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AP Chemistry Unit 7 – Thermodynamics

AP Chemistry Unit 7 – Thermodynamics. Lesson 7 – Entropy Book Sections: 19.1-19.3. First Law of Thermodynamics. You will recall from Chapter 5 that energy cannot be created nor destroyed. Therefore, the total energy of the universe is a constant.

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AP Chemistry Unit 7 – Thermodynamics

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  1. AP ChemistryUnit 7 – Thermodynamics Lesson 7 – Entropy Book Sections: 19.1-19.3

  2. First Law of Thermodynamics • You will recall from Chapter 5 that energy cannot be created nor destroyed. • Therefore, the total energy of the universe is a constant. • Energy can, however, be converted from one form to another or transferred from a system to the surroundings or vice versa.

  3. Spontaneous Processes • Spontaneous processes are those that can proceed without any outside intervention. • The gas in vessel B will spontaneously effuse into vessel A, but once the gas is in both vessels, it will not spontaneously return to vessel B.

  4. Spontaneous Processes Processes that are spontaneous in one direction are nonspontaneous in the reverse direction.

  5. Spontaneous Processes • Processes that are spontaneous at one temperature may be nonspontaneous at other temperatures. • Above 0 C it is spontaneous for ice to melt. • Below 0 C the reverse process is spontaneous.

  6. Reversible Processes In a reversible process the system changes in such a way that the system and surroundings can be put back in their original states by exactly reversing the process.

  7. Sample Problem • Predict whether the following processes are spontaneous as described, spontaneous in the reverse direction, or in equilibrium: • (a) When a piece of metal heated to 150 °C is added to water at 40 °C, the water gets hotter. • (b) Water at room temperature decomposes into H2 (g) and O2 (g), • (c) Benzene vapor, C6H6(g), at a pressure of 1 atm condenses to liquid benzene at the normal boiling point of benzene, 80.1 °C

  8. Entropy • Entropy (S) is a term coined by Rudolph Clausius in the 19th century. • Entropy can be thought of as a measure of the randomness of a system. • It is related to the motion in molecules.

  9. Entropy • Like total energy, E, and enthalpy, H, entropy is a state function. • Therefore, S = SfinalSinitial

  10. qrev T S = Entropy For a process occurring at constant temperature (an isothermal process), the change in entropy is equal to the heat that would be transferred if the process were reversible divided by the temperature: Positive entropy = more randomness Negative entropy = less randomness

  11. Sample Problem • The element mercury, Hg, is a silvery liquid at room temperature. The normal freezing point of mercury is –38.9 °C, and its molar enthalpy of fusion is ΔHfusion = 2.29 kJ/mol. What is the entropy change of the system when 50.0 g of Hg(l) freezes at the normal freezing point?.

  12. Entropy on the Molecular Scale • Ludwig Boltzmann described the concept of entropy on the molecular level. • Temperature is a measure of the average kinetic energy of the molecules in a sample.

  13. Entropy on the Molecular Scale • Molecules exhibit several types of motion: • Translational: Movement of the entire molecule from one place to another. • Vibrational: Periodic motion of atoms within a molecule. • Rotational: Rotation of the molecule on about an axis or rotation about  bonds.

  14. Entropy on the Molecular Scale • Boltzmann envisioned the motions of a sample of molecules at a particular instant in time. • This would be akin to taking a snapshot of all the molecules. • He referred to this sampling as a microstate of the thermodynamic system.

  15. Entropy on the Molecular Scale • Each thermodynamic state has a specific number of microstates, W, associated with it. • Entropy is S = k lnW where k is the Boltzmann constant, 1.38  1023 J/K.

  16. S= k ln lnWfinal lnWinitial Entropy on the Molecular Scale • The change in entropy for a process, then, is S = k lnWfinalk lnWinitial • Entropy increases with the number of microstates in the system.

  17. Entropy on the Molecular Scale • The number of microstates and, therefore, the entropy tends to increase with increases in • Temperature. • Volume. • The number of independently moving molecules.

  18. Sample Problem • Predict whether ΔS is positive or negative for each of the following processes, assuming each occurs at constant temperature: • H2O(l)  H2O(g) • Ag+(aq) + Cl-(aq)  AgCl(s) • 4 Fe(s) + 3 O2(g)  2 Fe2O3(s) • N2(g) + O2(g)  2 NO(g)

  19. Sample Problem • Predict which sample of matter has greater entropy: • 1 mol of NaCl(s) or 1 mol of HCl(g) • 2 mol of HCl(g) or 1 mol of HCl(g) • 1 mol of HCl(g) or 1 mol of Ar(g)Challenging!

  20. Entropy and Physical States • Entropy increases with the freedom of motion of molecules. • Therefore, S(g) > S(l) > S(s)

  21. Solutions Generally, when a solid is dissolved in a solvent, entropy increases.

  22. Entropy Changes • In general, entropy increases when • Gases are formed from liquids and solids; • Liquids or solutions are formed from solids; • The number of gas molecules increases; • The number of moles increases.

  23. HW: 8.66, 8.70, 8.72 (Read 19.1-19.3) • This week: • Tuesday: Entropy (19.1-19.3) • Wed-Fri: Midterm Exams • Ionic Equations Due Jan. 11 • Midterm Exam (Units 1-6): Jan. 14

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