 2B  AB2 K (equilibrium constant)






(A) ratef = kf[A][B]2; rater = kr[AB2]

at equilibrium, ratef = rater


15-2 (Continuous)

  • The Equilibrium Constant and Activities

  • Activity: The effective concentration for obtaining the (true) thermodynamic constants, the activity of a species is obtained as the product of an activity coefficient and the ratio of the stoichiometic concentration (or pressure) to that of a reference state.

X: activity coefficient, [X]: actual concentration,

Co: concentration of reference state.

15-2 (Continuous)

  • A General Expression for K

15-3 Relationships Involving Equilibrium Constants

  • Relationship of K to the Balanced Chemical Equation

    • Reverse an equation, invert the value of K.

    • Multiply the coefficients in a balanced equation by a common factor, raise the equilibrium constant to the corresponding power.

    • The balanced chemical equation must be written when citing a value for K.


  • Combining Equilibrium Constant Expressions

    • Individual equations are combined (that is, added), their equilibrium constants are multiplied to obtain the equilibrium constant for the overall


  • Equilibria Involving Gases

  • Concentration equilibrium constant (KC) using molar concentrations ([ ]), unit in M:

  • Pressure equilibrium constant (KP), using partial pressure (p), unit in bar:

  • Conversion between KC and KP:

    KP = KC(RT)Δn(R=0.08314 bar L K–1 mol–1)

    Δn = mol of gaseous Products - mol of gaseous Reactants

    In this case: Δn = (g+h) – (a+b)

    * P = (n/V)RT


  • Equilibria Involving Pure Liquids and Solids

    • The equilibrium constant expression does not include terms of pure solid and liquid phases.


C(s) + H2O(g) CO(g) + H2(g)

CaCO3(s) CaO(s) + CO2(g)

Kc = [CO2] KP = pCO2 = KC(RT)

H2O(l) H2O(g)

Kc = [H2O(g)] KP = pH2O = KC(RT)

15-4 The Magnitude of an Equilibrium Constant

  • A very large numerical value of K, if K > 1010, signifies that a reaction goes to completion.

  • A very small numerical value of K, if K < 10–10, signifies that the forward reaction occurs only to a very slight extent.

  • Although a reaction may be thermodynamically favored, it may kinetically controlled.

  • Thermodynamic equilibrium constant expression uses dimensionless unit known as activities.

15-5 The Reaction Quotient, Q: Predicting the Direction of Net Change

  • For an instantaneous conditions, the expression having the same form as K is called the reaction quotient, Q.

    • The Q value is not constant for a reaction, but is useful for predicting the direction in which a net change must occur to establish equilibrium.

    • If Q < K, a net change occurs in the forward direction, that is, from left to right.

    • If Q > K, a net change proceeds in the reverse direction, that is, from right to left.

    • If Q = K, the equilibrium is reached.

15-6 Altering Equilibrium Conditions: Le Châtelier’s Principle

  • Le Châtelier’s Principle: When a change (concentration, temperature, pressure, or volume) is imposed on a system at equilibrium, the system responds by attaining a new equilibrium condition that offsets the impact of the change.

(Continuous) Principle

  • Effect of Changing the Amounts of Reacting Species on Equilibrium

    • At equilibrium, Q = K.

    • If one of the reactants is added or one of the products is removed, resulted in Q < K, the concentrations of all species must change to make Q once again = K.

    • If Q < K, the reaction is pushed (shift) to the right.

(Continuous) Principle

  • Effect of Changes in Pressure or Volume on Equilibrium

    • Add an inert gas to the constant-volume reaction mixture: total pressure increased, but the partial pressure of the reactants and products do not change, no effect on the equilibrium condition.

    • Change the pressure by changing the volume of the system: For example, if the volume is reduced, i.e., concentration increased, an equilibrium shifts in the direction producing the smaller number of moles of gas. (If no change in the number of moles of gas in a reaction, no effect on an equilibrium).

(Continuous) Principle

  • Effect of Temperature on Equilibrium

  • The equilibrium constant (K) depend on temperature:

  • For endothermic reaction (△Ho > 0), increase T, increase K, shift right for restoring the equilibrium.

  • For exothermic reaction (△Ho < 0), increase T, decrease K, shift left for restoring the equilibrium.

(Continuous) Principle

  • Effect of a Catalyst on Equilibrium

    • The role of a catalyst is to change the mechanism of a reaction to lower activation energy.

    • Catalyst dosenot effectthe equilibrium.

15-7 Equilibrium Calculations: PrincipleSome Illustrative Examples

(Continuous) Principle


(Continuous) Principle

(Continuous) Principle

(Continuous) Principle