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For each pair, which state is more stable? Pencil on desk vs. raised in the air

THERMODYNAMICS Review of Energy and Enthalpy Changes (Ch. 5). For each pair, which state is more stable? Pencil on desk vs. raised in the air Skier at top of mountain vs. bottom Toaster at 25 o C vs. hot toaster Liquid water at 25 o C vs. ice at 25 o C

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For each pair, which state is more stable? Pencil on desk vs. raised in the air

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  1. THERMODYNAMICS Review of Energy and Enthalpy Changes (Ch. 5) • For each pair, which state is more stable? • Pencil on desk vs. raised in the air • Skier at top of mountain vs. bottom • Toaster at 25oC vs. hot toaster • Liquid water at 25oC vs. ice at 25oC • Salt water vs. pure water in contact with solid salt • MgO vs. Mg metal in contact with O2 • In each case, how was energy transferred in going to the more stable state?

  2. Energy Changes: Heat and Work • Heat = q = energy transferred due to a difference in temperature. • +q means heat is added to the system • Work = w = action of force through a distance (often P∆V) • +w means work is done on the system • Find one pair in which energy was transferred as heat. • Find one pair in which energy was transferred as work. What force was involved? • Find a second pair in which a different force was involved.

  3. STATE FUNCTIONS Astate function is a quantity that only does not depend on the process by which the system was prepared Example:Youraltitude(height above sea level) does not depend on the route you took to class this morning. State functions are written as uppercase letters (E, H, P, V, T, S…) Changes in state functions are path-independent: q and w are not state functions but ∆E (= q + w) is a state function reactants 1 ∆E 2 products

  4. Energy Changes • ∆E = Efinal state - Einitial state • ∆E (kJ/mol) • O2 (g)  2 O atoms +498.3 • Water  ice at 25oC -6.0 • Water + salt  salty water -3.9 • Si (s) + O2 (g) SiO2 (s) -908 • Which set of reactants/products has the biggest difference in energy? • Which set has the lowest absolute energy? • If we assign O atoms an energy of zero, what is the energy of O2? • Is the lowest energy state always most stable?

  5. First Law of Thermodynamics ∆Euniverse = 0 Energy is conserved ∆Esystem + ∆Esurroundings = ∆Euniverse so ∆Esystem = - ∆Esurroundings Heat and work: ∆Esystem = q + w and for PV work at const. pressure, ∆Esystem = q – P∆V

  6. Energy Changes - Heat and Work ∆ ∆ Dry ice (CO2) is heated to room temperature at const. P or at const. V • What are the signs of q and w for each? • Does Esystem increase or decrease? How do you know? • Which system has higher E at the end?

  7. When natural gas burns: • CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l) • Reaction produces heat and light • Is energy conserved? • What are the signs of q and w for the overall reaction? • Where does the energy come from? • Is energy stored or released when bonds are broken?

  8. Enthalpy (H) • ∆H = heat transferred at const. P • If (+), endothermic (need to add heat) • If (-), exothermic (heat is given off) • Classify as endo- or exothermic: • Ice melting • Water boiling • Wood burning • H is a state function – changes are path-independent • H = E + PV (sums and products of state functions are also state functions)

  9. Standard Enthalpy of Formation • ∆Hof: • ∆H for making a compound from elements in their standard states • Standard state is the most stable form (pure solid, pure liquid, or gas at P = 1 atm) • For solutes in solution, standard state is usually 1 M • There are tables of ∆Hof • ∆Horxn =  ∆Hof(products) –  ∆Hof(reactants)

  10. Standard Enthalpy of Formation of Oxides Reaction∆Hof (kJ/mol) 2 Li + 1/2 O2 Li2O -598 2 Na + 1/2 O2 Na2O -414 Ca + 1/2 O2 CaO -635 Mg + 1/2 O2 MgO -602 2 Al + 3/2 O2 Al2O3 -1676 2 Fe + 3/2 O2 Fe2O3 -824 2 Cr + 3/2 O2 Cr2O3 -1128 2 Ag + 1/2 O2 Ag2O -30 2 Cu + 1/2 O2Cu2O -167 C + O2 CO2 -394 1/2 N2 + O2 NO2 +34 Cl2 + 1/2 O2Cl2O +80

  11. What are the least and most stable oxides in the table? What periodic trends do you see for the ∆Hof’s of the oxides? How do oxidation states relate to ∆Hof? Predict ∆Hof for Au2O, K2O, and SrO. Why does Mg burn in dry ice?

  12. SPONTANEOUS REACTIONS A spontaneous reaction is one that can proceed in the forward direction under a given set of conditions. Note: spontaneity has nothing to do with the rate at which a reaction occurs. CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g) Spontaneous CO2(g)  C(s) + O2(g) Not Spontaneous 2 Fe2O3(s)  4 Fe(s) + 3 O2(g) Not Spontaneous What determines whether these reactions are spontaneous?

  13. SPONTANEITY AND WORK Useful work can be extracted from a spontaneous process 2 H2 + O2 2 H2O can be used to drive a rocket Water in a tower  Water on the ground Can do work (drive a turbine) Work must be done to drive a non-spontaneous process 2 H2O  2 H2 + O2 Energy (electrolysis work) must be put in. Water on the ground  Water in a tower Work is done to pump water

  14. Spontaneity and H • Not all spontaneous reactions are exothermic • Examples: • At +10°C H2O(s)  H2O(l) H > 0 • Ba(OH)2•8H2O(s) + 2NH4SCN(s)  • Ba(SCN)2(aq) + 2NH3(aq) + 10H2O(l) H > 0 NH4Cl(s) + H2O(l)  NH4Cl(aq) H > 0

  15. QUESTIONS • Why are some endothermic processes spontaneous? • Evaporation of water • Dissolution of NH4+NO3- • What makes an ideal gas expand into a vacuum • (q = 0, w = 0)? • What distinguishes "past" from "future" in chemistry and physics?

  16. ENTROPY A thermodynamic parameter that measures the disorder or randomness in a system. The more disordered a system, the greater its entropy. Entropy is a state function -- its value depends only upon the state of the system (not how it got there). We are usually concerned with the change in entropy (S ) during a process such as a chemical reaction. S = S final - S initial

  17. Which "reactions" have S > 0? Cards arranged in order  Random order after shuffling the deck Messy dorm room  Cleaned up room CO2 + H2O + Minerals  Tree 1 mole of gas  1 mole of gas in a 1 L flask in a 2 L flask NH4Cl(s)  NH4+(aq) + Cl-(aq)

  18. ENTROPY Gases have a lot more entropy than solids or liquids. Reactions that form gases usually have S > 0 S (+ or - ?) H2O (l, 25oC)  H2O(g) CaCO3 (s)  CaO(s) + CO2(g) N2(g) + 3 H2(g)  2 NH3(g) N2(g) + O2(g)  2 NO(g) Au(s) at 298K  Au(s) at 1000K Ag+(aq) + Cl-(aq)  AgCl(s)

  19. Adding heat increases entropy Entropy and heat: S = qrev/T e.g., for melting ice, S = Hfus/273 K (endothermic)  + + m.p.= 0oC Entropy units: J/mol-K (same units as specific heat) S(1 mole HCl(g)) >> S(1 mole NaCl(s)) why? S(2 moles HCl(g)) = 2 S(1 mole HCl(g)) “ S(1 mole HCl(g)) > S(1 mole Ar(g)) “ S(1 mole N2(g) at 300K) > S(1 mole N2(g) at 200K) “

  20. ENTROPY and PHASE CHANGES Sgas >> Sliq > Ssolid Svap = qrev = Hvap T T During phase changes temperature and pressure are constant. (Heat transfer is reversible, so H = q = qrev) Calculate the entropy change when 1 mole of liquid water evaporates at 100oC (Hvap = +44 kJ/mol)

  21. 2nd Law of Thermodynamics: The total entropy in the universe is increasing. Suniverse > 0 Suniverse = Ssystem + Ssurroundings > 0 3rd Law: The entropy of every pure substance at 0K (absolute zero temperature) is zero. S = 0 at 0 K

  22. 3rd LAW ENTROPY Entropy is a state function - its value depends only on the initial and final states. And S = 0 at T = 0 K (3rd Law) This means we can measure absolute entropy S (not just ∆S) So (rxn) = So (products) - So (reactants) N2(g) + 3 H2(g) 2 NH3(g) So (25oC) = 191.5 130.58 192.5 J/mol-K So (rxn) = (2x192.5) - (191.5 + 3x130.58) = -198.3 J/mol-K

  23. CRITERIA FOR SPONTANEITY Hsystem < 0 Exothermic reactions are usually spontaneous. S Reactions are always spontaneous if Ssystem + Ssurroundings > 0 (2nd Law) We need a new state function (G) that can predict spontaneity from just the system. This is called the Free Energy: G = H - TS (T = absolute temperature (in K)) Gsystem predicts rxns at constant T and P: G < 0 Spontaneous G > 0 Not spontaneous G = 0 Reaction at equilibrium

  24. EFFECT OF TEMPERATURE ON SPONTANEITY G = H - TS H and S do not change much with temperature, but ∆G does. H S GSpontaneous? - + - - + + + -

  25. Is this reaction spontaneous at room temp? At 1100 oC? CaCO3(s)  CaO(s) + CO2(g) So92.88 39.75 213.6 J/molK  Hfo1207.1 635.3 393.5 (kJ/mol) Calculate the boiling point of bromine Hovap = 31.0 (kJ/mol) Sovap = 92.9 (J/mol-K)

  26. STANDARD FREE ENERGY OF FORMATION Gf = Free energy change in forming one mole of a compound from its elements, each in their standard states. Standard States: Solid Pure solid Liquid Pure liquid Gas P = 1 atm Solution 1M solution Elements Gf = 0 Temperature Usually 25C Grxn = Gf(prods) -Gf(reactants) Units:kJ/mole.

  27. COMBUSTION OF SUCROSE • C12H22O11(s) + 8 KClO3(s) 12 CO2(g) + 11 H2O(l) + 8 KCl(s) • Is this reaction spontaneous? • ∆Gof(sucrose) = -1544.3 kJ/mol • ∆Gof(KClO3, s) = -289.9 ” • ∆Gof(CO2, g) = -394.4 " • ∆Gof(H2O, l) = -237.1 " • ∆Gof(KCl, s) = -408.3 " • How does H2SO4 affect the reaction rate? • What are two possible reasons for using high T to carry out a reaction?

  28. EXTENT OF REACTIONS • So far we have considered only standard conditions and ∆Ho, ∆So, ∆Go. • What happens under other conditions? • G = G + RTlnQ = G + 2.303RTlog10Q • aA + bB  cC + dD Q = ([C]c[D]d) = REACTION • ([A]a[B]b) QUOTIENT • Example: 2NO2 (g)  N2O4 (g) Go298 = 5.4 kJ/mol • Partial pressure of NO2 is 0.25 atm and partial pressure of N2O4 is 0.6 atm. What is ∆G?

  29. RELATIONSHIP BETWEEN G AND G How do concentrations affect Q and G? QG ( or ) Add more reactants Take away products Take away reactants

  30. ANALOGY BETWEEN POTENTIAL ENERGY AND FREE ENERGY Go slope = 0 At equilibrium point, G = 0 for interconverting reactants  products Note: This does NOT mean Go = 0

  31. Nitrogen Fixation N2(g) + 3 H2(g) = 2 NH3(g) ∆Gof: 0 0 -16.7 kJ/mol ∆Gorxn = What is the sign of ∆So? ∆Ho? Would higher or lower P favor products? " " " " T " " ?

  32. RELATIONSHIP BETWEEN • GAND Keq • At equilibrium, G = 0: • G = 0 = G + 2.303RTlogQ • = G + 2.303RTlogKeq • G = -2.303RTlogKeq • G > 0 Keq < 1 • G < 0 Keq > 1 • G = 0 Keq = 1

  33. Solubility Equilibrium Guess the solubilites of these salts in water: KClO3(s)  K+(aq) + ClO3-(aq) ∆Go > 0 NaF(s)  Na+(aq) + F-(aq) ∆Go ≈ 0 NaCl(s)  Na+(aq) + Cl-(aq) ∆Go < 0 What is the solubility product Ksp of AgBr? AgBr(s) Ag+(aq) + Br(aq) Gfo Ag+ 77.1 kJ/mol Gfo Br104 kJ/mol Gfo AgBr 96.9 kJ/mol

  34. Spontaneity and Equilibrium At 10°C H2O(l)  H2O(s)Spontaneous Can get the system to do work! At + 10°C H2O(l)  H2O(s) Non-Spontaneous Must do work to get this to go. But at 0°C H2O(l) H2O(s)Equilibrium Heat in and out is reversible q = qrev When a chemical system is at equilibrium, reactants and products can interconvert reversibly.

  35. RELATIONSHIP BETWEEN G AND WORK For a spontaneous process, G = Wmax= The maximum work that can be obtained from a process at constant T and P. For a non-spontaneous process, G = Wmin = The minimum work that must be done to make a process go at constant T and P.

  36. Calculating W from G What is the maximum work available from the oxidation of 1 mole of octane under standard conditions (P = 1 atm)? C8H18(l) + 12.5 O2(g) 8CO2(g) + 9H2O(l) Gfo 17.3 0 394.4 237.1 kJ/mol

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