International Baccalaureate Chemistry - PowerPoint PPT Presentation

International Baccalaureate Chemistry

Topic 7 – Chemical Equilibrium

• In general, equilibrium is the state in which the rate of the forward process/reaction equals the rate of the reverse process/reaction.

• Liquid - Vapor Equilibria
• The following equation represents Bromine liquid with Bromine gas in equilibrium.

Br(l)→ Br(g)

• At equilibrium, the rate of evaporation equals the rate of condensation.
• NOTE: This does not mean that the number of molecules in each state is the same!

• Solute – Solution Equilibria
• The following equation represents solid sodium chloride in equilibrium with its ions in a saturated solution:

NaCl(s) ↔ Na+(aq) + Cl-(aq)

• At equilibrium, the rate of dissolving equals the rate of crystalization.
• NOTE: This does not mean that the number of molecules in each state is the same!

• Nature of the reactants
• Type of Liquid (in liquid – vapor equilibria)
• Type of Solid (in solute – solution equilibria)
• Temperature
• In liquid – vapor equilibriamore particles will be found in the gas phase with an increase in temperature.
• In a solute-solution equilibria, more particles will be found in the dissolved phase with an increase in temperature.

In a physical equilibrium, the

system must be closed!

• Many chemical reactions are reversible and never go to completion.
• Consider the following reaction: (Haber Process)

N2 + 3H2↔ 2NH3

• If a certain quantity of each reactant is placed in a closed container, initially, N2 and H2 are the only species present in the container.
• N2 and H2 will begin to react at their maximum rate because their concentrations are at a maximum.

Forward Reaction → maximum rate

Reverse Reaction → no rate

• As time passes, the forward reaction will start to decrease because [N2] and [H2] are decreasing (used up).
• As this happens, an opposing change begins to occur.
• NH3 begins to decompose into N2 and H2 through the reverse reaction.
• The rate of the reverse reaction will steadily increase (as the [NH3] is increasing).
• Eventually, the forward reaction rate becomes equal to the reverse reaction rate.
• The system is said to be in a state of Dynamic Equilibrium.

• Equilibrium can be approached from both directions.
• The rate of the forward reaction equals the rate of the reverse reaction and the concentrations of all reactants and products remains constant (NOT EQUAL!)
• The system must be closed.
• Macroscopic properties remain constant (although microscopically there is lots going on!)

• For all chemicals, Kc = [Products]

[Reactants]

Equilibrium Constant

Equilibrium Expression

For the reaction aA + bB ↔ cC + dD, The Equilibrium Constant expression is:

Kc = [Products] = [C]c [D]d

[Reactants] [A]a [B]b

Where: a,b,c,d are the coefficients from the chemical equation

• The coefficients from the equation become a power in the equilibrium expression.
• Only gasses and aqueous solutions are involved in the equilibrium expression since solids and liquids cannot be expressed as concentrations.
• NOTE:
• Solids cannot be compressed so their density cannot be changed.
• Therefore, the molar concentrations cannot be changed.
• Liquids cannot be compressed either so that they have a constant density and molar concentration.
• If there is another liquid present that can dilute the first liquid, then the liquid is not “pure” and can have its concentration changed by dilution.

• N2(g) + 3H2(g)↔ 2NH3(g)
• Cu(s) + 2AgNO3(aq) ↔ 2Ag(s) + Cu(NO3)2(aq)
• P4(s) + 5O2(g) ↔ P4O10(s)
• Br2(l) + H2(g) ↔ 2HBr(g)
• CH3COCH3(l) + Cl2(g) ↔ CH3COCH2Cl(l) + HCl(g)

1. Kc = [NH3]2 2. Kc = [Cu(NO3)2] 3. Kc = 1

[N2] [H2]3 [AgNO3]2 [O2]5

4. Kc = [HBr]2 5. Kc = [HCl]

[H2] [Cl2]

• Whenever Kc is large (Kc > > 1):
• Products are favored and the reaction goes almost to completion.
• Kc = [Products] Kc = [BIG #] > > 1

[Reactants] [small #]

• Whenever Kc is small (Kc < < 1):
• Reactants are favored and the reaction hardly proceeds.
• Kc = [Products]Kc = [small #] < < 1

[Reactants] [BIG #]

• Note: The value of Kcis affected only by temperature.

• Changes will occur in a closed system if it has been disturbed or stressed.
• To explain these changes, we use

Le Chatelier’s Principle which states:

• “Whenever a stress is applied to a system at equilibrium, it will shift to relieve that stress applied.”

• Change in Concentration
• Change in Temperature
• Change in Volume or Pressure
• Addition of a Catalyst

Affect Equil. Position

Does Not Affect Equil. Position

(does not cause a shift)

• Only affects substances that are (g) or (aq)
• Or if two liquids are present (l)
• Example 1:

Fe3+(aq) + SCN-(aq) ↔ FeSCN2+(aq)

• If we add more [Fe3+] to the reaction mixture, the reaction will shift forward to use up what is added.
• The result will result in a decrease of [SCN-] and an increase in [FeSCN2+] product produced.
• Conclusion: Increasing the concentration of a reactant favors the forward reaction.
• Eventually, a new equilibrium point will be reached and no further change in concentration will occur.
• The same rules apply if a concentration is decreased.
• To offset the stress, more of what is removed must be produced.

• Example 2: N2(g)+ H2(g) ↔ NH3(g)
• If the [NH3] is decreased, the system will shift forward in an attempt to replace ammonia that was removed.
• As a result, N2 and H2 decrease.
• Example 3: Identify the shift (forward or reverse) that occurs for the following reaction and indicate the results of the shift and identify the remaining reactants.

2A(s) + 3B(g) ↔ 2E(g) + 4D(g)

• [B] is increased
• [E] is increased
• Amount of A is increased
• [D] is decreased

• At equilibrium, the temperature is constant.
• If the reaction vessel is cooled, the reaction will shift to produce heat.
• If the reaction vessel is heated, the reaction will shift to use up heat.

• Example 1: Endothermic Reaction
• Co(H2O)62+(aq) + 4Cl- + energy ↔ CoCl42-(aq) + 6H2O(l)
• If we disturb the system by increasing the temperature, both the forward and reverse reaction rates increase, however, in this reaction the forward will increase more than the reverse. Why?
• Forward reaction – uses heat (cools)
• Reverse reaction – gives off heat (warms)
• Results:
• [Co(H2O)62+] and [Cl-] Decrease
• [CoCl42-] increases
• The moles of water will also increase

• Example 2: Exothermic Reaction
• 3H2(g) + N2(g)↔ 2NH3(g) + 91 kJ
• If we disturb the system by increasing the temperature, both the forward and reverse reaction rates increase, however, in this reaction the reverse reaction will increase more than the forward.
• Forward reaction – gives off heat (warms)
• Reverse reaction – uses heat (cools)
• Results
• [H2] and [N2] increase
• [NH3] Decreases

• Generally, a temperature increase will
• favor endothermic reactions
• Not favor exothermic reactions.

• Only gasses are affected by pressure/volume changes.
• Recall:
• If the volume of a container is decreased (pressure increased), all concentrations increase.
• This will increase both the forward and reverse reactions, but the reaction will try to offset the increase in pressure .
• To decrease the pressure, the reaction will favor the reaction which produces fewer gas particles.

• Example 1:
• 3H2(g) + N2(g) ↔ 2NH3(g) + 91 kJ
• If the volume of the container is decreased, pressure is increased.
• 4 moles of gas (3H2 and 1N2) vs. 2 moles of gas (2NH3)
• Results
• Forward reaction is favored since the reaction is producing fewer gas molecules and alleviates the pressure.
• [H2] and [N2] decrease
• [NH3] increase
• What happens to the same reaction if the volume is increased?

• Example 2:
• H2(g) + I2(g)↔ 2HI(g)
• If you were to either increase or decrease the pressure of the system, it would not make a difference in the reaction rate as it cannot get rid of stress in either direction (2 moles of gas on either side)
• Example 3:
• NH3(g) + HCl(g) ↔ NH4Cl(g)
• If the pressure is decreased, which reaction is favored and what are the results?

• Example 4:
• 2A(g) + B(g) ↔ 3D(s) + C(g)
• Stress: Increase the pressure by decreasing the volume:
• Reaction favored: _______________
• [A] will ________________________
• [B] will ________________________
• [C] will ________________________
• Amount of [D] will _______________

• A catalyst will increase both forward and reverse rate equally.
• Concentrations of all substances remain constant.
• Catalysts do not affect the position of equilibrium.

• Using the resources available to you, describe and explain the application of equilibrium and kinetics concepts to the Haber process and Contact process.
• Start with the IB text and supplement with other texts and/or internet to make a complete set of notes.