Sherril Soman Grand Valley State University

1. Lecture Presentation Chapter 17Free Energy and Thermodynamics Sherril Soman Grand Valley State University

2. First Law of Thermodynamics • You can’t win! • The first law of thermodynamics is that energy cannot be created or destroyed. • The total energy of the universe cannot change. • But you can transfer it from one place to another. • DEuniverse = 0 = DEsystem + DEsurroundings

3. First Law of Thermodynamics • Conservation of energy • For an exothermic reaction, “lost” heat from the system goes into the surroundings. • There are two ways energy is “lost” from a system: • Converted to heat, q • Used to do work, w • Energy conservation requires that the energy change in the system is equal to the heat released plus work done. • DE = q + w • DE = DH + PDV • DE is a state function. • Internal energy change independent of how done

4. The Energy Tax • You can’t break even! • To recharge a battery with 100 kJ of useful energy will require more than 100 kJ because of the second law of thermodynamics. • Every energy transition results in a “loss” of energy. • An “energy tax” demanded by nature • Conversion of energy to heat, which is “lost” by heating up the surroundings

5. Heat Tax Fewer steps generally results in a lower total heat tax.

6. Thermodynamics and Spontaneity • Thermodynamics predicts whether a process will occur under the given conditions. • Processes that will occur are called spontaneous. • Nonspontaneous processes require energy input to go. • Spontaneity is determined by comparing the chemical potential energy of the system before the reaction with the free energy of the system after the reaction. • If the system after reaction has less potential energy than before the reaction, the reaction is thermodynamically favorable. • Spontaneity ≠ fast or slow

7. Reversibility of Process • Any spontaneous process is irreversible because there is a net release of energy when it proceeds in that direction. • It will proceed in only one direction. • A reversible process will proceed back and forth between the two end conditions. • Any reversible process is at equilibrium. • This results in no change in free energy. • If a process is spontaneous in one direction, it must be nonspontaneous in the opposite direction.

8. Comparing Potential Energy The direction of spontaneity can be determined by comparing the potential energy of the system at the start and the end.

9. Thermodynamics versus Kinetics Figure 17.4

10. Diamond → Graphite Graphite is more stable than diamond, so the conversion of diamond into graphite is spontaneous. But don’t worry, it’s so slow that your ring won’t turn into pencil lead in your lifetime (or through many of your generations).

11. Spontaneous Processes • Spontaneous processes occur because they release energy from the system. • Most spontaneous processes proceed from a system of higher potential energy to a system at lower potential energy. • Exothermic • But there are some spontaneous processes that proceed from a system of lower potential energy to a system at higher potential energy. • Endothermic • How can something absorb potential energy, yet have a net release of energy?

12. Melting Ice When a solid melts, the particles have more freedom of movement. More freedom of motion increases the randomness of the system. When systems become more random, energy is released. We call this energy, entropy.

13. Melting Ice Melting is an endothermic process, yet ice will spontaneously melt above 0 °C. More freedom of motion increases the randomness of the system. When systems become more random, energy is released. We call this energy, entropy.

14. Water Evaporating

15. Factors Affecting Whether a Reaction Is Spontaneous • There are two factors that determine whether a reaction is spontaneous. They are the enthalpy change and the entropy change of the system. • The enthalpy change, DH, is the difference in the sum of the internal energy and PV work energy of the reactants to the products. • The entropy change, DS, is the difference in randomness of the reactants compared to the products.

16. Enthalpy Change DH generally measured in kJ/mol • Stronger bonds = more stable molecules • A reaction is generally exothermic if the bonds in the products are stronger than the bonds in the reactants. • Exothermic =energy released; DH is negative. • A reaction is generally endothermic if the bonds in the products are weaker than the bonds in the reactants. • Endothermic =energy absorbed; DH is positive. • The enthalpy change is favorable for exothermic reactions and unfavorable for endothermic reactions.

17. Entropy • Entropyis a thermodynamic function that increases as the number of energetically equivalent ways of arranging the components increases, S. • S generally J/mol • S = k ln W • k = Boltzmann constant = 1.38 × 10−23 J/K • W is the number of energetically equivalent ways a system can exist. • Unitless • Random systems require less energy than ordered systems.

18. Salt Dissolving in Water

19. W These are energetically equivalent states for the expansion of a gas. It doesn’t matter, in terms of potential energy, whether the molecules are all in one flask, or evenly distributed. But one of these states is more probable than the other two.

20. Macrostates → Microstates This macrostate can be achieved through several different arrangements of the particles.

21. Macrostates → Microstates These microstates all have the same macrostate. So there are six different particle arrangements that result in the same macrostate.

22. Macrostates and Probability There is only one possible arrangement that gives state A and one that gives state B. There are six possible arrangements that give state C. The macrostate with the highest entropy also has the greatest dispersal of energy. Therefore, state C has higher entropy than either state A or state B. There is six times the probability of having the state C macrostate than either state A or state B.

23. Changes in Entropy, DS DS = Sfinal− Sinitial • Entropy change is favorable when the result is a more random system. • DS is positive. • Some changes that increase the entropy are as follows: • Reactions whose products are in a more random state • Solid more ordered than liquid; liquid more ordered than gas • Reactions that have larger numbers of product molecules than reactant molecules • Increase in temperature • Solids dissociating into ions upon dissolving

24. DS For a process where the final condition is more random than the initial condition, DSsystem is positive and the entropy change is favorable for the process to be spontaneous. For a process where the final condition is more orderly than the initial condition, DSsystem is negative and the entropy change is unfavorable for the process to be spontaneous. DSsystem = DSreaction = Sn(S°products) −Sn(S°reactants)

25. Entropy Change in State Change • When materials change state, the number of macrostates it can have changes as well. • The more degrees of freedom the molecules have, the more macrostates are possible. • Solids have fewer macrostates than liquids, which have fewer macrostates than gases.

26. Entropy Change and State Change Figure 17.5 pg 823

27. The SecondLaw of Thermodynamics • The second law of thermodynamics says that the total entropy change of the universe must be positive for a process to be spontaneous. • For reversible process DSuniv = 0 • For irreversible (spontaneous) process DSuniv > 0 • DSuniverse = DSsystem + DSsurroundings • If the entropy of the system decreases, then the entropy of the surroundings must increase by a larger amount. • When DSsystem is negative, DSsurroundings must be positive and big for a spontaneous process.

28. Heat Flow, Entropy, and the Second Law When ice is placed in water, heat flows from the water into the ice. According to the second law, heat must flow from water to ice because it results in more dispersal of heat. The entropy of the universe increases.

29. Heat Transfer and Changes in Entropy of the Surroundings • The second law demands that the entropy of the universe increase for a spontaneous process. • Yet processes like water vapor condensing are spontaneous, even though the water vapor is more random than the liquid water. • If a process is spontaneous, yet the entropy change of the process is unfavorable, there must have been a large increase in the entropy of the surroundings. • The entropy increase must come from heat released by the system; the process must be exothermic!

30. Entropy Change in the System and Surroundings When the entropy change in a system is unfavorable (negative), the entropy change in the surroundings must be favorable (positive), and large in order to allow the process to be spontaneous. Cartoon from page 825

31. Heat Exchange and DSsurroundings • When a system process is exothermic, it adds heat to the surroundings, increasing the entropy of the surroundings • When a system process is endothermic, it takes heat from the surroundings, decreasing the entropy of the surroundings. • The amount the entropy of the surroundings changes depends on its original temperature. • The higher the original temperature, the less effect addition or removal of heat has.

33. Temperature Dependence of DSsurroundings • When heat is added to surroundings that are cool it has more of an effect on the entropy than it would have if the surroundings were already hot. • Water freezes spontaneously below 0 °C because the heat released on freezing increases the entropy of the surroundings enough to make DS positive. • Above 0 °C the increase in entropy of the surroundings is insufficient to make DS positive.

34. Quantifying Entropy Changes in Surroundings • The entropy change in the surroundings is proportional to the amount of heat gained or lost. • qsurroundings = −qsystem • The entropy change in the surroundings is also inversely proportional to its temperature. • At constant pressure and temperature, the overall relationship is as follows:

35. Gibbs Free Energy and Spontaneity • It can be shown that −TDSuniv = DHsys−TDSsys. • The Gibbs free energy, G,is the maximum amount of work energy that can be released to the surroundings by a system for a constant temperature and pressure system. • Gibbs free energy is often called the chemical potential because it is analogous to the storing of energy in a mechanical system. • DGsys = DHsys−TDSsys • Because DSuniv determines if a process is spontaneous, DG also determines spontaneity. DSuniv is positive when spontaneous, so DG is negative.

36. Gibbs Free Energy, DG • A process will be spontaneous when DG is negative. • DG will be negative when • DH is negative and DS is positive. • Exothermic and more random • DH is negative and large and DS is negative but small. • DH is positive but small and DS is positive and large. • Or high temperature • DG will be positive when DH is positive and DS is negative. • Never spontaneous at any temperature. • When DG = 0 the reaction is at equilibrium.

37. Free Energy Change and Spontaneity

38. DG, DH, and DS

39. Standard Conditions • The standard stateis the state of a material at a defined set of conditions. • Gas = pure gas at exactly 1 atm pressure. • Solid or Liquid = pure solid or liquid in its most stable form at exactly 1 atm pressure and temperature of interest. • Usually 25 °C • Solution = substance in a solution with concentration 1 M.

40. The Third Law of Thermodynamics:Absolute Entropy • The absolute entropyof a substance is the amount of energy it has due to dispersion of energy through its particles. • The third law states that for a perfect crystal at absolute zero, the absolute entropy = 0 J/mol ∙ K. • Therefore, every substance that is not a perfect crystal at absolute zero has some energy from entropy. • Therefore, the absolute entropy of substances is always positive.

41. Standard Absolute Entropies • S° • Extensive • Entropies for 1 mole of a substance at 298 K for a particular state, a particular allotrope, a particular molecular complexity, a particular molar mass, and a particular degree of dissolution

42. Standard Absolute Entropies

43. Relative Standard Entropies: States • The gas state has a larger entropy than the liquid state at a particular temperature. • The liquid state has a larger entropy than the solid state at a particular temperature.

44. Relative Standard Entropies: Molar Mass • The larger the molar mass, the larger the entropy. • Available energy states are more closely spaced, allowing more dispersal of energy through the states.

45. Relative Standard Entropies: Allotropes • The less constrained the structure of an allotrope is, the larger its entropy • The fact that the layers in graphite are not bonded together makes it less constrained.

46. Relative Standard Entropies: Molecular Complexity • Larger, more complex molecules generally have larger entropy. • More energy states are available, allowing more dispersal of energy through the states.

47. “Places” for Energy in Gaseous NO

48. Relative Standard Entropies Dissolution • Dissolved solids generally have larger entropy, distributing particles throughout the mixture.

49. The Standard Entropy Change, DS • The standard entropy change is the difference in absolute entropy between the reactants and products under standard conditions. DSºreaction = (∑npSºproducts) − (∑nrSºreactants) • Remember, although the standard enthalpy of formation, DHf°, of an element is 0 kJ/mol, the absolute entropy at 25 °C, S°, is always positive.

50. Calculating DG • At 25 °C • At temperatures other than 25 C • Assuming the change in DHoreaction and DSoreaction is negligible • Or