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Chapter 7

Chapter 7. Chemical Reactions: Energy, Rates, and Equilibrium. Energy and Chemical Bonds. Chapter 6 Kept a careful accounting of atoms as they rearranged themselves Reactions also involve a transfer of energy. Energy and Chemical Bonds. Two fundamental kinds of energy.

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Chapter 7

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  1. Chapter 7 Chemical Reactions: Energy, Rates, and Equilibrium

  2. Energy and Chemical Bonds • Chapter 6 • Kept a careful accounting of atoms as they rearranged themselves • Reactions also involve a transfer of energy

  3. Energy and Chemical Bonds • Two fundamental kinds of energy. • Potential energy is stored energy. • Kinetic energy is the energy of motion. • Law of Conservation of Energy • Energy can be converted from one kind to another but never destroyed

  4. Energy and Chemical Bonds • A chemical • Potential - attractive forces in an ionic compound or sharing of electrons covalent compound • Kinetic – (often in form of heat) occurs when bonds are broken and particles allowed to move

  5. Heat Changes during Chemical Reactions • Bond dissociation energy: The amount of energy that must be supplied to break a bond and separate the atoms in an isolated gaseous molecule.

  6. Heat Changes during Chemical Reactions • Bond breakage requires energy to be added to the system (+ energy) • Bond formation gives off energy as the bonds form (- energy)

  7. Heat Changes during Chemical Reactions • Heat of reaction: (Enthalpy) • Represented by DH • is the difference between the energy absorbed in breaking bonds and that released in forming bonds • Endothermic: • More energy is required than released. • DH is positive • Exothermic: • More energy is released than required • DH is negative

  8. Exothermic and Endothermic Reactions

  9. Problem • Br2 (l) → Br2 (g) ΔH = 7.4 kcal/mol • Is this reaction endothermic or exothermic • Is this reaction spontaneous with respect to enthalpy? • 2C8H18 + 25O2 → 16CO2 + 18H2O + 239.5 kcal • Is this reaction endothermic or exothermic? • What is the sign of ΔH?

  10. Why do Chemical Reactions Occur? Free Energy • Events that lead to the system having less energy are said to be spontaneous with respect to enthalpy • Exothermic reactions are spontaneous • Why would endothermic reactions occur? • Free Energy (ΔG) • Enthalpy – ΔH – heat of reaction • Entropy (S)

  11. Entropy • Entropy – measures the spreading out of energy – universe moves toward disorder • Entropy favored system is one that goes from a concentrated area of energy to the energy being more spread out • ΔS is positive • Unfavorable process involves concentrating the energy into less area • ΔS is negative

  12. Why do Chemical Reactions Occur? Free Energy

  13. Problem • Identify each of the following as entropy favored or disfavored. For each state the sign of the ΔS. • Assembling a jig-saw puzzle • I2 (g) + 3F2 (g) → 2 IF3 (g) • A precipitate forming when two solutions are mixed • Demolition of a building • CS2(g) + 4H2(g) →  CH4(g) + 2H2S(g) • 2HBr(g)  → H2(g) + Br2(g)

  14. Why do Chemical Reactions Occur? Free Energy • Free Energy

  15. Why do Chemical Reactions Occur? Free Energy

  16. Problem • H2 (g) + Br2 (l) → 2 HBr (g) • Is this reaction spontaneous with respect to entropy? • If the ΔH = -17.4 kcal/mol is the reaction spontaneous with respect to enthalpy? • If the ΔH = -17.4 kcal/mol and ΔS = 27.2 cal/mol K, is the reaction spontaneous with respect to free energy? • What is the value of ΔG at 300 K?

  17. Problem • Given the reaction: 8 Al(s) + 3 Fe3O4(s) --> 4 Al2O3(s) + 9 Fe(s) + 3350 kJ • Is the reaction endothermic or exothermic? • The sign of ΔH should be positive or negative? • According to enthalpy, is the reaction favored or not favored? • According to entropy, is the reaction favored or not favored? • The sign of ΔS should be positive or negative? • Calculate Gibb’s free energy for this reaction at 25oC if ΔS=215.1 J/K and has the sign you determined in part e. • Is the reaction favored according to free energy?

  18. How do Chemical Reactions Occur? Reaction Rates • DG indicates whether a reaction will occur • But how fast will it occur? • To what extent does the reaction occur?

  19. Rates of Reaction • Rate of Reaction • How fast does a reaction go? • Properly oriented collisions • Sufficient energy to break the bonds of the reactants • Factors affecting collisions and energy • Concentration of reactants • Temperature of system

  20. How do Chemical Reactions Occur? Reaction Rates • Orientation

  21. How do Chemical Reactions Occur? Reaction Rates • Sufficient energy • Energy of activation

  22. Effects of Temperature, Concentration, and Catalysts on Reaction Rates

  23. Effects of Temperature, Concentration, and Catalysts on Reaction Rates • A third way to speed up a reaction is to add a catalyst—a substance that accelerates a chemical reaction but is itself unchanged in the process. • A catalyzed reaction has a lower activation energy.

  24. Problem

  25. Chemical Equilibrium • Equilibrium • To what extent a reaction occurs

  26. Reversible Reactions and Chemical Equilibrium • Many reactions result in complete conversion from reactant to product. • Many however do not

  27. Chemical Equilibrium

  28. Equilibrium Equations and Equilibrium Constants • Consider the following general equilibrium reaction: aA + bB + …  mM + nN + … • Where A, B, … are the reactants; • M, N, …. are the products; • a, b, ….m, n, …. are coefficients in the balanced equation. • At equilibrium, the composition of the reaction mixture obeys an equilibrium equation.

  29. Equilibrium Equations and Equilibrium Constants • The value of K varies with temperature.

  30. Problem • Write an equilibrium constant equation for • N2(g) + 3H2(g) → 2NH3(g) • FeCl3(aq) + 3NaOH(aq) → Fe(OH)3(s) + 3NaCl(aq)

  31. Equilibrium Equations and Equilibrium Constants • K larger than 1000: Reaction goes essentially to completion. • K between 1 and 1000: More products than reactants are present at equilibrium. • K between 1 and 0.001: More reactants than products are present at equilibrium. • K smaller than 0.001: Essentially no reaction occurs.

  32. Problem • Indicate the primary substance or substances in the reaction vessel given the K values of the reactions • 2CO(g) + O2(g) → 2CO2(g) K = 1.4 x 102 • H2O (l) + HNO2(aq) → H3O+(aq) + NO2-(aq) K = 4.50 x 10-4

  33. LeChatelier’s Principle: The effect of Changing Conditions on Equilibia • Le Châtelier's Principle: When a stress is applied to a system at equilibrium, the equilibrium shifts to relieve the stress. • The stress can be any • change in concentration • pressure • Volume • temperature that disturbs original equilibrium.

  34. LeChatelier’s Principle

  35. Le Chatelier’s Principle: The Effect on Changing Conditions on Equilibria

  36. Problem • Methanol can be synthesized by combining carbon monoxide and hydrogen.CO(g) + 2H2(g)   →  CH3OH(g) ΔH°rxn = -90.7 kJ • What happens when • The temperature is raised by 50oC? • The pressure is raised? • Methanol is added? • Hydrogen is removed?

  37. Optional Homework • Text – 7.17, 7.18, 7.19, 7.20, 7.22, 7.23, 7.30, 7.38, 7.40, 7.46, 7.48, 7.54, 7.56, 7.58, 7.62, 7.64, 7.66, 7.68, 7.80 • Chapter 7 Homework online

  38. Required Homework • Assignment 7

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