1 / 92

Chemical Equilibrium Chapter 16

Chemical Equilibrium Chapter 16. Hein and Arena . Version 1.1. Chapter Outline. 16.1 Reversible Reactions. 16.8 Effect of Catalysts on Equilibrium. 16.2 Rates of Reaction. 16.9 Equilibrium Constants. 16.3 Chemical Equilibrium. 16.10 Ion Product Constant for Water.

jacie
Download Presentation

Chemical Equilibrium Chapter 16

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. Chemical EquilibriumChapter 16 Hein and Arena Version 1.1

  2. Chapter Outline 16.1 Reversible Reactions 16.8Effect of Catalysts onEquilibrium 16.2 Rates of Reaction 16.9Equilibrium Constants 16.3 Chemical Equilibrium 16.10 Ion Product Constant forWater 16.4 Le Chatelier’s Principle 16.11Ionization Constants 16.5Effect of Concentration onEquilibrium 16.12 Solubility Product Constant 16.6Effect of Volume onEquilibrium 16.13 Acid-Base Properties ofSalts 16.7Effect of Temperature onEquilibrium 16.14 Buffer Solutions:The Control of pH

  3. Reversible Reactions

  4. reversible reaction A chemical reaction in which the products formed react to produce the original reactants.

  5. cooling 2NO2(g)→ N2O4 (g) heating N2O4(g)→ 2NO2 (g) The reaction between NO2 andN2O4 is reversible. N2O4 is formed N2O4 decomposes when heated forming NO2

  6. → 2NO2(g)N2O4 (g) → reaction to the right reaction to the left

  7. Rates of Reaction

  8. chemical kinetics The study of reaction rates and reaction mechanisms.

  9. The rate of a reaction is variable. It depends on: • concentrations of the reacting species • reaction temperature • presence or absence of catalysts • the nature of the reactants

  10. Forward reaction A + B → C + D Reverse reaction C + D → A + B The concentration of A and B decreases with time lowering the rate of the forward reaction. The concentration of C and D increases with time increasing the rate of the reverse reaction. 16.2

  11. Chemical Equilibrium

  12. equilibrium: a dynamic state in which two or more opposing processes are taking place at the same time and at the same rate. chemical equilibrium: the state in which the rate of the forward reaction equals the rate of the reverse reaction in a chemical change. At equilibrium the concentrations of the products and the reactants are not changing.

  13. NaCl(s) Na+(aq) + Cl-(aq) → → A saturated salt solution is in equilibrium with solid salt. salt crystalsare dissolving Na+ and Cl-are crystallizing At equilibrium the rate of salt dissolution equals the rate of salt crystallization.

  14. Le Chatelier’s Principle

  15. This generalization, known as LeChatelier’s Principle, states In 1888, the French chemist Henri LeChatelier set forth a far-reaching generalization on the behavior of equilibrium systems. If a stress or strain is applied to a system in equilibrium, the system will respond in such a way as to relieve that stress and restore equilibrium under a new set of conditions.

  16. Effect of Concentration on Equilibrium

  17. For most reactions the rate of reaction increases as reactant concentrations increase. • The manner in which the rate of reaction changes with concentration must be determined experimentally.

  18. An equilibrium is disturbed when the concentration of one or more of its components is changed. As a result, the concentration of all species will change and a new equilibrium mixture will be established.

  19. → → A + BC + D The system is at equilibrium results in C and D being produced faster than they are used. results in A and B being used faster than they are produced. increases the rate of the forward reaction Increasing the concentration of B

  20. → → A + BC + D The system is again at equilibrium In the new equilibrium concentration of A has decreased concentrations of B, C and D have increased After enough time has passed, the rates of the forward and reverse reactions become equal.

  21. Percent Yield

  22. The forward reaction is 79% complete at equilibrium. H2(g) +I2(g) 2HI(aq) 1 mol 0.21 mol 0 mol 0.21 mol 0 mol 1 mol 0 mol 2 mol 1.58 mol → → Initial Concentrations Final Concentrations in theAbsence of Equilibrium Equilibrium Concentrations At equilibrium the rate of HI formation equals the rate of HI decomposition. H2 + I2 combine to form HI HI decomposes to form H2 + I2 700 K

  23. Comparison of Equilibria

  24. Equilibrium shifts to left decrease Cl2 concentration decrease H3O+ concentration increase H2O concentration increase Cl- concentration increase HOCl concentration Equilibrium shifts to left Cl2(aq) +2H2O(l) HOCl(aq) + H3O+(aq) + Cl-(aq) → Equilibrium shifts to left Equilibrium shifts to right Equilibrium shifts to right → Effect of Concentration Changeson the Chlorine Water Equilibrium

  25. Equilibrium shifts to left NaC2H3O2(aq) → Na+(aq) + C2H3O2(aq) HC2H3O2(aq) +H2O(l) H3O+(aq) + C2H3O2(aq) 1 L 0.100 M HC2H3O2 1 L 0.100 M HC2H3O2 1 L 0.100 M HC2H3O2 → → Equilibrium pH = 2.87 Equilibrium pH = 4.74 Equilibrium pH = 5.05 Effect of C2H3O2Concentration Changes on pH Add 0.100 mol NaC2H3O2 Add 0.200 mol NaC2H3O2

  26. Effect of Volume on Equilibrium

  27. Changes in volume significantly affect the reaction rate only when one or more of the reactants or products is a gas and the reaction is run in a closed container. • The effect of decreasing the volume is to increase the concentrations of any gaseous reactants or products.

  28. Equilibrium shifts to left → → Decrease Volume increases CO2 concentration CaCO3(s) CaO(s) + CO2(g)

  29. → Equilibrium shifts to right → Increase Volume decreases CO2 concentration CaCO3(s) CaO(s) + CO2(g)

  30. In a system composed entirely of gases, a decrease in the volume of the container will cause the reaction and the equilibrium to shift to the side that contains the smallest number of molecules.

  31. N2(g) + 3H2(g) 2NH3(g) 1 mol 3 mol 2 mol 6.02 x 1023 molecules 1.81 x 1024 molecules 1.20 x 1024 molecules 2.41 x 1024 molecules → → Equilibrium shifts to the right towards fewer molecules. Decrease Volume

  32. N2(g) + O2(g) 2NO(g) 1 mol 1 mol 2 mol 6.02 x 1023 molecules 6.02 x 1023 molecules 1.20 x 1024 molecules 1.20 x 1024 molecules → → Decrease Volume Equilibrium does not shift. The number of molecules is the same on both sides of the equation.

  33. Effect of Temperature on Equilibrium

  34. When the temperature of a system is raised, the rate of reaction increases. The rate of the reaction that absorbs heat is increased to a greater extent, and the equilibrium shifts to favor that reaction. In a reversible reaction, the rates of both the forward and the reverse reactions are increased by an increase in temperature.

  35. At 1000oC moles CO2moles CO C(s) + CO2(g) + heat 2CO(g) → Equilibrium shifts to right → Heat may be treated as a reactant in endothermic reactions. At room temperature very little CO forms.

  36. Effect of Catalystson Equilibrium

  37. A catalyst is a substance that influences the rate of a reaction and can be recovered essentially unchanged at the end of the reaction. A catalyst does not shift the equilibrium of a reaction. It affects only the speed at which the equilibrium is reached.

  38. Energy Diagram for an Exothermic Reaction Activation energy: the minimum energy required for a reaction to occur. A catalyst speeds up a reaction by lowering the activation energy. A catalyst does not change the energy of a reaction. 16.3

  39. 2KClO3(s) → 2KCl + 3O2(l) The laboratory preparation of oxygen uses manganese dioxide as a catalyst to increase the rate of the reaction. AlCl3 PCl3(l)+ S(s) → PSCl3(l) Very little thiophosphoryl chloride is formed in the absence of a catalyst because the reaction is so slow. In the presence of an aluminum chloride catalyst the reaction is complete in a few seconds. MnO2 Δ

  40. Equilibrium Constants

  41. At equilibrium the rates of the forward and reverse reactions are equal, and the concentrations of the reactants and products are constant.

  42. The equilibrium constant (Keq) is a value representing the unchanging concentrations of the reactants and the products in a chemical reaction at equilibrium.

  43. aA + bB cC + dD → → For the general reaction at a given temperature

  44. 3H2 + N2 2NH3 → → For the reaction

  45. 4NH3 + 3O2 2N2+ 6H2O → → For the reaction

  46. The magnitude of an equilibrium constant indicates the extent to which the forward and reverse reactions take place. 3H2 + I2 2HI3 At equilibrium more product than reactant exists. At equilibrium more reactant than product exists. COCl2 CO+ Cl2 → → → →

  47. When the molar concentrations of all species in an equilibrium reaction are known, the Keq can be calculated by substituting the concentrations into the equilibrium constant expression.

  48. Calculate the Keq for the following reaction based on concentrations of PCl5 = 0.030 mol/L, PCl3 = 0.97 mol/L and Cl2 = 0.97 mol/L at 300oC. PCl5(g) PCI3(g) + Cl2(g) → →

  49. Ion Products Constant for Water

  50. H2O + H2O H3O+ + OH- H2O H+ + OH- → → → → Water autoionizes to a slight degree. or more simply At equilibrium [H+] = [OH-] = 1.00 x 10-7 mol/L The water equilibrium constant, Kw, is called the ion product constant for water. Kw = [H+][OH-] = 1.00 x 10-14 T = 25oC Kw = (1 x 10-7)(1 x 10-7) = 1.00 x 10-14

More Related