Thermochemistry

1. Thermochemistry Ch. 5 in textbook (omit Section 5.8) Phsstudent.blogspot.com

2. Heat vs. Work A. Energy (E) • Energy = ability to transfer heat or do work • Kinetic Energy = energy of an object in motion; depends on the mass and velocity of the object • Potential Energy= energy of an object due to its composition or position relative to other objects; thought of as stored energy because it is converted to kinetic energy when changes occur Techgage.com

3. B. Work • A force (push/pull) acting over a distance • Formula: w = F x d • Energy can never be in the form of work ONLY, since some heat will be lost due to friction Artandwork.com

4. C. Heat (q) • The transfer of energy from high to low temperature • Heat always flows from high to low • Frictional heat is generated when particles move against one another; this heat is “lost” in the sense that it can no longer be utilized Jcwinnie.biz

5. D. Relating Energy, Work, and Heat • The overall energy change is represented by the following equation: • ΔE = q + w • Units of energy: Joules (J), calories (cal) or Calories (Cal) • Joule = Newtons x meters • 1 N = 1 kg x m/s2 • Therefore 1 J = 1 kg x m2/s2 • 1 cal = amount of heat needed to raise 1 gram of water 1 ºC • 1000 cal = 1 kcal = 1 Cal (dietary calorie) Vat19.com

6. The 1st Law of Thermodynamics A. Definitions • System = the specific part of the universe on which we are focusing • Surroundings = everything else in the universe other than the system • Internal Energy = total amount of energy (kinetic and potential) in the system Splung.com

7. B. The Law • Although the energy of the system and surroundings may change, the overall energy of the universe does not (nor does the energy of the schuniverse) • This is also known as the Law of Conservation of Energy: energy is neither created nor destroyed, only transferred Stutterlaw.com

8. Changes in Internal Energy • We don’t necessarily know the internal energy of the system before and after a change. However, because the 1st law deals with net changes to the system, we only care about the actual value of ΔE • If ΔE is (+), then the internal energy of the system increases • If ΔE is (-), then the internal energy of the system decreases Apchemistry.wikispaces.com Ctaikido.com

9. D. ΔE, q, & w (again)

10. HW: 5.14, 15.16 (a), 15.18

11. E. State Functions • A state function is a property that depends on the current state of a system, not the pathway or history of the system • q and w are NOT state functions, but their sum q + w = ΔE is since ΔE only depends on the final and initial states of the system • ΔE can be gained or lost as all heat or heat and work (can never be just work, due to friction); either way, the ΔE will be the same at the end, regardless of the combination of heat and work Designyoucanwear. spreadshirts.com Enduringamerica.com

12. Analogy from the textbook: Chicago is 596 ft above sea level while Denver is 5280 ft above sea level; you can take any route from Chicago to Denver (distance is not a state function), but the altitude change will always be 4684 ft (altitude is a state function) Ohdeedoh.com

13. III. Enthalpies of Reaction • When the heat lost or gained by a system is at a constant pressure (qp), it is equal to the enthalpy change (ΔH) of the system • Also, since very little work is done in a chemical reaction, we assume that ΔE = ΔH for a chemical system Clickandlearn.com HW: 5.28

14. IV. Calorimetry • Used to experimentally determine enthalpy values • We have already used the mathematical formula: q = mCΔT • Read this section for context, if needed Vat19.com Hartfordphysics.wikispaces.com HW: 5.40

15. V. Hess’s Law • If a reaction is carried out in a series of steps, the heat of reaction (ΔH) can be calculated by adding up the enthalpy changes of the individual steps • This is especially useful when determining the heat of reaction directly is difficult • Sometimes we have to rearrange and modify the given steps so that they “add up” and “cancel out” to the correct final reaction • If we reverse a step, we change the sign of ΔH • If we multiply the coefficients, we do the same to ΔH Todayinsci.com En.wikibooks.org Link

16. Ex) Calculate the ΔH for the following reaction: 2C(s) + H2(g) → C2H2(g) Given the following reactions: C2H2(g) + 5/2 O2(g) → 2CO2(g) + H2O(l) ΔH = -1299.6 kJ C(s) + O2(g) → CO2(g) ΔH = -393.5 kJ H2(g) + 1/2 O2(g) → H2O(l) ΔH = -285.8 kJ

17. Hess is watching you! Todayinsci.com HW: 5.54

18. VI. Enthalpies of Formation • Also known as heats of formation (ΔHf); heat change associated with forming a compound from its constituent elements • To standardize these values, we use standard enthalpies (all reactants and products are in their most stable, standard states at atmospheric pressure and room temp (25 ºC)) • The standard enthalpy of formation (ΔHfº) is the enthalpy change that forms 1 mol of a compound from its elements, with all substances in their standard states • FYI, the standard state of oxygen is O2 not O3 and the standard state of carbon is graphite not diamond • By definition, the ΔHfº of the most stable form of an element is 0 Vat19.com Mnhe.com

19. We can now apply Hess’s Law to calculate the heat of reaction using the heats of formation of all reactants and products! • Ex) C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l) Wiki.injuryboard.com

20. Hess is STILL watching you! Todayinsci.com HW: 5.60

21. Vat19.com

22. Chemical Thermodynamics Part 1 Ch. 19 In Text Toothpastefordinner.com

23. I) SpontaneityA) Definitions • In any physical or chemical process, there is always a preferred direction said to be spontaneous • We tend to think of this as a process that occurs without any outside intervention • The direction that is NOT preferred is called nonspontaneous Educonline.com

24. Spontaneity is dependent on temperature, not just the process • Ex) Water Fehd.gov.hk

25. B) Reversibility • A reversible process is one in which the original state can be restored by exactly reversing the change • There is no net change in the system or the surroundings after the change is reversed • This is ONLY true for a system at equilibrium • Ex) Ice at its melting point Freshnessmag.com

26. An irreversible process is one in which a different path (different values of q and w) must be taken to return to the original state • Although the system may be restored, the surroundings are changed • Ex) Expansion of a gas • Ex) Ice above or below its melting point Freefoto.com HW: 19.6, 19.8

27. If a process is spontaneous, then it must be irreversible! • Spontaneous processes are not necessarily fast, just the preferred direction… • In general, exothermicreactions are spontaneous; however some endothermic reactions are spontaneous, so what’s driving them to be spontaneous?

28. II) Entropy • In an isothermal expansion of a gas, no heat or work is done to expand a gas but the increase in entropy causes the diffusion to be spontaneous Ch.ntu.edu.tw

29. Entropy is a state function (just like internal energy, ΔE, and enthalpy, ΔH) • ΔS = Sf – Si • ΔS = qrev/T (at const. T) • ΔS has the units J/K • If ΔS = +, then the products have more entropy than the reactants (look for gases!) • If ΔS = -, then the products have less entropy than the reactants (look for solids!)

30. III) The 2nd Law of Thermodynamics • The entropy of the universe can NEVER decrease, instead it must remain 0 or increase • Unlike the 1st law in which energy is conserved, entropy is NOT conserved Zazzle.com

31. ΔSuniv = ΔSsys + ΔSsurr • Reversible process: ΔSuniv = ΔSsys + ΔSsurr = 0 • Irreversible process: ΔSuniv = ΔSsys + ΔSsurr >0 • ex) rusting Share.ehs.uen.org

32. In an isolated system (the system cannot exchange heat, work, or matter with the surroundings): • ΔSsys = 0 • ΔSsys > 0 Tutorvista.com HW: 19.22, 19.28

33. IV) The 3rd Law of Thermodynamics • Each type of molecular motion (vibrational, translational, and rotational) is related a degree of freedom which translates into an increase in entropy • The fewer the bonds, the lower the IMFs, and the higher the temperature, the greater the number of degrees of freedom • The greater the bonds, the higher the IMFs, and the lower the temperature, the lower the number of degrees of freedom Flickr.com

34. The Law states that the entropy of a pure crystalline solid at absolute zero is 0 (S @ 0 K = 0) • This is the theoretical definition for “perfect” order where there are no degrees of freedom • NOTE: although entropy is temp. dependent, it is NOT related to enthalpy Entiremonthfree.com

35. V) Entropy Calcs • Standard Molar Entropy: Sº = J/mol•K • Sº for elements are NOT zero • Sº for gases are greater than those of solids and liquids • Sº increases with increasing molar mass • Sºincreases with increasing number of atoms in the formula Uwsp.edu Link

36. ΔSº = ΣSºproducts – ΣSºreactants • Ex) Calculate the standard entropy change for the synthesis of ammonia from its elements. Physics.ubc.ca HW: 19.34 (a) & (d)

37. Last.fm Chemical Thermodynamics Part 2 (Ch. 19 in Text)

38. VI) Gibbs Free Energy • Spontaneous changes are favored by an increase in entropy and a decrease in energy(exothermic) • Spontaneity and its relation to these 2 factors was quantified by J.W. Gibbs in the late 1800s • The state function known as free energy predicts the degree of spontaneity as follows: • ΔG = ΔH - TΔS Chemistry.about.com

39. If ΔG = -, the forward rxn is spon. • If ΔG = +, the forward rxn is nonspon. and work must be done by the surroundings on the system for it to occur • If ΔG = 0, the rxn is at equilibrium Cafepress.com Link

40. Standard Free Energies of Formation • Tells us the free energy change when a compound is formed from its elements in their standard states • As with ΔHºf, ΔGºf for an element is 0 • ΔGºf = ΣGºproducts – ΣGºreactants • Can also be used to calculate the standard free energy change of any reaction (remember how Hess’s Law was applied to standard enthalpies of formation in the last chapter) Teachers.northallegan.org

41. Ex) Using Appendix C, calculate the standard free energy for the combustion of methane. Collegesurfing.com

42. VII) Free Energy and Temperature • Assuming standard states and conditions, the Gibbs equation becomes: • ΔGº = ΔHº - TΔSº • Notice that the spontaneity (or the degree of spontaneity) may be dependent on the temperature of the system Alohapoolandspaservice.com

43. ΔGº = ΔH º - TΔSº • Under what conditions is a reaction ALWAYS spontaneous? • Under what conditions is a reaction NEVER spontaneous? Brandautopsy.typepad.com

44. ΔGº = ΔH º - TΔSº • Ex) H2O(s) → H2O(l) • Ex) H2O(l) → H2O(s) Marionsilver.wordpress.com

45. Ex) Calculate the standard free energy change for the formation of ammonia at 500. ºC. Luckymojo.com HW: 19.38, 19.50, 19.52

46. VIII) Free Energy Under Nonstandard Conditions • ΔG = ΔGº + RT ln Q • where ΔG = the free energy change under nonstandard conditions • R = 8.31 J/mol • T = the absolute temp. • Q = the reaction quotient Scificool.com

47. Ex) Calculate ΔG at 500. ºC for a reaction mixture that consists of 1.0 atm N2, 3.0 atm H2, and 0.50 atm NH3. Djurnal.wordpress.com

48. At equilibrium, ΔG = 0 (no preferred direction) and Q = K • Thus, ΔG = ΔGº + RT ln Q 0 = ΔGº + RT ln K ΔGº = -RT ln K If ΔGº = -, then K>1 If ΔGº = 0, then K=1 If ΔGº = +, then K<1 Cafepress.com HW: 19.58, 19.60

49. Solving for K: • K = e-ΔGº/RT • Ex) Find Kp at 500. ºC for the Haber process. Tgnfr.wordpress.com HW: 19.66

50. Ballyhooligan.com