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Equilibrium

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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Equilibrium

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  1. Slide 1 Equilibrium

  2. 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)

  3. 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.

  4. 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)

  5. 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

  6. 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.

  7. 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.

  8. 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]

  9. 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

  10. 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

  11. 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?

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