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Chemical Equilibrium

Chemical Equilibrium. Dynamic Equilibrium. Reversible reactions. Always reach the same equilibrium concentration whether you start with the reactants or products The forward and reverse reaction proceed at equal rates (k 1 = k -1 ) N 2 + 3 H 2 2 NH 3

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Chemical Equilibrium

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  1. Chemical Equilibrium Dynamic Equilibrium

  2. Reversible reactions • Always reach the same equilibrium concentration whether you start with the reactants or products • The forward and reverse reaction proceed at equal rates (k1 = k-1) N2 + 3 H2 2 NH3 • YouTube - Belousov-Zhabotinsky Reaction

  3. Equilibrium does not mean equal amounts at equilibrium! H2(g) + I2(g) 2 HI(g) At equilibrium: [H2] = 0.022 M, [I2] = 0.022 M, [HI] = 0.156 M • There is more HI at equilibrium (0.156 M) than there is of H2 and I2 (both at 0.022M). • The product side of this reaction is favored.

  4. The Equilibrium Constant Expression • The equilibrium constant expression relates the concentrations of the products to the reactants aA + bB cC + dD • [C]c[D]d • K = [A]a[B]b

  5. Reaction Quotient, Q • When a reaction is not at equilibrium the concentration of the of the products to reactants is called the mass action expression and is equal to the reaction quotient. • When Q = Kc the reaction is in equilibrium • [C]c[D]d • Q = [A]a[B]b

  6. Equilibrium Constant Facts • K is independent of the original concentrations. • K does depend on the temperature of the reaction. • K does not include solids or liquids because their concentration is not changing. • K’ is the reverse reaction equilibrium constant • K’ = 1/K • K for the overall reaction is equal to the product of the individual step reaction constants

  7. Equilibrium Constant Facts 7) Multiplying the Coefficients by a Factor PCl3 + Cl2 PCl5 Kc 2PCl3 + 2 Cl2 2 PCl5 Kc’ Kc’ = Kc2

  8. The Equilibrium Constant • We can calculate the equilibrium constants by combining two or more reactions for which the value of K is known. N2(g) + O2(g)  2 NO (g) K1 = 2.3 x 10-19 2 NO(g) + O2(g)  2 NO2(g) K2 = 3 x 106 __________________________ N2(g) + 2 O2(g) 2 NO2(g) K = K1 x K2 = (2.3 x 10-19)(3 x 106) = 7 x 10-13

  9. Equilibrium Law for Gases, Kp • To use Kp, everything must be a gas. • Just like Kc, Kp always has the same value (provided you don't change the temperature), irrespective of the amounts of A, B, C and D you started with. • Kp has exactly the same format as Kc, except that partial pressures are used instead of concentrations.

  10. The Haber Process N2 (g) + 3H2 (g) 2NH3(g) • The Kp expression: Kp = (PNH3)2 (PN2) (PH2)3 • Remember: PA = the partial pressure of A, which is the mole fraction of A x the total pressure PA = XA x PT

  11. Kp andKc • The relationship between Kc and Kp is derived from the Ideal Gas Law: Kp = Kc x (RT) Dn • R = ideal gas constant = 0.0821 L-atm/mol-K • T = temperature of the system in Kelvin • Dn= moles of gas products - moles of gas reactants

  12. The Value of Kc or Kp • A very large Kc or Kp indicates the reaction goes to completion, meaning the reaction is not reversible • A very small Kc or Kp indicates the forward reaction does not occur to any real extent. • Double-digit negative/positive powers of ten meets the requirement for very small/large.

  13. Le Châtelier’s Principle • If an outside influence upsets an equilibrium the system undergoes a change in a direction that counter-acts the disturbing influence, and returns the system to equilibrium.

  14. Conditions that Disturb an Equilibrium Reaction • changing the concentration of one of the components (either by adding more or removing some) • changing the pressure of the system • changing the temperature

  15. Changing the Concentration • Adding or removing a reactant or product shifts the equilibrium to consume the added substance or replace the substance removed. CoCl42- + 6 H2O Co(H2O)62+ +4Cl- • Le Chatelier Principle

  16. Changing the Temperature • Increasing the temperature shifts the equilibrium towards the endothermic direction and decreasing the temperature shifts toward the exothermic direction. 2 NO2(red-brown)  N2O4 (colorless)+heat Effect of Temperature on an Equilibrium Reaction Nitrogen dioxide

  17. Changing the Pressure • If the volume decreases the equilibrium shifts the reduce the number of molecules of gas to counter the increased pressure and visa-versa for increasing the volume. • Le Chatelier

  18. Factors that do not Change the Equilibrium of a Reaction • Catalyst- only increases the rate at which the reaction achieves equilibrium. • Inert gas added (at constant volume) – does not change the concentration of the reactants or products.

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