Energy, Rates, and Equilibrium in Chemical Reactions

1. Chapter 7 Energy, Rates, and Equilibrium

2. Energy The capacity to do work

3. The Two Types of Energy Potential:due to position or composition – One must compare what can be made from what you have in order to determine the potential. Kinetic:due to motion of the object KE = 1/2 mv2 (m = mass, v = velocity)

4. Food Calories • Food is stored chemical potential energy. • Energy is produced by reacting the food we eat with oxygen. • Some of the energy can be used to do work!

5. Exothermic Reactions • Some reactions give off heat. • Once the magnesium begins to burn, the heat given off makes more magnesium burn. • Link to Magnesium burning

6. Endothermic Reactions • Some reactions use heat during a reaction. • The hot water supplies the heat needed to sublime the solid carbon dioxide into gaseous carbon dioxide. • Link to Sublimation Video

7. Energy • Potential - Stored energy • It is not the energy in the bonds of the compound you have in your hands that makes it have potential energy. • It is the energy of the new bonds that are formed in a chemical reaction that makes for the potential energy in the reactants.

8. Energy Diagrams • It is the difference which determines the potential

9. Covalent Bond Strength • Most simply, the strength of a bond is measured by determining how much energy is required to break the bond. • This is the bond enthalpy. • The bond enthalpy for a Cl—Cl bond is measured to be 242 kJ/mol.

10. DH0 = 436.4 kJ H2 (g) H (g) + H (g) DH0 = 242.7 kJ Cl2 (g) Cl (g) + Cl (g) DH0 = 431.9 kJ HCl (g) H (g) + Cl (g) DH0 = 498.7 kJ O2 (g) O (g) + O (g) O DH0 = 941.4 kJ N2 (g) N (g) + N (g) O Bond Energies Single bond < Double bond < Triple bond N N The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy. Bond Energy 9.10

11. Average Bond Enthalpies • This table lists the average bond enthalpies (kJ/mol) for many different types of bonds. • Average bond enthalpies are positive, because bond breaking is an endothermic process.

12. Energy • The energy required to break a bond is equal to the energy liberated when making a bond. • LAW of conservation of energy – • Neither created nor destroyed during physical or chemical changes

13. Enthalpies of Reaction • Compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed. Hrxn = (bonds broken)(bonds formed)

14. Enthalpies of Reaction CH4(g) + Cl2(g) CH3Cl(g) + HCl(g) In this example, One C—H bond and one Cl—Cl bond are broken; one C—Cl and one H—Cl bond are formed. So, Hrxn = [D(C—H) + D(Cl—Cl)  [D(C—Cl) + D(H—Cl) = [(413 kJ) + (242 kJ)]  [(328 kJ) + (431 kJ)] = (655 kJ)  (759 kJ) = 104 kJ

15. Free Energy, Enthalpy, Entropy • DH = Total heat energy of a system at constant pressure. • Not all of this can be used to do work. • DG = That part which can do work • DS x T = That part which cannot do work. • DG = DH - T DS • If DG is favorable, the process will proceed and work can be extracted from the process. • If DG is unfavorable, work must be continually added to make the process proceed.

16. Reaction Rates • Even if energy can be extracted from a system, how fast the energy is produced can make or break the process literally. • Example: Atomic bomb vs. nuclear energy • Link

17. Reaction Rates • The rates of chemical reactions are affected by the following factors • molecular collisions • activation energy • nature of the reactants • concentration of the reactants • temperature • presence of a catalyst • On the following screens, we examine these factors one at a time

18. Molecular Collisions • two species must collide to react • calculations show that the rate things collide is far greater than the rate at which they react • Conclusion: most collisions do not result in a reaction • a collision that does result in a reaction is called an effective collision • there are two main reasons why some collisions are effective and others are not; • energy • orientation

19. Molecular Collisions • Activation energy: the minimum energy required for a reaction to take place • energy is required for reactions to begin even if they give off energy during the process • this energy comes from collisions • if the collision energy is large, there is sufficient energy to break the necessary bonds, and reaction takes place • if the collision energy is too small, no reaction occurs

20. Molecular Collisions • Orientation at the time of collision • the colliding particles must be properly oriented for bond breaking and bond making • for example, to be an effective collision between H2O and HCl, the oxygen of H2O must collide with the H of HCl so that the new O-H bond can form and the H-Cl bond can break

21. Energy Diagrams • Energy diagram for an exothermic reaction

22. Energy Diagrams • The reaction of H2 and N2 to form ammonia is exothermic • in this reaction, six covalent bonds are broken and six new ones are formed • breaking a bond requires energy, and forming a bond releases energy • in this reaction, the energy released in making the six new bonds is greater than the energy required to break the six original bonds; the reaction is exothermic

23. Energy Diagrams • Energy diagram for an endothermic reaction

24. Energy Diagrams • Transition state: a maximum on an energy diagram • the transition state for the reaction between H2O and HCl probably looks like this, in which the new O-H bond is partially formed and the H-Cl bond is partially broken

25. Factors Affecting Rate • Nature of reactants • in general, reaction between ions in aqueous solution are very fast (activation energies are very low) • in general, reaction between covalent compounds, whether in water or another solvent, are slower (their activation energies are higher) • Concentration • in most cases, reaction rate increases when the concentration of either or both reactants increases • for many reactions, there is a direct relationship between concentration and reaction rate; when concentration doubles the rate doubles

26. Factors Affecting Rate • Temperature • in virtually all reactions, rate increases as temperature increases • an approximate rule for many reactions is that for a 10°C increase in temperature, the reaction rate doubles • when temperature increases, molecules move faster (have more kinetic energy), which means that they collide more frequently; more frequent collisions mean higher reaction rates • not only do molecules move faster at higher temperatures, but the fraction of molecules with energy equal to or greater than the activation energy also increases

27. Factors Affecting Rate • The distribution of kinetic energies (molecular velocities) at two temperatures

28. Factors Affecting Rate • Catalyst: a substance that increases the rate of a chemical reaction without itself being used up

29. Factors Affecting Rate • Many catalysts provide a surface on which reactants can meet • the reaction of ethylene with hydrogen is an exothermic reaction • if these two reagents are mixed, there is no visible reaction even over long periods of time • when they are mixed and shaken with a finely divided transition metal catalyst, such as Pd, Pt, or Ni, the reaction takes place readily at room temperature

30. Reversible Reactions • Reversible reaction: one that can be made to go in either direction • if we mix CO and H2O in the gas phase at high temperature, CO2 and H2 are formed • we can also make the reaction take place the other way by mixing CO2 and H2 • the reaction is reversible, and we can discuss both a forward reaction and a reverse reaction

31. Reversible Reactions • Equilibrium: a dynamic state in which the rate of the forward reaction is equal to the rate of the reverse reaction • at equilibrium there is no change in concentration of either reactants or products • reaction, however, is still taking place; reactants are still being converted to products and products to reactants, but the rates of the two reactions are equal

32. Equilibrium Constants • Equilibrium constant, K: the product of the concentration of products of a chemical equilibrium divided by the concentration of reactants, each raised to the power equal to its coefficient in the balanced chemical equation • for the general reaction • the equilibrium constant is

33. Equilibrium Constants • Problem: write the equilibrium constant for this reversible reaction • solution: for this reaction, K is • note that no exponents are shown in this equilibrium constant; by convention the exponent “1” is not written

34. Equilibrium Constants • Problem: when H2 and I2 react at 427°C, the following equilibrium is reached • the equilibrium concentrations are [I2] = 0.42 mol/L, [H2] = 0.025 mol/L, and [HI] = 0.76 mol/L. Using these values, calculate the value of K • Solution: this K has no units because molarities cancel

35. Equilibrium and Rates • There is no relationship between a reaction rate and the value of K • reaction rate depends on the activation energy of the forward and reverse reactions; these rates determine how fast equilibrium is reached but not its position • it is possible to have a large K and a slow rate at which equilibrium is reached • it is also possible to have a small K and a fast rate at which equilibrium is reached • it is also possible to have any combination of K and rate in between these two extremes

36. LeChatelier’s Principle • LeChatelier’s Principle: when a stress is applied to a chemical system at equilibrium, the position of the equilibrium shifts in the direction to relieve the applied stress • We look at three types of stress that can be applied to a chemical equilibrium • addition of a reaction component • removal of a reaction component • change in temperature

37. LeChatelier’s Principle • Addition of a reaction component • suppose this reaction reaches equilibrium • suppose we now disturb the equilibrium by adding some acetic acid • the rate of the forward reaction increases and the concentrations of ethyl acetate and water increase • as this happens, the rate of the reverse reaction also increases • in time, the two rates will again become equal and a new equilibrium will be established

38. LeChatelier’s Principle • at the new equilibrium, the concentrations of reactants and products again become constant, but not the same as they were before the addition of acetic acid • the concentrations of ethyl acetate and water are now higher, and the concentration of ethanol is lower • the concentration of acetic acid is also higher, but not as high as it was immediately after we added the extra amount • the system has relieved the stress by increasing the components on the other side of the equilibrium • we say that the system has shifted to minimize the stress

39. LeChatelier’s Principle • Removal of a reaction component • removal of a component shifts the position of equilibrium to the side that produces more of the component that has been removed • suppose we remove ethyl acetate from this equilibrium • if ethyl acetate is removed, the position of equilibrium shifts to the right to produce more ethyl acetate and restore equilibrium • the effect of removing a component is the opposite of adding one

40. LeChatelier’s Principle • Problem: when acid rain attacks marble (calcium carbonate), the following equilibrium can be written how does the fact that CO2 is a gas influence the equilibrium? • Solution: CO2 gas diffuses from the reaction site, and is removed from the equilibrium mixture; the equilibrium shifts to the right and the marble continues to erode

41. LeChatelier’s Principle • Change in temperature • the effect of a change in temperature on an equilibrium depends on whether the forward reaction is exothermic or endothermic • consider this exothermic reaction • we can look on heat as a product of the reaction • adding heat (increasing the temperature) pushes the equilibrium to the left • removing heat (decreasing the temperature) pushes the equilibrium to the right

42. LeChatelier’s Principle • summary of the effects of change of temperature on a system in equilibrium

43. Chapter 7 End Chapter 7