1. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 1 Introductory Chemistry: A Foundation FIFTH EDITION by Steven S. Zumdahl University of Illinois

2. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 2 Equilibrium Chapter 16

3. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 3 Collision Theory of Kinetics Kinetics is the study of the factors that effect the speed of a reaction and the mechanism by which a reaction proceeds. In order for a reaction to take place, the reacting molecules must collide into each other. Once molecules collide they may react together or they may not, depending on two factors - ¬Whether the collision has enough energy to "break the bonds holding reactant molecules together"; ­Whether the reacting molecules collide in the proper orientation for new bonds to form.

4. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 4 Effective Collisions Collisions in which these two conditions are met (and therefore the reaction occurs) are called effective collisions. The higher the frequency of effective collisions the faster the reaction rate. When two molecules have an effective collision, a temporary, high energy (unstable) chemical species is formed - called an activated complex –Not a true molecule because its bonds aren’t complete

5. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 5 Activated Complex The difference in potential energy between the reactant molecules and the activated complex is called the activation energy, E a The larger the activation energy, the slower the reaction The energy to overcome the activation energy comes from the kinetic energy of the collision being converted into potential energy, or from energy available in the environment, i.e. heat. Different reactions have different activated complexes and therefore different activation energies

6. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 6 Factors Affecting Reaction Rate The kind of molecules and what condition the reactants are in. –Small molecules tend to react faster than large molecules; –Gases tend to react faster than liquids which react faster than solids; –Powdered solids more reactive than “blocks” More surface area for contact with other reactants –Certain types of chemicals are more reactive than others –Ions react faster than molecules Increasing temperature always increases reaction rate

7. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 7 The larger the concentration of reactant molecules, the faster the reaction will go. –Increases the frequency of reactant molecule contact Concentration of gases depends on the partial pressure of the gas; –Higher pressure = Higher concencentration Concentration of solutions depends on the solute to solution ratio (molarity). Factors Affecting Reaction Rate

8. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 8 Catalysts are substances that effect the speed of a reaction without being consumed. Most catalysts are used to speed up a reaction. Homogeneous = present in same phase Heterogeneous = present in different phase –Molecule gets adsorbed onto catalyst active site Enzymes Catalysts work by providing a pathway for the reaction with a lower activation energy Factors Affecting Reaction Rate

9. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 9 Molecular Interpretation of Factors Affecting Rate Increasing the temperature raises the kinetic energy of the reactant molecules causing them to collide more frequently and have more collisions with sufficient energy to form the activated complex The larger the concentration of reactant molecules, the more frequently they collide and the larger the number of effective collisions will be Therefore the speed will increase.

10. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 10 Molecular Interpretation of Factors Affecting Rate Catalysts work by providing an alternative pathway for the reaction with a lower activation energy Lowering the activation energy means more molecules have enough kinetic energy so that when they collide they can form the activated complex The result is the reaction goes faster

11. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 11 Reaction Dynamics If the products of a reaction are removed from the system as they are made, then a chemical reaction will proceed until the limiting reactants are used up. However, if the products are allowed to accumulate; they will start reacting together to form the original reactants - called the reverse reaction.

12. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 12 Reaction Dynamics The forward reaction slows down as the amounts of reactants decreases because the reactant concentrations are decreasing At the same time the reverse reaction speeds up as the concentration of the products increases. Eventually the forward reaction is using reactants and making products as fast as the reverse reaction is using products and making reactants. This is called chemical equilibrium. –rate forward = rate reverse

13. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 13 Chemical Equilibrium Equilibrium only occurs in a closed system!! When a system reaches equilibrium, the amounts of reactants and products in the system stays constant –the forward and reverse reactions still continue, but because they go at the same rate the amounts of materials don't change. There is a mathematical relationship between the amounts of reactants and products at equilibrium –no matter how much reactants or products you start with.

14. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 14 Equilibrium Constant Law Of Chemical Equilibrium In this expression, K is a number called the equilibrium constant. Do not include solids or liquids, only solutions and gases For a reaction, the value of K for a reaction depends on the temperature K is independent of the amounts of reactants and products you start with. aA + bB  cC + dD = [C] c [D] d [A] a [B] b K eq

15. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 15 Position of Equilibrium The relative concentrations of reactants and products when a reaction reaches equilibrium is called the position of equilibrium Different initial amounts of reactants (and or products) will result in different equilibrium concentrations but the same equilibrium constant

16. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 16 Position of Equilibrium If K is large then there will be a larger concentration of products at equilibrium than of reactants; we say the position of equilibrium favors the products. If K is small then there will be a larger concentration of reactants at equilibrium than of products; we say the position of equilibrium favors the reactants. The position of equilibrium is not affected by adding a catalyst.

17. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 17 Example – Determine the value of the Equilibrium Constant for the Reaction 2 SO 2 + O 2  2 SO 3 ¬Determine the Equilibrium Expression ­Plug the equilibrium concentrations into to Equilibrium Expression ®Solve the Equation 3.503.00SO 3 1.251.50O2O2 2.00SO 2 [Equilibrium][Initial]Chemical

18. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 18 Le Ch âtelier’s Principle Le Châtelier's Principle guides us in predicting the effect various changes have on the position of equilibrium it says that when a change is imposed on a system at equilibrium, the position of equilibrium will shift in the direction that will reduce the effect of that change

19. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 19 Concentration Changes and Le Châtelier’s Principle The position of equilibrium can be affected without changing the equilibrium constant. Adding a reactant will decrease the amounts of the other reactants and increase the amount of the products until a new position of equilibrium is found Position shifts toward the products –has the same K Removing a reactant will increase the amounts of the other reactants and decrease the amounts of the products. –Position shifts toward the reactants Removing a product can allow us to drive a reaction to completion!

20. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 20 Changing the pressure of one gas is like changing its concentration –Has the same effect as changing the concentration on the position of equilibrium Increasing the pressure on the system causes the position of equilibrium to shift toward the side of the reaction with the fewer gas molecules Decreasing the volume of the system increases its pressure –Reduces the pressure by reducing the number of gas molecules –Opposite effect happens if the system pressure is decreased Changing Pressure and Le Châtelier’s Principle

21. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 21 The equilibrium constant will change if the temperature changes Exothermic reactions release heat, Endothermic reactions absorb heat. For exothermic reactions, heating the system will decrease K –Think of heat as a product of the reaction –Therefore shift the position of equilibrium toward the reactant side For endothermic reactions, heating the system will increase K –Think of heat as a reactant –The position of equilibrium will shift toward the products Cooling an exothermic or endothermic reaction will have the opposite effects on K and equilibrium position Changing Temperature and Le Châtelier’s Principle

22. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 22 ¬Determine the Equilibrium Expression ­Plug the equilibrium concentrations and Equilibrium Constant into the Equilibrium Expression ®Solve the Equation ?3.00SO 3 1.251.50O2O2 2.00SO 2 [Equilibrium][Initial]Chemical Example – If the value of the Equilibrium Constant for the Reaction 2 SO 2 + O 2  2 SO 3 is 4.36, Determine the Equilibrium Concentration of SO 3

23. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 23 Solubility & Solubility Product Even “insoluble” salts dissolve somewhat in water –insoluble = less than 0.1 g per 100 g H 2 O The solubility of insoluble salts is described in terms of equilibrium between undissolved solid and aqueous ions produced A n X m (s)  n A + (aq) + m Y - (aq) Equilibrium constant called solubility product K sp = [A + ] n [Y - ] m If undissolved solid in equilibrium with the solution, the solution is saturated Larger K = More Soluble –for salts that produce same the number of ions

24. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 24 Example ¬Determine the balanced equation for the dissociation of the salt AgI(s)  Ag + (aq) + I - (aq) ­Determine the expression for the solubility product –Same as the Equilibrium Constant Expression K sp = [Ag + ][I - ] Calculate the solubility of AgI in water at 25°C if the value of K sp = 1.5 x 10 -16

25. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 25 ®Define the concentrations of dissolved ions in terms of x AgI(s)  Ag + (aq) + I - (aq) Stoichiometry tells us that we get 1 mole of Ag + and 1 mol I - for each mole of AgI dissolved Let x = [Ag + ], then [I - ] = x Example Calculate the solubility of AgI in water at 25°C if the value of K sp = 1.5 x 10 -16

26. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 26 ¯Plug the ion concentrations into the expression for the solubility product and solve for K sp [Ag + ] = [I - ] = x [Ag + ] = 1.2 x 10 -8 mol/L = [AgI] The solubility of AgI = 1.2 x 10 -8 M Example Calculate the solubility of AgI in water at 25°C if the value of K sp = 1.5 x 10 -16