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CHEM 163 Chapter 17

CHEM 163 Chapter 17. Spring 2009 Instructor: Alissa Agnello aagnello@sccd.ctc.edu. What affects reaction rates?. Chemical Equilibrium. Many reactions can go forward AND backwards. If the opposite reaction can occur,. then this reaction is a reversible reaction. What is equilibrium?.

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CHEM 163 Chapter 17

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  1. CHEM 163Chapter 17 Spring 2009 Instructor: Alissa Agnello aagnello@sccd.ctc.edu

  2. What affects reaction rates?

  3. Chemical Equilibrium Many reactions can go forward AND backwards If the opposite reaction can occur, then this reaction is a reversible reaction

  4. What is equilibrium? • …in terms of reaction rates? • …in terms of reactant and product concentrations? • Has the forward reaction stopped? • Has the reverse reaction stopped?

  5. Chemical Equilibrium • Reactions continue at equal (but opposite) rates • No further changes in concentrations of reactants or products occurs

  6. Chemical Equilibrium If fwd and rev reactions are both elementary steps, how would we write their rate laws? ratefwd = kfwd [N2O4]eq raterev = krev [NO2]2eq kfwd [N2O4]eq = krev [NO2]2eq

  7. Equilibrium Constants (K)

  8. Equilibrium Constants: Small K Small K value: • Greater concentration of reactants or products? • Reaction favors reactants

  9. Equilibrium Constants: Large K Large K value: • Greater concentration of reactants or products? • Reaction favors products

  10. Equilibrium Constants Will each of the following favor reactants or products? • Reaction with Kc = 2.9 x 10-12 • Reaction with Kc = 0.001 x 105

  11. Reaction Quotient (Q) • K derived from rates • Q derived from concentrations At a given temperature, a system will always return to the same [product] : [reactant] ratio Q What if K = Q?

  12. Reaction Quotients coefficients reactants products

  13. 2-minute practice Write a reaction quotient expression for the following: Compare your answer with your neighbors!

  14. Q for Overall reaction Many reactions take place in multiple steps. • Add together the steps to get overall reaction • MultiplyK for each step to get overall K • MultiplyQ for each step to get overall Q

  15. Other situations… In a reversible reaction: When the coefficients are multiplied by a common factor (n):

  16. If solids or liquids are present… • heterogeneous equilibrium • Q and K only related to concentration that change (gases)

  17. 2-minute exercise Rearrange the ideal gas law, so that concentration is isolated on one side. What is the relationship between pressure and concentration?

  18. Qp: using partial pressures concentration pressure

  19. Shortcut for relating Qp to Qc What is the ∆n for the reaction? nreactants = 2 nproducts = 2 Qp= Qc ∆n = 0 Kp=Kc Qp= Qc (RT)∆n If ∆n ≠ 0 Kp= Kc(RT)∆n

  20. Is a reaction at equilibrium? Compare Q and K! • Q < K: • Too much reactant • Equilibrium “shifts” towards products • Q > K: • Too much product • Equilibrium “shifts” towards reactants • Q = K: • At equilibrium

  21. Equilibrium Calculations • If equilibrium concentrations and/or equilibrium constant is known: • Write Q expression, plug in concentrations, solve • At equilibrium Q = K • Solving for K when concentrations not given: • Write Q expression • Set up table including: • Initial concentrations (or pressures) • Change • Equilibrium concentrations (or pressures)

  22. Mix together graphite and carbon dioxide (P = 0.458 atm) to create carbon monoxide. Once equilibrium is reached, the pressure in the vessel (from CO2 and CO) is 0.757 atm. CO2 (g) + C (graphite) ↔ 2CO (g) 0.458 0 Initial - x + 2x Change Equilibrium 0.458 - x + 2x Ptotal = 0.458 – x + 2x = 0.757

  23. Solving for “x” To solve for x, you may need to use the quadratic formula. Set up your equation: a*x2 + b*x + c = 0 • You will end up with two values. Which value is right? • Remember: • K can’t be negative • We can’t have a higher concentration of reactant than our initial concentration.

  24. Steps for Solving • Write balanced equation • Set up table • Solve for x by… • Setting up Q expression • Setting total pressures equal to final pressure • Solve for equilibrium concentrations or pressures • (using x) • To check: • plug in solved concentration or pressures into Q expression and compare to known K value

  25. 3-minute Practice Consider the reaction: Kc = 9.30 x 10-8 at 700 °C 2 H2S (g) ↔ H2 (g) + S2 (g) 2 0.45 mole H2S is placed in a 3.0 L container. Make a table for this situation. Calculate the equilibrium concentration of H2 (g) at 700 °C

  26. Le Châtelier’s Principle • At equilibrium, concentrations of substances do not change. • If a stress is put on the reaction at equilibrium, the equilibrium will shift to relieve the stress. What changes count as stress? • Concentration • Adding or removing reactant or product • Volume (Pressure) • Temperature

  27. Effect of Concentration Changes 2 2NO2 (g) NO (g) + O2 (g) • What is the effect on the concentration of each substance? • Add NO2? • Add NO? • Add O2? • Remove NO2? • Remove NO? • Remove O2?

  28. Calculations: adding/removing substances • Make a table! 0.159 0.598 Original Equil. + 0.1 Disturbance 0.159 New Initial 0.698 Change + x - 2x New Equil. 0.159 + x 0.698 - 2x Predict direction of shift

  29. Effect of Volume Changes • Each mole of gas exerts a certain pressure • Decrease the volume… • Increase the pressure… • Shifts to side of reaction with FEWER moles of gas • 2NO2 (g) 2NO (g) + O2 (g) • What happens if we increase the volume? • Shifts to side with MORE moles • What happens if we increase the pressure? • Shifts to side with FEWER moles

  30. Effect of Volume Changes With volume changes and concentration changes, Equilibrium shifts to relieve new stress …in order to return to equilibrium

  31. Effect of Temperature Changes • If T is increased, equilibrium shifts to remove the heat • If T is decreased, equilibrium shifts to create heat • Kc changes! • Unlike for concentration and volume changes • Endothermic: • Heat + • A (g) B (g) + C (g) • Increase the temperature? • Shifts to the products (to use up heat) • Kcincreases

  32. Effect of T on K • Exothermic (∆H°< 0): increasing T decreases Kc • Endothermic (∆H°> 0): increasing T increases Kc van’t Hoff Equation: R = 8.314 J/mol∙K

  33. 3-minute Practice Consider the exothermic reaction between nitrogen gas and hydrogen gas, creating ammonia gas (NH3). Write a balanced equation for this reaction. Which direction will the equilibrium shift if: • T is increased? • Ammonia is removed? • Volume of the container is decreased? • A catalyst is used? • Hydrogen is added?

  34. Chapter 17 Homework Due: Tuesday, 4/14 #18, 23, 29, 34, 45, 55, 63, 76, 77, 91, 114

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