1. Chapter 17-18 Thermodynamics and Electrochemistry
Review 4
Chapter 17-18
Thermodynamics and Electrochemistry

2. First Law of thermodynamics – Law of conservation of energy (of universe)
Enthalpy change: H (KJ/mol) is a state function, heat exchange at constant pressure.
Internal energy = total energy of the system
E = q + w q is the heat and w is work
SIGNS: + if being absorbed; + if work is done on the system)
Second Law of thermodynamics
Entropy (S, J/K): randomness, disorder, is a state function
Any spontaneous reaction is associated with an increased entropy of the universe (Suniverse 0)
For equilibrium process, Suniverse = 0
Suniverse = Ssystem + Ssurrounding
P   S, T   S, Dilution, dissolution   S,
More atoms   S (greater capacity of molecule takes up energy, eg: C2H2 vs. C2H4)
Standard state: gas at 1 atm, pure solid or pure liquid, aqueous solution with concentration of 1 M at 25oC
Third Law of Thermodynamics: At absolute zero, a substance has zero entropy and is a perfect crystalline.

3. Gibb’s Free energy: maximum energy available to do work, helps to determine whether a reaction is spontaneous or not.
Grxno =  m Gproductso -  n Greactantso
G = H - TS G  0 if reaction is spontaneous
G  0 if reaction is nonspontaneous
G = 0 if reaction is in equilibrium
Go = Ho - TSo For pure elements Gof = 0
At equilibrium: G = 0 so H = TS (phase transition)
G = Go at Standard state
Free energy and Equilibrium:
G = Go + RT ln Q
at equilibrium G = 0, Q = K
 Go = -RT ln K  Go < 0, K > 1
 Go > 0, K < 1
Under rare condition, Go = 0 when K=1
Go = Ho - TSo = -RT ln K
(Ho / R){(T2 – T1)/T1T2} = ln (K2/K1)

4. Oxidation states: page 175-181, section 4.9-4.10. See Chapter 4
Redox reactions:
Oxidation states: page , section See Chapter 4
Oxidation:  ON, lose electrons, reducing agent
Reduction:  ON, gain electrons, oxidizing agent
Balance redox reaction:
Write down half-reactions (oxidation & reduction)
Balance all elements except Oxygen and Hydrogen
Use H2O balance oxygen
Use H+ balance hydrogen
Balance charges with electrons
Make sure both reduction half and oxidation half have the same number of electrons
Combine two half reactions
Use OH- to balance out the H+ if the solution is basic
Chapter 18: Electrochemistry (Chang – Chap 19)

5. Galvanic Cells (Voltaic Cells): Produce electrical energy
Electrode: solid phase on which redox reaction occurs
An example: Zn in Cu2+ ( oC)
Zn (s)  Zn2+ (aq) + 2e- Oxidation (anode)
Cu2+ (aq) + 2e-  Cu (s) Reduction (cathode)
(Daniell Cell)
Cell diagram:
Anode // cathode
Zn (s) / Zn2+ (aq) // Cu2+ (aq) / Cu (s)
pt / H2(g) / H+ (aq) / Cl- (aq) / Cl2 (g) / pt ---inert electrode
SHE as standard reduction potential
Eocell = Eocathode - Eoanode
Eocell = EoOxidation + EoReduction
Eocell > 0 spontaneous
Depends on temperature and pressure
Do not multiply coefficient

6. Nernst Equation: E = Eo – (2.303 RT/nF) log Q
F = Faraday = one mole of charges of electron = C/mol electron
J = C V
1 Amp = 1 C/sec
E = Eo – ( V/n) log Q @ 25oC
Eo = ( V/n) log K @ 25oC
Go = -nFEo
Electrolytic Cells – Use electricity to make a non spontaneous reaction work
Used of electroplating, and recovery of metals from molten solutions
1 Watt = 1 J/sec
CaCl2  Ca Cl- Ca e Ca(s) Cl-  Cl2 (g) + 2e
Grams of Ca = Amp * time in seconds * (1 mol e /96500 C) * (1 mol Ca/2 mol e) * (molar mass Ca/1 mol Ca)
Volume of Cl2 at STP, using above find moles of Cl2
Amp * time in seconds * (1 mol e /96500 C) * (1 mol Cl2/2 mol e) = Moles of Cl2 (g)
Use PV =nRT to solve for V at STP