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Chapter 13

Chapter 13. Chemical Equilibrium. Chemical Equilibrium. Chemical equilibrium occurs in chemical reactions that are reversible. In a reaction such as: CH 4 (g) + H 2 O(g)  CO(g) + 3H 2 (g) The reaction can proceed in both directions CO(g) + 3H 2 (g)  CH 4 (g) + H 2 O(g).

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Chapter 13

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  1. Chapter 13 Chemical Equilibrium

  2. Chemical Equilibrium Chemical equilibrium occurs in chemical reactions that are reversible. In a reaction such as: CH4(g) + H2O(g)  CO(g) + 3H2 (g) The reaction can proceed in both directions CO(g) + 3H2 (g)  CH4(g) + H2O(g)

  3. An Equilibrium System CH4(g) + H2O(g)  CO(g) + 3H2 (g) • After some of the products are created the products begin to react to form the reactants • At equilibrium there is no net change in the concentrations of the reactants and products • The concentrations do not change but they are not necessarily equal

  4. Chemical Equilibrium • CH4(g) + H2O(g) <=> CO(g) + 3H2 (g) • At the beginning of the reaction, the rate that the forward direction is higher • As the reactants decrease the rate in the forward direction slows • As the products form, the rate in the reverse direction increases • When the two rates are the same equilibrium is achieved

  5. Dynamic Equilibrium An equilibrium is Dynamic CH4(g) + H2O(g)  CO(g) + 3H2 (g) The amount of products and the reactants are constant. (Note: The concentrations are not necessarily equal but constant. Both reactions are still occurring, but at the same rate)

  6. The Equilibrium Constant aA + bB  cC + dD The upper case letters are the molar concentrations of the reactants and products. The lower case letters are the coefficients that balance the equation.

  7. The Equilibrium Constant aA + bB  cC + dD

  8. Equilibrium Constant Calculations Example N2(g) + 3H2 (g)  2NH3 (g) At equilibrium, a one-liter container has 1.60 moles NH3, .800 moles N2, and 1.20 moles of H2. What is the equilibrium constant?

  9. Equilibrium Constant Calculations At equilibrium, a one-liter container has 1.60 moles NH3, .800 moles N2, and 1.20 moles of H2. What is the equilibrium constant?

  10. • If Kc > 103 products predominate over reactants. If Kc is very large, the reaction proceeds nearly to completion. • • If Kc < 10-3 reactants predominate over products. If is very small, the reaction proceeds hardly at all. • • If Kc is in the range 10-3 to 103 appreciable concentrations of both reactants and products are present.

  11. Equilibrium Constant Kc, Kp • For gases Kc and Kp are used. • KP same format as Kc except pressures used instead of concentrations.

  12. Equilibrium Constant Kc, Kp E.g. Write out the equilibrium expression for KP using the reaction below: N2(g) + 3H2(g)  2NH3(g) KP = ? E.g.3 Determine the equilibrium constant (KP) for the formation of one mole of ammonia if at 500K, PNH3 = 0.15 atm, PN2 = 1.2 atm. and PH2 = 0.81 atm.

  13. Relationship between Kc and Kp • Relationship between concentration and pressure obtained from the ideal gas law. • Recall PV = nRT or • Substitute for P in equilibrium expression. Consider the reaction: aA + bB  cC + dD • Use this relationship to relate KP and Kc

  14. Reaction quotient • The equilibrium constant is a constant ratio only when the system is in equilibrium. • If the system it not at equilibrium the ratio is known as a Reaction Quotient • If the reaction quotient is equal to the equilibrium constant then the system is at equilibrium

  15. Using Equilibrium Constants for other calculations If a solution is not at equilibrium the ratio of the right side over the left is called a reaction quotient. aA + bB  cC + dD

  16. Equilibrium Constants and calculations If Qc > Keq , the product side is too high and the equilibrium will shift to the left to restore equilibrium If Qc < Keq , the product side is too low and the equilibrium will shift to the right to restore equilibrium If Qc = Keq, System is at equilibrium

  17. Equilibrium Calculations –Using I. C. E. Models Equilibrium constants and concentrations can often be deduced by carefully examining data about initial and equilibrium concentrations Initial Change Equilbrium

  18. Equilibrium CalculationsICE Model problem 1 Hydrogen and iodine are in equilibrium with Hydrogen iodide to this reaction: H2 + I2 2HI Suppose that 1.5 mole of H2 and 1.2 mole of I2 are placed in a 1.0 dm3 container. At equilibrium it was found that there were 0.4 mole of HI. Calculate the equilibrium concentrations of [H2] and [I2] and the equilibrium constant.

  19. Equilibrium CalculationsICE Model Problem 1 Solution Hydrogen and iodine are in equilibrium with Hydrogen iodide to this reaction: H2 + I2 2HI Suppose that 1.5 mole of H2 and 1.2 mole of I2 are placed in a 1.0 dm3 container. At equilibrium it was found that there were 0.4 mole of HI. Calculate the equilibrium concentrations of [H2] and [I2] and the equilibrium constant. I C E Since 2x = 0.4, x = 0.2 [H2 ] 1.5 -x 1.5- x [H2 ] = 1.5 – 0.2 = 1.3 [ I2 ] 1.2 -x 1.2 –x [I2 ] = 1.2 – 0.2 = 1.0 [HI ] 0 +2x 0.4 Keq = [HI]2 . = (0.4)2 = 0.123 [H2 ] [I2 ] (1.3)(1.0)

  20. Equilibrium CalculationsICE Model Problem 2 Sulfur dioxide reacts with oxygen to produce sulfur trioxide according to this reaction: 2 SO2 + O2 2SO3 Suppose that 1.4 mole of SO2 and 0.8 mole of SO3 are placed in a 1.0 dm3 container. At equilibrium it was found that there were 0.6 dm3 of SO3. Calculate the equilibrium concentrations of [SO2] and [O2] and the equilibrium constant.

  21. Equilibrium CalculationsICE Model Problem 2 Solution Sulfur dioxide reacts with oxygen to produce sulfur trioxide according to this reaction: 2 SO2 + O2 2SO3 Suppose that 1.4 mole of SO2 and 0.8 mole of SO3 are placed in a 1.0 dm3 container. At equilibrium it was found that there were 0.6 dm3 of SO3. Calculate the equilibrium concentrations of [SO2] and [O2] and the equilibrium constant. I C E Since 2x = 0.6, x = 0.3 [SO2 ] 1.4 -2x 1.4-2x [SO2 ] =1.4 – 2( 0.3) = 0.8 [O2 ] 0.8 -x 0.8 –x [O2 ] = 0.8 – 0.3 = 0.5 [SO3] 0 +2x 0.6 Keq = [SO3]2 . = (0.6)2 = 0.281 [SO2 ]2 [O2 ] (0.8)2(0.5)

  22. Le Chatelier’s Principle • Le Chatelier's Principle states: When a system in chemical equilibrium is disturbed by a change of temperature, pressure, or a concentration, the system shifts in equilibrium composition in a way that tends to counteract this change of variable. • A change imposed on an equilibrium system is called a stress • The equilibrium always responds in such a way so as to counteract the stress

  23. Le Chatelier’s Principle • Types of stresses • Change in concentration of one or more reactants or products • Change in temperature • Change in pressure • Addition of a catalyst

  24. Effect of a Change in Temperature • An increase in the temperature causes the equilibrium to shift in the direction of the endothermic reaction • N2 (g) + 3 H2 (g)  2NH3 (g) DH =-92 kJ mol-1 • Since DH is negative the endothermic reaction is the reverse direction. An increase in temperature causes the reaction to shift to the left, resulting in an increase in N2 and H2 and a decrease in NH3

  25. Effect of a Change in Pressure • Pressure affects only gases in an equilibrium • PV = nRT • An increase in pressure causes the equilibrium to shift in the direction that has the fewer number of moles • N2 (g) + 3 H2 (g)  2NH3 (g) DH =-92 kJ mol-1 • An increase in pressure results in a an decrease in N2 and H2 and an increase in NH3

  26. Effect of a Change in one of the reactants or products • The equilibrium responds in such a way so as to diminish the increase • Substances on the same side of the arrow respond in opposite directions. • Substances on the opposite side of the arrow move in the same direction • N2 (g) + 3 H2 (g)  2NH3 (g) • An increase in [N2 ] results in a decrease in N2 and H2 and an increase in NH3

  27. Effect of a Catalyst • Catalysts affect both the forward and reverse directions equally • A catalyst does not change the concentrations but reduces the time required for the system to come to equilibrium

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