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Introduction to Thermodynamics

Introduction to Thermodynamics. Conservation of Energy First Law of Thermodynamics = energy cannot be created or destroyed, only converted between different forms Example: CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O + energy Reaction gives off energy as heat

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Introduction to Thermodynamics

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  1. Introduction to Thermodynamics • Conservation of Energy • First Law of Thermodynamics = energy cannot be created or destroyed, only converted between different forms • Example: CH4(g) + 2O2(g) CO2(g) + 2H2O + energy • Reaction gives off energy as heat • Potential energy stored in chemical bonds is lowered • Total energy is unchanged • Uses and Shortcomings • Lets us keep track of energy flow in processes • Does not tell us if or why a given process occurs • Does not tell us direction of a chemical reaction

  2. Entropy • The “Heat Tax” • Conversion of energy between forms is inefficient: • Usually, some energy is lost as heat • The fewer energy conversion, the better

  3. Direction of Processes and Reactions • Examples: • Ball at the top of a hill Ball at the bottom of a hill • Steel + H2O + O2 Rust • Gas in one part of a container Gas filling a container • Ice at 5 oC Water at 5 oC • What is in common? • Exothermic? Not ice melting or gas expanding • Increased Disorder = Increased Entropy = +DS • Entropy = S = driving force of spontaneous reactions = disorder or random • Probability (likelihood): there are many ways for objects/molecules to be disordered, but only a few to be ordered • Nature proceeds towards the most likely state = state with greatest number of energetically equivalent arrangements • Expansion of a Gas

  4. Possible arrangements of 4 gas molecules in a 2-bulb system • Microstates = possible configurations of a particular arrangement • Entropy selects most likely arrangement = 2 molecules in each bulb Microstates Order 1 microstate Ordered 4 microstates Somewhat Disordered 6 microstates Fully disordered ABCD --- ABC D ABD C ACD B BCD A AB CD CD AB AC BD BD AC AD BC BC AD B

  5. Probabilities

  6. Positional Entropy = entropy depending on configuration in space • Changes in state depend on positional entropy • Ssolid < Sliquid << Sgas c) Larger volume allows many more available positions for particles d) Dissolving a solid provides more volume for particles to occupy e) Example: What has highest positional entropy • 1 mole solid CO2 or 1 mole of gaseous CO2? • 1 mole N2 at 1 atm or 1 mole of N2 at 0.01 atm?

  7. Example: Predict the sign of the entropy change for • Dissolving solid sugar into water • Iodine vapor condensing to crystals on a surface 7) Second Law of Thermodynamics = in any spontaneous process, there is always an increase in the entropy of the universe • Energy is conserved = constant • Entropy is always increasing • DSuniverse = DSsystem + DSsurroundings • For a given process: if DSuniverse = + the process is spontaneous if DSuniverse = - the process is not spontaneous • Life = constant battle against entropy i) Large molecules are assembled from smaller ones • Organizing a cell is DSsystem = - the process is not spontaneous • Fortunately, it is DSuniverse that must be positive in a process iii) DSsurroundings = large + for life to occur

  8. Temperature and Spontaneity • Change in state: 1 mol = 18 ml H2O(l) 1 mol = 31 L H2O(g) • DSsurrounding depends on flow of heat into or out of the system • Heat increases the motion (randomness) of particles • Exothermic reactions release heat to surroundings DSsurrounding = + • Endothermic reactions absorb heat from surroundings DSsurrounding = - • Vaporization of water is endothermic DSsurrounding = - • DSuniverse = DSsystem + DSsurrounding = (+) + (-) = +/- ? • Depends on the temperature • If T > 100 oC, DSuniverse = + If T < 100 oC, DSuniverse = - • Temperature Effects • DSsurroundings depends on heat flow • Exothermic reactions usually favors spontaneity • Spontaneity usually lowers the energy of the starting material as it becomes product • The difference of these energies = heat released to surroundings

  9. Importance of Exothermicity of DSuniverse depends on Temperature • Adding heat to hot surroundings has little effect • Adding heat to cold surroundings has a large effect • Heat transfer is more important at low temperatures • In Summary • Sign of DSsurr depends on direction of heat transfer • Magnitude of DSsurr depends on T • The (-) is there because DH is for the system, which is opposite of DH of the surroundings • Example: Find DSsurr at 25 oC • Sb2S3(s) + 3Fe(s) 2Sb(s) + 3FeS(s) DH = -125 kJ/mol • Sb4O6(s) + 6C(s) 4Sb(s) + 6CO(g) DH = +778 kJ/mol

  10. Free Energy • Free Energy = G = H – TS • DGprocess = DH – TDS • Divide by –T • A process is spontaneous if DG = - • Chemists use DG rather than DS because we only need to know system • Example: Predicting Spontaneity using DG 1) H2O(s) H2O(l) DHo = 6030 J/mol, DSo = 22.1 J/K mol

  11. Classifying Processes/Reactions based on DH and DS 3) Example: At what T is Br2(l) Br2(g) spontaneous (1 atm) given that DH = 31.0 kJ/mol and DS = 93.0 J/Kmol • Spontaneous when DG = - • Set DG = 0 and solve for T • When T > 333K, TDS > DH and DG = - (Entropy controlled) • When T < 333K, TDS < DH and DG = + (Enthalpy controlled) • 333K is the boiling point of Br2(l)

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