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Equilibrium

Equilibrium. Chemistry 30. Equilibrium in Chemical Systems. Reactions do not always go forward to completion (reactants do not all react) Instead reach equilibrium: concentration of reactants and products becomes constant (but not equal) Use an arrow in the equation to show equilibrium. Ex)

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Equilibrium

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  1. Equilibrium Chemistry 30

  2. Equilibrium in Chemical Systems • Reactions do not always go forward to completion (reactants do not all react) • Instead reach equilibrium: concentration of reactants and products becomes constant (but not equal) • Use an arrow in the equation to show equilibrium

  3. Ex) Ex) Colorless Violet Colorless

  4. Dynamic equilibrium – a balance between the forward and reverse reactions occurring at the same rate with constant macroscopic properties (temp, ph, color etc).

  5. Although rates are the same, concentrations of reactants and products are different at equilibrium • equilibrium is reached when concentrations stop changing

  6. Demonstration of simulated chemical equilibrium

  7. Two Types of Equilibrium • Phase change equilibrium • Solid in a saturated solution ex) sugar and sugar water • Gas above a liquid ex) pop • Vapour above a solid ex) mothballs in a dresser

  8. Chemical equilibrium • Homogeneous: reactants and products are in the same phase or state Ex) C2H5OH(l) + CH3COOH(l)   CH3COOC2H5(l) + H2O(l) • Heterogeneous: reactants and products are in different phases or states Ex) CaCO(s) CaO(s) + CO2(g)

  9. 4 Conditions for Equilibrium • Forward rate and reverse rate of reaction are the same • Macroscopic properties are constant • Closed system • Equilibrium can be approached from either direction

  10. Classes of Chemical Reaction Criteria

  11. Equilibrium Constant • Ratio of the product and reactant concentrations is a constant value. • This value is called the equilibrium constant Kc

  12. Finding Kc • For the reaction: aA + bBcC+ dD Kc= [C]c[D]d = products [A]a[B]b reactants *Only substances whose concentrations or pressures change can be included in the equilibrium expression. Liquids and solids are not included.

  13. If K>1 the reaction favors products • If K<1 the reaction favors reactants • If K=1 there are equal concentrations of reactants and products

  14. Calculations with Equilibrium Constants • Case 1 – Given all the amounts of species at equilibrium, calculate the equilibrium constant. • Case 2 – Given the initial and final conditions, calculate the equilibrium constant.

  15. In each case, you will need: • The balanced chemical equation for the system at equilibrium • The equilibrium constant expression (remember to include only those that are solutions or gases, not liquids or solids)

  16. Case 1 Examples • Calculate the Kc for the following system at a constant temperature: 2NO(g) + O2(g) 2NO2(g) [NO]: 1.0 mol/L [O2]: 6.0 mol/L [NO2]: 2.0 mol/L

  17. For the reaction CO(g) + Cl2(g) COCl2(g) at a particular temperature, Kc= 5.1 x 105. At equilibrium, there are 0.30 mol of Cl2(g) and 0.072 mol of COCl2(g) in a 2.0L container. What is the equilibrium concentration of CO(g)?

  18. Case 2 Examples • Use the ICE tables • I= initial • C=change • E= equilibrium • Need to use molar ratio as well

  19. For the reaction given below, 3.00 mol of A and 4.00 mol of B are placed in a 5.00 L container. 2A(g) + B(g) C(g) At equilibrium, the concentration of A is 0.40 mol/L. Determine the value of Kc.

  20. Initially 2.0 mol of N2 and 4.0 mol of H2 were added to a 1.0 L container and the following reaction occurred. 3H2(g)+ N2(g) 2NH3(g) The equilibrium concentration of NH3 is 0.68 mol/L. What is the Kc?

  21. Le Chatelier’s Principle • when a system at equilibrium is disturbed by a change in a property, the system adjusts in a way that opposes this change. • Change can occur to: • Concentration • Pressure or volume • Temperature (this will also change Kc) • Add a Catalyst

  22. The application of Le Chatelier’s Principle involves three stages: • The initial state of equilibrium in the system • Stress exerted on the system by some change in property (temp, conc, pressure) that creates dis-equilibrium • Shift towards reactants or products, to re-establish a new state of equilibrium • Le Chatelier’s Principle can provide a method of predicting the reaction’s response to the change. • http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/lechv17.swf

  23. Concentration Changes • The addition of more reactant or the removal of a product will increase the yield of the product, shifting the equilibrium towards the right (products) A + B C • If the opposite occurs, the shift will be towards the left (reactants).

  24. Concentration change can occur by adding something that makes a precipitate. Ex) AgNO3(s) Ag+(aq) + NO3-(aq) • If you add NaCl, the Cl-(aq) reacts with Ag+ (aq) to form AgNO3(s) • Response is to shift to the right AgNO3(s) Ag+(aq) + NO3-(aq)

  25. 2. Pressure / Volume Change • Gases only • If pressure decreases, response is to increase pressure by shifting to the side with more moles of gas Ex) Decrease pressure and : A(g) + 2B(g) 1C (g)

  26. In a gaseous system, decreasing the volume of the container increases the pressure. • Adding more gas increases the pressure • Pressure change has no effect on equilibrium if moles are the same on both side.

  27. 3. Temperature Change • The energy in a chemical reaction is treated as if it were a reactant or product. Endothermic: reactants + energy products Exothermic:reactants products + energy • Energy can be added or removed by heating or cooling the system. Equilibrium shifts to minimize the change. • Will change Kc

  28. Ex) 2 SO3(g) + 97 kJ 2 SO2(g) + O2(g) • If you heat it up (the stress), response is to cool down and equilibrium shifts right. Kc increases • If you cool it down (the stress), response is to heat up and equilibrium shifts left. Kc decreases.

  29. 4. Catalyst • Decreases time to reach equilibrium • Lowers activation energy by the same amount in both the forward and reverse directions • Does not effect final equilibrium position

  30. Summary

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