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Thermodynamics

Thermodynamics. First Law of Thermodynamics. Energy is always conserved However, it can transfer between system and surrounding. ∆U = q + w Internal energy = KE + PE. Change in internal energy. work. Heat exchanged. What governs the spontaneity of a reaction?

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Thermodynamics

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  1. Thermodynamics

  2. First Law of Thermodynamics • Energy is always conserved However, it can transfer between system and surrounding. ∆U = q + w Internal energy = KE + PE Change in internal energy work Heat exchanged

  3. What governs the spontaneity of a reaction? Spontaneous reactions happen without outside intervention and have a definite direction (not reversible)

  4. Internal warmth Exothermic reactions have ΔH<0 (neg) Products have less energy than reactants Energy is released Enthalpy (H)

  5. Endothermic reactions take in heat. • Products have more energy than reactants. • ΔH>0 (pos)

  6. Second law of thermodynamics In an isolated system, natural processes are spontaneous when they lead to an increase in disorder, or entropy.

  7. Entropy (S): the disorder in a system • Reactions can occur spontaneously when a state of less order is achieved. • If ΔS is positive, disorder is increasing

  8. Lower entropy Higher entropy

  9. Predict the sign of ΔS • H2O (l) g H2O (g) • Ag+(aq) + Cl-(aq) g AgCl(s) • 4Fe(s) + 3O2(g) g 2Fe2O3(s) • N2(g) + O2(g) g 2NO(g) + - - Near 0

  10. Driving forces for chemical reactions • Toward lower enthalpy or potential energy • Maximum stability • Toward higher entropy • Maximum randomness

  11. Enthalpy of Formation (ΔHf) • The change in enthalpy when one mole of a compound is produced from the free elements. (kJ/mol) • If ΔHfis negative, produced by exothermic rx. “thermodynamically stable”, since much energy is required to decompose it • If ΔHfis positive, produced by endothermic rx and gives off energy when decomposing See Table A-6 p.859

  12. ΔHr = ∑ ΔHf (products) - ∑ ΔHf (reactants)

  13. Gibbs Free Energy (G) • ΔG = ΔH - T· ΔS • ΔG is the maximum amount of energy available from any chemical reaction. • If ΔG is negative, the reaction occurs spontaneously. “exergonic”-gives off work • If ΔG is positive, the reaction is not spontaneous. “endergonic”- takes in work • If ΔG is 0, it’s at equilibrium

  14. ΔG = ΔH - T· ΔS

  15. ΔGr = ∑ΔGf (products) - ∑ΔGf (reactants) Calculate the change in entropy and in free energy for the following reaction. CH4(g) + 2O2(g) g CO2(g) + 2H2O(l)

  16. ΔGf = -817.90 kJ • ΔS = -242.98 J/K

  17. Hess’s Law • The enthalpy change for a reaction is the sum of the enthalpy changes for all steps of the reaction. • The two-step reactions are: • C + 1/2 O2g CO,         ΔH° = -110 kJ/molCO + 1/2 O2g CO2,     ΔH° = -283 kJ/mol. • Adding the two equations together and cancel out the intermediate, CO, on both sides leads to • C + O2g CO2, ΔH° = (-110)+(-283)= -393kJ/mol.

  18. BaO(s) + H2SO4(l) g BaSO4(s) + H2O(l) Calculate the enthalpy of the above reaction from the following data. SO3(g) + H2O(l) g H2SO4 (l) ΔH° = -78.2 kJ BaO(s) + SO3(g) g BaSO4(s) ΔH° = -213.4 kJ H2SO4 (l) g SO3(g) + H2O(l) ΔH° = 78.2 kJ +BaO(s) + SO3(g) g BaSO4(s) ΔH° = -213.4 kJ H2SO4 + BaO g BaSO4 + H2O ΔHr°=-135.2kJ

  19. To get the right overall reaction. . . • Add them as is • Flip one and reverse the sign of the ΔH° • Multiply one by a factor. Multiply the ΔH° by the same factor.

  20. State function • The value depends only on the current state of the system, not on how it got there. • ΔG, ΔH, ΔS, and ΔE are state functions.

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