1 / 29

Unit 6

Unit 6. Equilibrium. What is equilibrium? . Equilibrium is NOT when all things are equal. Equilibrium is signaled by no net change in the concentrations of reactants or products. . Important notes about equilibrium.

willis
Download Presentation

Unit 6

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Unit 6 Equilibrium

  2. What is equilibrium? • Equilibrium is NOT when all things are equal. • Equilibrium is signaled by no net change in the concentrations of reactants or products.

  3. Important notes about equilibrium • At equilibrium, the concentrations of reactants and products no longer change with time. • For equilibrium to occur, neither reactants nor products can escape from the system. • At equilibrium a particular ratio or concentration terms equals a constant.

  4. The Equilibrium Constant • Consider the reaction: • aA + bBcC + dD • At equilibrium the equilibrium constant (keq) is calculated as…

  5. Examples • Write the equilibrium constant expressions for the following equations: • 2 O3(g)  3 O2(g) • 2 NO(g) + Cl2(g)  2 NOCl(g) • Ag+(aq) + 2 NH3(aq)  Ag(NH3)2+(aq)

  6. Calculating Keq

  7. Equilibrium Constant in Terms of Pressure kp • When all of the reactants and products of a chemical reaction are gasses we can write the equilibrium constant expression in terms of each components partial pressure.

  8. Keqvskp • keq and kp are often times numerically different even for the same reaction. • We can however use one of the two to calculate the other.

  9. Example • In the synthesis of ammonia from nitrogen and hydrogen, N2(g) + 3 H2(g)  2 NH3(g) • keq = 9.60 at 300oC. Calculate kp for this reaction at 300oC.

  10. Working With Equilibrium Constants • keq values for different reactions can range from very large numbers to very small numbers. • Consider the reaction: • If keq >> 1 the equilibrium lies to the right and products predominate. • If keq << 1 the equilibrium lies to the left and the reactants predominate.

  11. The Direction of the Chemical Equation and keq • Equilibrium can be approached from either direction. • Consider:

  12. Multi-Step Reactions • We have already seen cases where the overall reaction is actually a series of other reactions. • Consider the following two reactions:

  13. Heterogeneous Equilibria • A heterogeneous equilibrium is one where not all of the chemical species are in the same physical state. • Example: • Whenever a pure liquid or a pure solid is involved in a heterogeneous equilibrium, its concentration is not included in the keq expression.

  14. Example • Write the keq expression for the following reactions: • CO2(g) + H2(g) CO(g) + H2O(l) • SnO2(s) + 2 CO(g) Sn(s) + 2 CO2(g)

  15. Calculating Equilibrium Constants • The Haber process is a reaction that was first put to use by Fritz Haber as a way to produce ammonia. • One problem he had was that this reaction was an equilibrium reaction.

  16. Example • A chemists performs the Haber reaction at 472oC. The equilibrium mixture of the gases is found to contain 7.38 atm H2, 2.46 atm N2, and 0.166 atm of NH3. Calculate kp for the Haber process at this temperature.

  17. Often times we do not know the equilibrium concentrations of all of the chemical species in a reaction. • But if we know the equilibrium concentration of at least one species we can use stoichiometry to find the others (and eventually keq or kp) • Follow these steps: • Tabulate all the known initial and equilibrium concentrations. • Use these to calculate the change in concentration of that chemical species. • Use stoichiometry to calculate the change in the other chemical species. • Use the initial concentrations with the change in concentration to calculate the equilibrium concentration.

  18. Example • A closed system initially containing 1.0 x 10-3 M H2 and 2.0 x 10-3 M I2 at 448oC is allowed to reach equilibrium. Analysis of the equilibrium mixture shows that [HI] = 1.87 x 10-3 M. Calculate keq at 448oC.

  19. Applications of Equilibrium Constants • We have seen that the magnitude of keq tells us to which side the equilibrium lies. • We can also use keq to tell us which way a reaction will proceed if it isn’t already at equilibrium. • If Q = keq the reaction is at equilibrium. • If Q > keq the concentration of products is too large and the reaction will proceed to the left. • If Q < keq the concentration of reactants is too large and the reaction will proceed to the right.

  20. Example • At 448oC the equilibrium for the reaction: • H2(g) + I2(g)  2 HI(g) is 50.5. Predict the direction the reaction will proceed if we start with 0.02 mol of HI, 0.01 mol of H2 and 0.03 mol of I2 in a 2.00-L container.

  21. Calculating Equilibrium Concentrations • For the Haber process at 500oC the equilibrium constant is 1.45 x 10-5. An analysis of an equilibrium mixture of the reaction shows that the partial pressures of the reactants are H2 = 0.928 atm, N2 = 0.432 atm. Calculate the partial pressure of NH3.

  22. Calculating Equilibrium Concentrations using Initial Concentrations • Many times we will not know the equilibrium concentrations of at least one species. • Example: • A 1.0 L flask is filled with 1.0 mol H2 and 2.0 mol I2. The keq for this reaction at 448oC is 50.5. What are the equilibrium concentrations of H2, I2, and HI?

  23. Le Chatelier’s Principle • The Chatelier’s Principle states that: • If a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position to counteract the change. • In other words, keq , must remain constant.

  24. Changes in Reactant or Product Concentration. • N2(g) + H2(g)  2 NH3(g)

  25. Effects of Volume and Pressure Changes • If the volume of a gaseous system is decreased what would happen to the pressure of the system? • Reducing the volume of a gaseous equilibrium will cause the system to shift in the direction that decreases the total moles of gas. • Consider: • N2O4(g)  2 NO2(g)

  26. The Effect of Temperature Change • Changes in concentrations, volume, or pressure cause equilibrium shifts without changing the value of keq. • However changing the temperature of an equilibrium system will change the numeric value of keq. • When the temperature of an equilibrium system is increase, the system reacts as if we added a reactant to the endothermic reaction. The equilibrium will shift a such to consume that reactant.

  27. How keq changes with Temperature • If the forward reaction of an equilibrium is endothermic increasing the temperature increases the value of k. • If the forward reaction is exothermic increasing the temperature decreases the value of k.

  28. Example • Consider the equilibrium: N2O4(g)  2 NO2(g) In which direction will the equilibrium shift when (a) N2O4 is added. (b) NO2 is removed. (c) The total pressure of the system is increased. (d) The volume is increased. (e) The temperature is decreased.

More Related