# Chapter 17 Lecture Outline

## 17.2 Spontaneous and Nonspontaneous Processes

1. Spontaneous reactions (go forward)

2. [delta]H is part of the picture

3. [delta]H does not explain it all (ie: evaporation of water)

## 17.3 Entropy and the Second Law of Thermodynamics

1. Introduce entropy (S)
1. Dispersion of Energy ( internet ©1997, Saunders)
2. Dispersion of Matter ( internet ©1997, Saunders)

2. Effect Entropy
1. Phase of Matter ( internet ©1997, Saunders)
2. Temperature ( internet ©1997, Saunders)
3. Complexity of Molecule ( internet ©1997, Saunders)
4. Ionic Solids ( internet ©1997, Saunders)

3. Second law of thermodynamics "entropy of the universe increases for spontaneous processes"

4. [delta]S = Sfinal - Sinitial

## 17.5 Gibbs Free Energy

1. Free Energy and Spontaneity
1. Relate to Hiking
1. climbing a mountain
1. Takes work
2. But reward?
3. Reward must be > or = work for you to climb
2. Going into a valley
1. easy
2. reward or penalty?
3. tradeoff between easy and penalty

2. free energy is [delta]G

3. [delta]G = [delta]H - T*[delta]S

4. Relate [delta]G to spontaneous reactions

## 17.6 Entropy Changes in Chemical Reactions: Calculating dSorxn

1. Entropy at phase change
1. Equlibrium between phases
2. Difference in energy ([delta]H) offset by change in entropy (T*[delta]S)
3. T*[delta]S = [delta]H (At equlibrium)

2. Example with water
1. Boils at 100 oC
2. [delta]Hvap = 40.7 kJ/mole
3. Calculate [delta]Svap
4. Units
1. [delta]H (kJ mole-1)
2. [delta]S (J K-1 mole-1)

3. Calculating [delta]Srxn from absolute S
1. [delta]S for H2O(l) -> H2O (g)
2. Compare to above
3. (NOTE: S in table is at 298 K, changes some with temp.)

## 17.7 Free Energy Changes in Chemical Reactions: Calculating dGorxn

1. Calculating [delta]G from [delta]Hrxn, [delta]Srxn, and [delta]Gf
1. Note about [delta]H and [delta]S at different temperature

2. Work an example with formation of water
1. 2 H2 + O2 -> 2 H2O
2. [delta]Hrxn = -285.9 kJ/mole
3. [delta]Srxn = -265.7 J/mole K
4. [delta]Grxn = -206.7 kJ/mole at 25°C
5. Change temp, and vary [delta]G

3. For reactions in table have class calculate [delta]H, [delta]S, and [delta]G at 25 oC

 Reaction [delta]H, kJ/rxn [delta]S, J/(K rxn) [delta]Go, kJ/rxn (25oC) [delta]Go, kJ/rxn (300oC) 2 O3 -> 3 O2 -284.5 138.1 -325 -363.6 NO2 + O2 -> NO + O3 200.5 4.3 199.2 198.0 2 SO2 + O2 -> 2 SO3 -198.3 -189.0 -142 -103.8 NH4Cl -> NH3 + HCl 176.1 284.3 91.4 13.2

1. Look at trends in [delta]S (from reaction)
2. Look at effect of T on [delta]G, Find at 300 °C

## 17.8 Free Energy Changes for Nonstandard States: The Relationship between dGorxn and dGrxn

1. [delta]G = [delta]G° + RT lnQ (effect of concentration on [delta]G°)
2.  Reaction [delta]G°(kJ/mol) [delta]G (kJ/mol) 2 O3 (.001 bar) -> 3 O2 (0.2 bar) -327 -304 NO2 (10-6 bar) + O2 (0.2 bar) -> NO (10-6 bar) + O3(10-6 bar) 198 168 2SO2 (10-7bar) + O2 (0.2 bar) --> 2 SO3(10-7bar) -142 -138 NH4Cl -> NH3(10-8 bar) + HCl (10-8 bar) 91.1 -0.2

3. Solutions

## 17.9 Free Energy and Equilibrium: Relating dGorxn and to the Equilibrium Constant (K)

1. [delta]H = T [delta]S
2. [delta]G = 0
3. Q = K
4. [delta]G° = -RT lnK
5. Look at how Q effects [delta]G°
1. For a spontaneous forward reaction [delta]G < 0
2. For a non-spontaneous forward reaction [delta]G > 0
3. Relate to figure 20.9

6. Temperature effect on [delta]G and on K
1. 2 NO2 --> N2O4 at low Temp ( Internet ©1997, Saunders)
2. 2 NO2 --> N2O4 at high Temp ( Internet ©1997, Saunders)

7. Varying Pressure for the reaction, calculate [delta]G:
8. H2 + Br2 --> 2 HBr

 H2 (bar) Br2 (bar) HBr (bar) [delta]G (kJ/mole) 1 1 1 -106.9 10-6 10-6 100 -15.6 10-8 10-8 500 15.2