1 / 52

Chapter 19 Chemical Thermodynamics

Chemistry, The Central Science , 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten. Chapter 19 Chemical Thermodynamics. John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc. Thermodynamics.

kimball
Download Presentation

Chapter 19 Chemical Thermodynamics

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 19Chemical Thermodynamics John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

  2. Thermodynamics • Thermodynamics is concerned with the question: can a reaction occur? • HW: Work problems in the chapter as we finish each section. For tomorrow 19.7 to 19.17 odd • BE PREPARE TO PRESENT YOUR EXAMPLES AT THE BEGINNING OF THE CLASS!!!

  3. First Law of Thermodynamics-Energy is conserved-. • 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.

  4. 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 separate again.

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

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

  7. 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.

  8. Irreversible Processes • Irreversible processes cannot be undone by exactly reversing the change to the system. • The system can go back but the surroundings are changed. • Spontaneous processes are irreversible.

  9. q T Entropy • Entropy (S) is a term coined by Rudolph Clausius in the 19th century. • Clausius was convinced of the significance of the ratio of heat delivered and the temperature at which it is delivered,

  10. Entropy • Entropy can be thought of as a measure of the randomness of a system. • It is related to the various modes of motion in molecules.

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

  12. 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:

  13. Second Law of Thermodynamics The total entropy of the universe increases for any spontaneous (irreversible) processes, and does not change for reversible processes.

  14. Second Law of Thermodynamics In other words: For reversible processes: Suniv = Ssystem + Ssurroundings = 0 For irreversible processes: Suniv = Ssystem + Ssurroundings > 0

  15. Entropy is not conserved: Suniv is increasing! • For a reversible process: Suniv = 0. • For a spontaneous process (i.e. irreversible): Suniv > 0. • Note: the second law states that the entropy of the universe must increase in a spontaneous process. It is possible for the entropy of a system to decrease as long as the entropy of the surroundings increases. • For an isolated system, Ssys = 0 for a reversible process and Ssys > 0 for a spontaneous process.

  16. Entropy • Suppose a system changes reversibly between state 1 and state 2. Then, the change in entropy is given by • at constant T where qrev is the amount of heat added reversibly to the system. (Example: a phase change occurs at constant T with the reversible addition of heat.)

  17. Calculating entropies for phase changes • Q rev= D H • Then D S = D H / T

  18. Entropy and the Second Law of Thermodynamics • The Spontaneous Expansion of a Gas • Why do spontaneous processes occur? • Consider an initial state: two flasks connected by a closed stopcock. One flask is evacuated and the other contains 1 atm of gas. • The final state: two flasks connected by an open stopcock. Each flask contains gas at 0.5 atm. • The expansion of the gas is isothermal (i.e. constant temperature). Therefore the gas does no work and heat is not transferred.

  19. The Spontaneous Expansion of a Gas • Why does the gas expand?

  20. The Spontaneous Expansion of a Gas • Consider the simple case where there are two gas molecules in the flasks. • Before the stopcock is open, both gas molecules will be in one flask. • Once the stopcock is open, there is a higher probability that one molecule will be in each flask that both molecules being in the same flask.

  21. The Spontaneous Expansion of a Gas • When there are many molecules, it is much more probable that the molecules will distribute among the two flasks than all remain in only one flask.

  22. Entropy on the Molecular Scale • Ludwig Boltzmann described the concept of entropy on the molecular level. He used statistical thermodynamics – statistics and probability are used to link macro and micro world. • Temperature is a measure of the average kinetic energy of the molecules in a sample.

  23. 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.

  24. 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.

  25. 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.

  26. S= k ln Wfinal Winitial 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.

  27. 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.

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

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

  30. There is a balance between energy and entropy considerations. • When an ionic solid is placed in water two things happen: • the water organizes into hydrates about the ions (so the entropy decreases), and • the ions in the crystal dissociate (the hydrated ions are less ordered than the crystal, so the entropy increases).

  31. The Molecular Interpretation of Entropy • A gas is less ordered than a liquid that is less ordered than a solid. • Aqueous ions are less ordered than pure solids and liquids, but more ordered than gases • Any process that increases the number of gas molecules leads to an increase in entropy. • When NO(g) reacts with O2(g) to form NO2(g), the total number of gas molecules decreases, and the entropy decreases.

  32. 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.

  33. Examples – Determine the sign of ΔS for each of the following: • Na (s) + ½ Cl2 (g)  NaCl (s) • N2 (g) + 3 H2 (g)  2 NH3 (g) • 2 H2 (g) + O2 (g)  2 H2O (l) • H2O (l)  H2O (g) • NaCl (s)  Na+ (aq) + Cl- (aq)

  34. Third Law of Thermodynamics The entropy of a pure crystalline substance at absolute zero is 0.

  35. Energy is required to get a molecule to translate, vibrate or rotate. • The more energy stored in translation, vibration and rotation, the greater the degrees of freedom and the higher the entropy. • In a perfect crystal at 0 K there is no translation, rotation or vibration of molecules. Therefore, this is a state of perfect order (zero entropy).

  36. Boiling corresponds to a much greater change in entropy than melting. • Entropy will increase when • liquids or solutions are formed from solids, • gases are formed from solids or liquids, • the number of gas molecules increase, • the is temperature increased.

  37. Entropy Changes in Chemical Reactions • Absolute entropy can be determined from complicated measurements. • Standard molar entropy, S: entropy of a substance in its standard state. Similar in concept to H. • Units: J mol-1 K-1. Note units of H: kJ mol-1. • Standard molar entropies of elements are not zero. • For a chemical reaction which produces n moles of products from m moles of reactants:

  38. Standard Entropies • These are molar entropy values of substances in their standard states. • Standard entropies tend to increase with increasing molar mass.

  39. Standard Entropies Larger and more complex molecules have greater entropies. More atoms in the molecule allow for a greater degree of freedom.

  40. Examples – Calculate ΔS for each of the following reactions: • CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (g) • N2 (g) + 3 H2 (g)  2 NH3 (g) • 2 SO3 (g)  2 SO2 (g) + O2 (g) • HCl (g)  H+ (aq) + Cl- (aq)

  41. qsys T Ssurr = Entropy Changes in Surroundings • Heat that flows into or out of the system changes the entropy of the surroundings. • For an isothermal process: • At constant pressure, qsys is simply H for the system.

  42. Entropy Change in the Universe • The universe is composed of the system and the surroundings. • Therefore, Suniverse = Ssystem + Ssurroundings • For spontaneous processes Suniverse > 0

  43. Hsystem T Entropy Change in the Universe • This becomes: Suniverse = Ssystem + Multiplying both sides by T, TSuniverse = Hsystem  TSsystem

  44. Gibbs Free Energy • TDSuniverse is defined as the Gibbs free energy, G. • When Suniverse is positive, G is negative. • Therefore, when G is negative, a process is spontaneous.

  45. Gibbs Free Energy • If DG is negative, the forward reaction is spontaneous. • If DG is 0, the system is at equilibrium. • If G is positive, the reaction is spontaneous in the reverse direction.

  46. DG = SnDG(products)  SmG(reactants) f f Standard Free Energy Changes Analogous to standard enthalpies of formation are standard free energies of formation, G. f where n and m are the stoichiometric coefficients.

  47. Free Energy Changes At temperatures other than 25°C, DG° = DH  TS How does G change with temperature?

  48. There are two parts to the free energy equation: H— the enthalpy term TS — the entropy term The temperature dependence of free energy, then comes from the entropy term. Free Energy and Temperature

  49. Free Energy and Temperature

More Related