1. 1 The Laws of Thermodynamics in Review 1.The internal energy* of an isolated system is constant. There are only two ways to change internal energy – heat and work. 2.The entropy of the universe increases in any spontaneous process. 3.The entropy of a perfect crystal approaches zero as the temperature approaches absolute zero. * Internal energy (U) is the energy stored in a system. It includes the energy of chemical bonds and intermolecular forces as well as kinetic energy.

2. 2 1 st Law: Heat and Work in Review Heat and work are path functions (not state functions). Heat at a constant pressure is called Enthalpy. By definition, at constant pressure: Ex) Imagine a reaction done in a balloon. The volume of the balloon may change over the course of the reaction. i) If the balloon shrinks, work has been done on the system. ii) If the balloon expands, the system has done work. The amount of work done on the system depends on the pressure and on the change in volume. At constant pressure:

3. 3 1 st Law: Heat and Work (Review) Recalling that and that at constant pressure, we can relate enthalpy and internal energy: Therefore, enthalpy under constant pressure can be defines as:

4. 4 2 nd Law: Gibbs Free Energy Recall that at constant temperature, the entropy change is: As at constant P, Hence at constant P and T: The heat absorbed by the system will be the same as the heat released by the surroundings, hence:

5. 5 2 nd Law: Gibbs Free Energy Recall that: Then: Therefore: For any spontaneous reaction, -T  S universe will be negative.

6. 6 2 nd Law: Gibbs Free Energy -T  S universe is defined in terms of the system only, and is used to define a new term, the Gibbs free energy (G), as: i) At constant temperature and pressure, since enthalpy is only meaningful at constant pressure. ii) For any spontaneous reaction,  G will be negative. iii) Gibbs free energy is a state function.

7. 7 Gibbs Free Energy and the Standard State Changes in Gibbs free energy are determined by reaction conditions. Hence free energy values are always made in reference to standard states. Standard temperature and pressure (STP) are defined as: i) T = 25 °C ii) P = 1 bar (100 kPa) Standard states are defined as: i) A solid is in the standard state if it is at STP. ii) A liquid is in the standard state if it is at STP. iii) A gas is in the standard state if it is at STP and behaves ideally. iv) A solution is in the standard state if it is has one solute with a concentration of 1 mol/L, is at STP and behaves ideally. Note that “a liquid” is not the same as “a solution”.

8. 8 Gibbs Free Energy and the Standard State Standard molar Gibbs free energy of formation (  f G° or  f G m °) is generally tabulated alongside standard molar enthalpy of formation and standard entropy. It refers to the free energy change when one mole of a compound is made from pure elements in their standard states. e.g. If a substance is a pure element in its standard state, this means that its standard molar Gibbs free energy of formation must be 0.

9. 9 Gibbs Free Energy and Predicting Spontaneity As the Gibbs free energy is a state function it is used to calculate the free energy change for any reaction (  r G°) in the standard state. e.g.

10. 10 Gibbs Free Energy and Predicting Spontaneity Note the similarity to Hess’s Law. Important notes: The state of matter is important. i) Only one allotrope of any element can have  f G° = 0. Most often the one that is most stable under STP. ex) graphite, O 2(g), etc.; white phosphorus are exceptions. ii) As ions cannot exist in isolation hence without counterions; Therefore the standard state H + (aq) (1 M at STP) is used as the reference for ions. ie.  f G°(H + (aq) ) = 0.

11. 11 Gibbs Free Energy and Predicting Spontaneity A reaction that we know is spontaneous occurs when potassium metal is added to water. What is the free energy change for this reaction under standard conditions? e.g.

12. 12 Gibbs Free Energy and Predicting Spontaneity A reaction that we know doesn’t proceed so easily is the reaction of hydrogen and nitrogen to make ammonia. What is the free energy change for this reaction under standard conditions?

13. 13 Gibbs Free Energy and Activity The Haber-Bosch process isn’t performed under standard conditions so this free energy value doesn’t apply in this case. Reaction conditions: 450 o C, 100 atm, catalyst: magnetite (Fe 3 O 4 ) In fact, many reactions are run under nonstandard conditions. To deal with reactions under nonstandard conditions, we need to consider the activity (a) of the substances involved. Activity is a measure of a substance’s deviation from the standard state: i) Any substance in the standard state will have a = 1. Recall that the standard state requires that gases and solutes behave ideally. ii) Activity is dimensionless and therefore has no units.

14. 14 Gibbs Free Energy and Activity For ideal substances, activity is relatively straightforward: i) Pure solids have a = 1 ii) Ideal gases have a = P/P° where P° = 1 bar iii) Ideal solutes have a = c/c° where c° = 1 mol/L iv) Ideal solvents have a = X where X is the mole fraction of solvent molecules. For nonideal substances, activity must be modified by an activity coefficient (  ): i) Nonideal solutes have a =  c/c° where c° = 1 mol/L ii) Activity coefficients are concentration-dependent as they account for such nonideal behaviour arising from intermolecular forces between different particles of solute. Thus, the activity coefficient will tend towards 1 with dilution.

15. 15 Gibbs Free Energy & Nonstandard Conditions To calculate the free energy of a reaction under nonstandard conditions, we use the free energy under standard conditions and the deviation from the standard state conditions: where  r G m is the molar Gibbs free energy change under the given conditions, R is the ideal gas constant, T is the given temperature (in K) and Q is the reaction quotient: There is a similarity between the formula for reaction quotient (Q) and the formula for equilibrium constant (K). The equilibrium constant only describes the system at equilibrium.

16. 16 Gibbs Free Energy & Nonstandard Conditions Calculate the molar Gibbs free energy change for the rusting of iron under atmospheric conditions. Q = 10.39  G = -736.4 kJ/mol

17. 17 Gibbs Free Energy & Nonstandard Conditions Calculate the molar Gibbs free energy change for dissolving carbon dioxide in a soft drink under atmospheric conditions. Q = 2.66  G=10.8 kJ/mol

18. 18 Gibbs Free Energy & Nonstandard Conditions Repeat the calculation for dissolving carbon dioxide in a soft drink with a higher pressure for carbon dioxide (5.0 bar). Q = 0.0202  G=-1.26 kJ/mol

19. 19 Free Energy Change and Equilibrium Recall that the free energy change for a reaction under any conditions (  r G) can be calculated from the free energy change of the same reaction under standard conditions (  r Gº): where Q is the reaction quotient: We also know that: i) The forward reaction is spontaneous if  r G < 0. ii) The reverse reaction is spontaneous if  r G > 0. What if  r G = 0?

20. 20 Free Energy Change and Equilibrium When  r G = 0, we can say that: Which rearranges to give: The forward and reverse reactions are in balance, ie at equilibrium! Therefore: Q = K Hence:

21. 21 Free Energy Change and Equilibrium I) If we know the equilibium constant for a reaction, we can calculate its standard molar free energy change: II) If we know the standard molar free energy change for a reaction, we can therefore calculate its equilibrium constant: Note that these equations refer to  r Gº (for standard conditions) NOT to  r G (which is ZERO at equilibrium!)

22. 22 Q vs. K Consider a very generic equilibrium reaction: i) When does Q = K?at equilibrium K =Q = a B /a A ii) When is Q a A (eq) or a B Q < K Reaction proceeds forward When is Q > K? When there is excess B or a depletion of A w.r.t equilibrium a A a B (eq) => Q > K Reaction proceeds in reverse A  B

23. 23 Free Energy Change and Equilibrium Consider the dimerization of NO 2 : Calculate the equilibrium constant for this reaction at 25 ºC. K = 6.74

24. 24 Free Energy Change and Equilibrium Is the dimerization of NO 2 spontaneous at 25 ºC if the partial pressure of NO 2 is 0.4 bar and the partial pressure of N 2 O 4 is 1.8 bar?

25. 25 Free Energy Change and Equilibrium: Acids The acid ionization constant, K a, is the equilibrium constant for the reaction in which an acid ionizes. e.g. pK a is a measure of strength of an acid. An acid with a low pK a is stronger than an acid with a high pK a. The pK a value for an acid comes from its acid ionization constant: Also an acid will be weaker than another acid whose K a value is smaller. (Even a “big” K a value is pretty small. K a values range from about 10 -50 to 10 10.)

26. 26 Free Energy Change and Equilibrium: Acids Calculate the standard molar free energy of formation for HF (aq).

27. 27 Free Energy Change and Equilibrium:Solubility The solubility product, K sp, is the equilibrium constant for the reaction in which an ionic compound dissolves in water. e.g. Ex) Calculate the solubility of lead(II) iodide in water at 25 ºC. Solubility is reported in #moles of solid soluble in 1L water.

28. 28 Free Energy Change and Equilibrium:Solubility Simple K sp calculations is that they almost never give good results. Multiple simultaneous equilibria that must be considered, not just the one for K sp. When PbI 2 is dissolved in water, two complex ions also form: In this case the K sp calculation does gives almost the same answer as a complete calculation factoring in both complex ions; because, the K f values for both complex ions are relatively small (compared to some K f values in the 10 30 range). Usually the complex ion formation is significant enough that it must be considered. Otherwise in the solubility calculation will give very poor values.

29. 29 Free Energy Change and Equilibrium:Solubility Complex ions are not the only factor complicating solubility calculations. Ex) What equilibria would you have to consider if you wanted to calculate the solubility of carbon dioxide in water? How would you increase the solubility?