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Equilibrium - PowerPoint PPT Presentation


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Slide 1. Equilibrium. Slide 2. Equilibrium Constant Kp. Because gas pressures are easily measured, equilibrium equations for gas-phase reactions are often written as partial pressures rather than molarity concentrations.

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slide2

Slide 2

Equilibrium Constant Kp

  • Because gas pressures are easily measured, equilibrium equations for gas-phase reactions are often written as partial pressures rather than molarity concentrations.
  • For example the equilibrium equation for the decomposition of N2O4 can be written
  • N2O4 2NO2
  • Kp = (PNO2)2
  • (PN2O4)
slide3

Slide 3

** Just a little note with the last slide…The equilibrium equations for Kp and Kc

have the same form except that the expression for Kp contains partial pressures instead of molar concentrations.

** For Kp as with Kc the units are omitted.

slide4

Slide 4

It can be shown that the values of Kp and Kc for a general gas-phase reaction are related by the equation:

For the reaction aA + bB cC + dD

Kp = Kc(RT)∆n

R = the gas constant 0.0821 L∙atm/K∙mol

T = absolute temperature in kelvin

∆n = is the sum of the coefficients of the gaseous reactions : for the above reaction (c+d) –(a+b)

example
Example

Slide 5

N2O4(g) 2NO2(g)

For the decomposition of 1 mol of N2O4 to 2 mol NO2, ∆n = 2-1=1 and Kp=Kc(RT)1

H2(g) + I2(g) 2HI

For the reaction of 1 mol of hydrogen with 1 mol of iodine to give 2 mol of hydrogen iodine. ∆n= 2-(1+1) = 0 and

Kp=Kc(RT)0 or Kp = Kc

slide6

Slide 6

Homogeneous Equilibria

  • So far every example dealt with reactants and products where all were in the same phase.
  • We can use K in terms of either concentration or pressure.
  • Units depend on reaction.
slide7

Slide 7

Heterogeneous Equilibria

  • If the reaction involves pure solids or pure liquids the concentration of the solid or the liquid doesn’t change.
  • As long as they are not used up they we can leave them out of the equilibrium expression.
slide8

Slide 8

Example

  • H2(g) + I2(s) 2HI(g)
  • K = [HI]2 [H2][I2]
  • But the concentration of I2 does not change.
  • K[I2]= [HI]2 =K’ [H2]
slide9

Slide 9

The Reaction Quotient

  • Tells you the direction the reaction will go to reach equilibrium
  • Calculated the same as the equilibrium constant, but for a system not at equilibrium
  • Q = [Products]coefficient [Reactants] coefficient
  • Compare value to equilibrium constant
slide10

Slide 10

What Q tells us

  • If Q<K
    • Not enough products
    • Shift to right
  • If Q>K
    • Too many products
    • Shift to left
  • If Q=K system is at equilibrium
slide11

Slide 11

Example

  • for the reaction
  • 2NOCl(g) 2NO(g) + Cl2(g)
  • K = 1.55 x 10-5 M at 35ºC
  • In an experiment 0.10 mol NOCl, 0.0010 mol NO(g) and 0.00010 mol Cl2 are mixed in 2.0 L flask.
  • Which direction will the reaction proceed to reach equilibrium?