international baccalaureate chemistry l.
Skip this Video
Loading SlideShow in 5 Seconds..
International Baccalaureate Chemistry PowerPoint Presentation
Download Presentation
International Baccalaureate Chemistry

Loading in 2 Seconds...

play fullscreen
1 / 28

International Baccalaureate Chemistry - PowerPoint PPT Presentation

  • Uploaded on

International Baccalaureate Chemistry. Topic 7 – Chemical Equilibrium. Equilibrium. In general, equilibrium is the state in which the rate of the forward process/reaction equals the rate of the reverse process/reaction. Physical Equilibria. Liquid - Vapor Equilibria

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 'International Baccalaureate Chemistry' - dawn

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
international baccalaureate chemistry

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.
physical equilibria
Physical Equilibria
  • 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!
physical equilibria4
Physical Equilibria
  • 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!
what determines the equilibrium point of a substance
What Determines the Equilibrium Point of a Substance?
  • 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!

reversible reactions
Reversible Reactions
  • 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

n 2 h 2 nh 3
N2 + H2 NH3
  • 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.
characteristics of a system in a state of equilibrium
Characteristics of a System in a State of 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!)
the equilibrium constant expression
The Equilibrium Constant Expression
  • For all chemicals, Kc = [Products]


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.
practice write k c expressions for the following reactions
Practice: Write Kc Expressions for the Following Reactions
  • 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]

k c values
Kc Values
  • 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 in equilibrium
Changes in Equilibrium
  • 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.”
types of stresses that can be placed on a system at equilibrium
Types of StressesThat Can Be Placed On a System at Equilibrium
  • 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)

change in concentration
Change in Concentration
  • 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.
change in concentration17
Change in Concentration
  • 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
changes in temperature
Changes in Temperature
  • 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.
changes in temperature19
Changes in Temperature
  • 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
changes in temperature20
Changes in Temperature
  • 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
changes in temperature21
Changes in Temperature
  • Generally, a temperature increase will
    • favor endothermic reactions
    • Not favor exothermic reactions.
change in volume or pressure
Change in Volume or Pressure
  • 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.
change in volume or pressure23
Change in Volume or Pressure
  • 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?
change in volume or pressure24
Change in Volume or Pressure
  • 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?
change in volume or pressure25
Change in Volume or Pressure
  • 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 _______________
addition of a catalyst
Addition of a Catalyst
  • A catalyst will increase both forward and reverse rate equally.
  • Concentrations of all substances remain constant.
  • Catalysts do not affect the position of equilibrium.
haber process contact process
Haber Process/Contact Process
  • 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.