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Chapter 17 Thermodynamics

Chapter 17 Thermodynamics. 17.1 – Spontaneous Processes and Entropy. First Law of Thermodynamics: The energy of the universe is constant Energy can be neither created nor destroyed (but it can be transferred)

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Chapter 17 Thermodynamics

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  1. Chapter 17Thermodynamics

  2. 17.1 – Spontaneous Processes and Entropy • First Law of Thermodynamics: • The energy of the universe is constant • Energy can be neither created nor destroyed (but it can be transferred) • In physics, if you hold a ball on the roof of the school, it has a lot of potential energy due to its height. Then you drop it. As it falls, it loses potential energy and gains kinetic energy.

  3. 17.1 – Spontaneous Processes and Entropy • In chemistry, if you burn petrol in your car (assume petrol = octane = C8H18), you convert the chemical potential energy stored in the bonds of octane into mechanical energy, which drives your car. It is an exothermic reaction 2C8H18 + 25O2 16CO2 + 18H2O ΔE < 0

  4. 17.1 – Spontaneous Processes and Entropy 2C8H18 + 25O2 Energy ΔE < 0 16CO2 + 18H2O

  5. 17.1 – Spontaneous Processes and Entropy • An exothermic reaction does not ensure that the octane and oxygen react spontaneously. (If they did, we would be in trouble!) • A reaction that occurs spontaneously does not mean that it occurs quickly. Spontaneous means that it occurs without outside intervention.

  6. 17.1 – Spontaneous Processes and Entropy • Kinetics answers “how fast will this reaction go?” • Depends on activation energy, temperature, concentration of reactants, and the presence of a catalyst • Thermodynamics answers “will this reaction go?” • Depends on … we will discover in this chapter

  7. 17.1 – Spontaneous Processes and Entropy

  8. 17.1 – Spontaneous Processes and Entropy • Entropy, S • A measure of disorder, or randomness • A spontaneous process increases the entropy of the universe • Naturally, “things” tend towards disorder. • Cannot be measured directly, rather ΔS is used • Can be used to predict spontaneity of a reaction • Usually measured in J/K

  9. 17.1 – Spontaneous Processes and Entropy • Entropy, S • “Any method involving the notion of entropy, the very existence of which depends on the second law of thermodynamics, will doubtless seem to many far-fetched and may repel beginners as obscure and difficult of comprehension.” - Josiah Willard Gibbs (1873).

  10. 17.1 – Spontaneous Processes and Entropy • Entropy, S • It describes the number of arrangements (positions or energy levels) that are available to a system existing in a given state. Hence, it is closely related to probability. • The more ways a particular state can be achieved, the more likely you are to find the system in that state. Nature spontaneously proceeds towards states that have higher probabilities of existing

  11. 17.1 – Spontaneous Processes and Entropy • An example attempting to simplify entropy • Take four O2 molecules, O2A, O2B, O2C, and O2D. Put them in the following vessel: vacuum

  12. 17.1 – Spontaneous Processes and Entropy • An example attempting to simplify entropy • We know that if we were to open the valve, the gas molecules would spontaneously move to the vacuum vacuum

  13. 17.1 – Spontaneous Processes and Entropy • An example attempting to simplify entropy • But how many would move? Would it be possible to take a snapshot where the arrangement is 2-2? 1-3? 3-1? 0-4? 4-0? vacuum

  14. 17.1 – Spontaneous Processes and Entropy • An example attempting to simplify entropy • Let’s call each of the possibilities in each arrangement a microstate. How many microstates are possible in each arrangement? For example, there is only 1 microstate in the 4-0 arrangement: A B vacuum C D

  15. 17.1 – Spontaneous Processes and Entropy • An example attempting to simplify entropy • Another microstate exists in the 0-4 arrangement: A B C D

  16. 17.1 – Spontaneous Processes and Entropy • An example attempting to simplify entropy • What about the 3-1 arrangement? There are 4 possible microstates: • The same is true of the 1-3 arrangement A A D B B C C D A D A D B B C C

  17. 17.1 – Spontaneous Processes and Entropy • An example attempting to simplify entropy • There are 6 microstates in the 2-2 arrangement: C C A A D B B D B A A D B C C D B C A D A D B C

  18. 17.1 – Spontaneous Processes and Entropy • An example attempting to simplify entropy • The probability of finding all of the gas molecules in one bulb is very low, compared to the even distribution. • We say that if the probability of a particular microstate is low, the event will not occur spontaneously • This probability is called positional probability

  19. 17.1 – Spontaneous Processes and Entropy • Entropy of different phases • There are many possibilities for the positions of molecules in 1 mole of gas. More so than in a liquid, or a solid. In general:

  20. 17.1 – Spontaneous Processes and Entropy • Entropy at different volumes/pressures • If you increase the volume of a gas (or decrease the pressure), more microstates become available to the gas.

  21. 17.1 – Spontaneous Processes and Entropy • Entropy in solution formation • When you mix two pure substances, there become many more microstates available for the solution than separated. Therefore the entropy change is expected to be positive.

  22. 17.1 – Spontaneous Processes and Entropy • Predicting the sign of ΔS • Example, what is the sign of the entropy change in the following processes? • A. Solid sugar is added to water to form a solution • B. Iodine vapor condenses on a cold surface to form iodine crystals

  23. 17.2 – Entropy and the Second Law of Thermodynamics • Second Law of Thermodynamics: • In any spontaneous process, the entropy of the universe increases • In the universe, energy is conserved, but the entropy is not – it is constantly increasing.

  24. 17.2 – Entropy and the Second Law of Thermodynamics • Recall from Chapter 6 • The system is the part of the universe on which we are focusing • The surroundings are everything else – the vessel, the table, the air molecules in the lab, etc…

  25. 17.2 – Entropy and the Second Law of Thermodynamics • Second Law of Thermodynamics: ΔSuniverse = ΔSsystem + ΔSsurroundings • If ΔSuniverse > 0, then entropy increases and the process occurs spontaneously • If ΔSuniverse< 0, then the reverse process is spontaneous • If ΔSuniverse= 0 , the process, or its reverse process, has no tendency to occur. The system is at equilibrium

  26. 17.2 – Entropy and the Second Law of Thermodynamics • Second Law of Thermodynamics: • How does this apply to living cells, where large molecules are spontaneously assembled from smaller ones? • A process for which ΔSsystem is negative can be spontaneous if the associated ΔSsurroundings is larger, and positive. i.e., ΔSsurroundings > 0 > ΔSsystemleads to a ΔSuniverse > 0

  27. 17.3 – The Effect of Temperature on Spontaneity • Water example on spontaneity • We know that at atmospheric pressure, water spontaneously evaporates (liquid  gas) at temperature greater than 100oC. Under these conditions, entropy increases: Summary: ΔSuniverse > 0 ΔSsystem > 0 ΔSsurroundings < 0 (vaporizing is endothermic)

  28. 17.3 – The Effect of Temperature on Spontaneity • Water example on spontaneity • At temperatures lower than 100oC, water vapor spontaneously condenses to liquid water. Under these conditions, entropy of the system decreases. Summary: ΔSuniverse > 0 ΔSsystem < 0 ΔSsurroundings > 0 (condensing is exothermic)

  29. 17.3 – The Effect of Temperature on Spontaneity • Water example on spontaneity • In these two scenarios, the temperature is the only variable that has changed. This has resulted in a difference in the ΔSsurroundings • Entropy changes in the surroundings are primarily determined by heat flow.

  30. 17.3 – The Effect of Temperature on Spontaneity • Temperature • An exothermic process in the system increases the entropy (by  kinetic energy) of the surroundings • The magnitude of ΔSsurroundings depends on the temperature at which the heat is transferred. If the temperature of the surroundings is high, the heat transferred will not make a big difference in the kinetic energy of the surrounding molecules. If temperature is low, the heat transfer has a bigger impact

  31. 17.3 – The Effect of Temperature on Spontaneity • Factors affecting ΔSsurroundings • The sign of ΔSsurroundings depends on whether the reaction in the system is exothermic (ΔSsurr > 0) or endothermic (ΔSsurr < 0) • The magnitude of ΔSsurroundingsdepends on the temperature. The impact of the transfer of a given quantity of energy as heat to or from the surroundings will be greater at lower temperatures

  32. 17.3 – The Effect of Temperature on Spontaneity • Summary: • Exothermic process: • Endothermic process:

  33. 17.3 – The Effect of Temperature on Spontaneity • Summary: • At constant pressure, quantity of heat (q) is equal to enthalpy, ΔH. Taking into account the sign of enthalpy (+ve for endothermic, -ve for exothermic), we can summarize these results as:

  34. 17.3 – The Effect of Temperature on Spontaneity • Summary:

  35. 17.3 – The Effect of Temperature on Spontaneity • Example,

  36. 17.4 – Free Energy • (Gibbs) Free Energy, G • Named for Josiah Willard Gibbs • Is useful in dealing with the temperature dependence on spontaneity • At constant temperature and pressure • Since there are no subscripts, these refer to the thermodynamic state of the system.

  37. 17.4 – Free Energy • (Gibbs) Free Energy, G • The true determining factor for spontaneity: • If ΔG > 0, reaction is not spontaneous • If ΔG < 0, reaction is spontaneous • If ΔG = 0, reaction is at equilibrium

  38. 17.4 – Free Energy • (Gibbs) Free Energy, G • Divide both sides by –T, • Recall that at constant T and P, ΔSsurr = –ΔH/T, • Recall that ΔSuni = ΔSsurr+ ΔSsys

  39. 17.4 – Free Energy • (Gibbs) Free Energy, G • This equation tells us that if a process is carried out at constant T and P, the reaction will be spontaneous if ΔG is negative (recall a spontaneous process has a positive ΔSuni)

  40. 17.4 – Free Energy • (Gibbs) Free Energy, G • Recall that the degree symbol (o) means that all substances are in their standard states • There are standard tables that show all of these thermodynamic properties for all substances. They are used the same way we used ΔHformationin chapter 6

  41. 17.4 – Free Energy • (Gibbs) Free Energy, G • Example • A vessel contains 1 mole of liquid dinitrogentetraoxide (N2O4). Its boiling point is at 21oC. N2O4(l) ⇌ N2O4(g) ΔHo = 30kJ mol-1, ΔSo = 102 J mol-1 K-1

  42. 17.4 – Free Energy

  43. 17.5 – Entropy Changes in Chemical Reactions • Third Law of Thermodynamics: • The absolute entropy, S, of a perfect crystal, at 0K, is zero. • This is an unattainable ideal, never been observed. It is used as a reference point for the determination of entropy. • For Gibbs free energy and enthalpy, you cannot measure absolute values, only changes. For entropy, you can measure absolute values.

  44. 17.5 – Entropy Changes in Chemical Reactions • Entropy change, ΔS • Standard entropy values that are found in a table represent the entropy change when going from absolute 0 to SATP (0K  298K) • Because entropy is an extensive (depends on the amount) state function (does not depend on path taken to get there), we can calculate the entropy change in any reaction via:

  45. 17.5 – Entropy Changes in Chemical Reactions • Example, • Predict the sign of ΔS, then calculate the value of ΔS of the following reaction: H2(g) + ½ O2(g) H2O(l)

  46. 17.6 – Free Energy and Chemical Reactions • Gibbs free energy is not measured directly in a reaction • It tells us nothing about the rate of a reaction, only the eventual equilibrium position • There are three ways to calculate it, we’ve already done the first (using ΔG = ΔH – TΔS)

  47. 17.6 – Free Energy and Chemical Reactions • The second method is similar to what we just did with entropy change, but it uses ΔGf, similar to how we calculate enthalpies of formation (another state function)

  48. 17.6 – Free Energy and Chemical Reactions • Example,

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