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Chapter 15

Chapter 15. Chemical equilibria. Closed vial of NO 2 : NO 2 (g) + NO 2 (g)  N 2 O 4 (g) brown colorless Brown color becomes less intense, then, after some time period, color change stops. Concentrations of products and reactants remain constant with time (equilibrium)

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Chapter 15

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  1. Chapter 15 Chemical equilibria

  2. Closed vial of NO2: NO2(g) + NO2(g)  N2O4(g) brown colorless Brown color becomes less intense, then, after some time period, color change stops. Concentrations of products and reactants remain constant with time (equilibrium) Applies to all reactions Chemical Equilibrium

  3. H2O(g) + CO(g)  H2(g) + CO2(g) Start off with closed flask of CO and H2O. Equilibrium sets up (chemical reaction occurring). What happens to amounts of H2O(g) + CO(g) ? What happens to amounts of H2(g) + CO2(g) ? Reaction still occurring past dotted line (although concentrations don’t change): Dynamic Equilibrium Dynamic Equilibrium

  4. Dynamic Equilibrium H2O(g) + CO(g)  H2(g) + CO2(g)

  5. Law of mass action jA + kB  lC + mD K = [C]l [D]m Equilibrium Expression [A]j [B]k Concentrations of species at equilibrium K = equilibrium constant H2O(g) + CO(g)  H2(g) + CO2(g) 2NO2(g)  N2O4(g) Equilibrium Expression and K

  6. 3H2(g) + N2(g)  2NH3(g) Value for K is always the same for a particular reaction, Value of Equilibrium Constant

  7. Equilibrium position (does equilibrium lie towards products or reactants?); is this important (Synthesis of Aspirin) Large K K1 Small K Magnitude of K

  8. Writing equilibrium expressions Always predicted from the balancedequation: 4NH3(g) + 7O2(g)  4NO2(g) + 6H2O(g) I2(g) + H2(g)  2HI(g) C4H10(g) + O2(g)  CO2(g) + H2O(g)

  9. Calculating values of K

  10. 3H2(g) + N2(g)  2NH3(g) K = [NH3]2Equilibrium constant found in terms of [H2]3 [N2] concentrations of species (Kc) K can also be found in terms of partial pressures(Kp) PV = nRT Kp = PNH32 (PH23)(PN2) Relationship between Kp and Kc: Kp = Kc(RT)ng Eq. laws for gaseous reactions

  11. More than one phase exists in reaction mixture Thermal decomposition of CaCO3(s)  CaO(s) + CO2(g) Concentrations of pure solids and pure liquids are always constant (thus, can be removed from eq. expression) Heterogeneous Equilibria

  12. Write equilibrium expressions (in terms of both K and Kp for each of the following): CaO(s) + SO2(g)  CaSO3(s) Decomposition of solid phosphorous pentachloride to liquid phosphorous trichloride and chlorine gas Deep blue solid copper(II)sulfate pentahydrate is heated to drive off water vapor to form white solid copper(II)sulfate Heterogeneous Equilibria

  13. In which direction will a particular reaction shift to reach equilibrium? 3H2(g) + N2(g)  2NH3(g) If [NH3]0 = 0, shift to right to achieve equilibrium. If [H2]0 or [N2]0 = 0, shift to left to achieve equilibrium. If initial concentrations of all three species are nonzero, which way will shift occur to achieve equilibrium? More difficult to predict. Use Reaction Quotient, Q Q is obtained by applying law of Mass Action to initial concentrations of species involved. Reaction Quotient, Q

  14. 3H2(g) + N2(g)  2NH3(g) Q = [NH3]o2 [H2]o3 [N2]o Then, to determine in which direction a system will shift to reach equilibrium, compare values of Q and K: Q = K System is at equilibrium Q > K Ratio of initial conc. of products to initial conc. of reactants is too large. System shifts to the left to reach equilibrium. Q < K Ratio of initial conc. of products to initial conc. of reactants is too small. System shifts to the right to reach equilibrium Reaction Quotient, Q

  15. Calculating Equilibrium Concentrations • ‘ICE’ Tables

  16. ‘If an outside influence upsets an equilibrium, the system undergoes a change in the direction that counteracts the disturbing influence, and, if possible, returns the system to equilibrium.’ Outside influences? Adding / removing a reactant / product Changing volume/pressure of gaseous reactions Changing T Le Chatelier’s Principle

  17. Eq. shifts in direction that will partially consume a reactant or product added Eq. shifts in direction that will partially replace a reactant or product removed 3H2(g) + N2(g)  2NH3(g Add some N2 to above equilibrium; which way will it shift to re-establish equilibrium? Adding/removing a reactant or product

  18. Reducing volume (what happens to the pressure?) – eq. shifts to side with smaller # of gas molecules 3H2(g) + N2(g)  2NH3(g Increasing volume? What would happen here? Changing volume/pressure

  19. 2NO2(g)  N2O4(g) brown colorless Changing volume/pressure

  20. Increasing T shifts an equilibrium in direction that produces an endothermic change (need to know energy involved in reaction to predict which direction this is) 3H2(g) + N2(g)  2NH3(g) H= -46.19 kJ/mol Exothermic or endothermic reaction? Increasing T: Decreasing T: Changing T

  21. Is this reaction exo- or endothermic? Changing T 2NO2(g)  N2O4(g) brown colorless

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