Thermodynamics

1. Thermodynamics Ch. 19

2. Why thermodynamics? • So far we have answered TWO really important questions regarding chemical reactions... • HOW FAST does a reaction make products? • This is the study of Chemical Kinetics • HOW MUCH products are made? • This is the study of Chemical Equilibrium

3. Why thermodynamics? • Now we will answer probably the most important question of all (at least to chemists and engineers)... • DOES THE REACTION EVEN HAPPEN!?!?!??? • The answer to this question is based on the study of Chemical Thermodynamics (the energy relationship between reactants and products)

4. Spontaneity • Spontaneous processes are processes that occur in a definite direction without any outside interventions on the system • Examples: • Spontaneous Non-spontaneous • Egg falls and breaks Broken egg goes back together • Ice melts at room T Water freezes at room T

5. Spontaneity • Unlike reactions at equilibrium (in which both the forward and reverse reactions occur)... • If the forward reaction is spontaneous, the reverse reaction is nonspontaneous! • If the reverse reaction is spontaneous, the forward reaction is nonspontaneous! • *Both directions may happen, but one requires "outside intervention" (an energy input)

6. Spontaneity • Temperature affects spontaneity • Example: ice melts spontaneously at 25oC, but not at -10oC • Other things also affect spontaneity, but they are beyond the scope of this class

7. Enthalpy & Spontaneity • Remember enthalpy? • At constant pressure (such as in open glassware in the lab) enthalpy is equal to the heat transfer between the system and the surroundings • We can't measure enthalpy directly, so we always find the change in enthalpy, DH. • DH = Hfinal– Hinitial

8. Enthalpy & Spontaneity • Remember how we get DH values? • Measure it in the lab (calorimetry) • Hess' Law (indirectly from other reactions) • Molar heats of formation (tabulated data from the work of others) • Thermochemical equations • Enthalpy diagrams

9. Molar Heats of Formation Revisited • Molar heats of formation are enthalpy changes for substances formed directly from elements • Like this, where CH4 is formed directly from its elements • Example: • C + 4H  CH4 • NOT like this, where CH4 is formed from reacting other substance together elements • CO2+ 2H2O  CH4 + 2O2

10. Molar Heats of Formation Revisited • Molar heats of formation are based on a "standard state" reaction happening at 1 atm and 298 Kelvins • Elements are assigned a molar enthalpy value of zero • DHorxn = SnDHof(products) - SnDHof(reactants) • where "n" and "m" are stoichiometric coefficients

11. Molar Heats of Formation Revisited EXAMPLE: 2H2S(g) + 3O2(g)  2H2O(l) + 2SO2(g) What is DHorxn if:

12. Summary of Enthalpy and Spontaneity • Things in our universe want to be low energy, right? • There is a tendency to go from high energy to low energy • So if DH = Hfinal - Hinitial, then -DH values are "favored" Remember… if DH < 0 Exothermic if DH > 0 Endothermic

13. Summary of Enthalpy and Spontaneity • So...a decrease in enthalpy (an exothermic process) is a favorable contribution toward spontaneity. • So why does ice melt spontaneously at room temperature when melting is an endothermicprocess??!?!?! • There must be other things that also contribute to spontaneity…

14. Entropy & Spontaneity • Just like there is a trend in our universe for things to become lower in energy, there is also a trend in our universe for things to become more disordered • The amount of disorder (randomness) in a system is called entropy. • When a system becomes more disordered, we say its entropy increases. • Entropy is represented by the letter "S”

15. Entropy & Spontaneity • We can't calculate exact entropy values, so (just like enthalpy) we calculate changes in entropy, DS DS = Sfinal- Sinitial If DS > 0, entropy increased (became more disordered) If DS < 0, entropy decreased (became more ordered)

16. Entropy & Spontaneity • The 1st Law of Thermodynamics: • Energy is conserved • We can't create it or destroy it • The 2nd Law of Thermodynamics: • Entropy is NOT conserved... it is increasing • In other words, we CAN create entropy • Things tend to go from Order Disorder (…think about the cleanliness of your room!!!)

17. Entropy & Spontaneity • Examples of Order and Disorder: • A collection of marbles: If you put a bunch of marbles in a box, will they form orderly stacks of marbles or will they spread out sort of randomly?

18. Entropy & Spontaneity • Examples of Order and Disorder: • A new deck of poker cards: All the cards are ordered in sequence and suit. What happens to their order when you throw the deck into the air? Do the cards all land in the same order of sequence and suit?

19. Entropy & Spontaneity • Examples of Order and Disorder: • Think about a perfect crystal of ice. All the H2O molecules are specifically arranged and turned a certain way to allow H-O intermolecular bonding. But when the ice melts, molecules start to move more freely and randomly. And when liquid water is boiled, think about the rapid, random motion of the gaseous water molecules.

20. Entropy & Spontaneity • …so things tend to go spontaneously from Order to Disorder • If DS = Sfinal - Sinitial, then our universe "favors" +DS values right? Because Sfinalwill be larger than Sinitial. There's more disorder at the end of the process than at the beginning! • THEREFORE... an increase in entropy is a favorable contribution toward spontaneity.

21. Entropy & Spontaneity • What about that spontaneous melting of ice? • The enthalpy change for it is positive (+DH) which is NOT favorable • But the entropy change for it is positive (+DS) which IS favorable

22. Entropy & Spontaneity • What about that spontaneous melting of ice? • We know the process happens simply based on our experience so evidently the influence of entropy wins out over enthalpy (in this case) • So the spontaneity of a chemical reaction depends on the balance between enthalpy and entropy. The net effect of both contributions is seen in the equation for something called "Free Energy”

23. Some Common Entropy Changes • 1. Changes in phase • Entropy increases (+DS) as you go from a solid  liquid  gas SOLID Crystal structure is very ordered Repeating pattern of atoms Very little particle movement LIQUID Less ordered More particle movement as particles can slide past each other GAS Constant, random, rapid motion of Particles very disordered and unpredictable motion

24. Some Common Entropy Changes • 2. Solids dissolving • Entropy increases (+DS) when solids dissolve • NaCl crystal is very ordered NaCl(aq) is more disordered • and doesn’t allow much Na+ and Cl- ions are free to • movement move about Cl-Na-Cl-Na-Cl-Na Na-Cl-Na-Cl-Na-Cl Cl-Na-Cl-Na-Cl-Na Na-Cl-Na-Cl-Na-Cl Na+ Na+ Cl- Na+ Cl- Cl- Na+ Na+ Cl-

25. Some Common Entropy Changes • 3. Increasing the number of gaseous molecules during a reaction • Causes an increase in entropy (+DS) • 2NH3(g) N2(g) + 3H2(g) • 2 moles of gas have 4 moles of gas have • less disorder; there more disorder; there • are fewer particles are more particles • moving in random moving in random • paths paths

26. Some Common Entropy Changes • 4. Gas diffusion increases entropy • 5. Higher temperatures increase entropy

27. Some Common Entropy Changes • EXAMPLE 1: Predict the sign of DS for • a candle burning • butter melting • ammonia vapor condensing • tea dissolving in water • EXAMPLE 2: Predict the sign for • 2Cl(g) Cl2(g) • 2H2(g) + O2(g)  H2O(l)

28. Calculating DS Values • We will calculate "standard molar entropies" by looking up values in the appendix of our textbook just like when we calculated enthalpy values. (NOTE: The units for entropy are J/mol K) • Standard molar entropy values are based on the 3rd Law of Thermodynamics which states: All elements and compounds in a perfect crystal at 0 Kelvins have an entropy of 0

29. Calculating DS Values • DSorxn = SnDSof(products) - SnDSof(reactants) • where "n" and "m" are stoichiometric coefficients

30. Free Energy • The spontaneity of a reaction involves two thermodynamic concepts: • Entropy • Enthalpy

31. Free Energy • In general... • Spontaneity is favored by an increase in entropy (+DS) • (Things in our universe spontaneously get more disordered) • Spontaneity is favored by a decrease in enthalpy (-DH) • (Things in our universe spontaneously lose energy)

32. Free Energy • Josiah Willard Gibbs (1839 - 1903) proposed a relationship that shows how entropy and enthalpy can be combined to predict reaction spontaneity • He said spontaneity was a function of the "free energy" possessed by a reaction

33. Free Energy • DG = DH - TDS • DG Free energy of a system • DH Enthalpy • TAbsolute temperature (Kelvins) • DS Entropy • When temperature and pressure are constant: • If G < 0 forward reaction is spontaneous • If G > 0 forward reaction is nonspontaneous (reverse reaction is spontaneous) • If G = 0 reaction is at equilibrium

34. Free Energy • “Free energy” is the portion of the energy change of a spontaneous reaction that is free to do work • The rest of the energy is released as heat • If the process is NON-spontaneous, the value of DG will represent the amount of work required to make the process happen

35. Calculating DG Values • DGorxn= SnDGof(products) - SnDGof(reactants) • where "n" and "m" are stoichiometric coefficients • So you'll have to look up DG values in a table in your book • Again, pure elements have a value of zero!

36. Calculating DG Values EXAMPLE: Is the following process spontaneous? 2H2O2(l)  2H2O(l) + O2(g)

37. Free Energy & Temperature This is guaranteed to be on the AP Test in some form!!!

38. Free Energy & Temperature In the last two situations above, you can see that as temperature changes, the process goes from being spontaneous to nonspontaneous or vice versa.

39. Free Energy & Temperature • The temperature where this happens is called the crossover temperature and is found by solving for T when DG = 0 (at equilibrium) • 0 = DH - TDS

40. Free Energy & Temperature EXAMPLE: Consider the reaction: 2As(s) + 3F2(s)  2AsF3(l) If DH = -1897.9 kJ and DS = -0.318 kJ/K, at what temperature is the reaction spontaneous?

41. Free Energy & The Equilibrium Constant • The DG values from tables are all at standard conditions which means: • Solids are pure, liquids are pure, gases are at 1 atm, aqueous solutions are 1M, and elements have zero free energy • The problem is...most reactions don't happen at "standard conditions"

42. Free Energy & The Equilibrium Constant To calculate Free Energy at non-standard conditions, use: DG = DGo + RTlnQ where R = 8.314 J/mol K T = absolute temp in Kelvins Q = the reaction quotient

43. Free Energy & The Equilibrium Constant • At equilibrium, DG = 0 and Q = K, so.... • DG = DGo + RTlnQ • 0 = DGo + RTlnK • DGo = -(RTlnK) • …and it follows that • K = e-DG/RT • so the more negative "DG" is, the larger "K" will be If… DG < 0, then K > 1 DG = 0, then K = 1 DG > 0, then K < 1