Redox Reactions OIL RIG
by Cindy Powell
1. Voltaic cells
1.1. Spontaneous: Uses chemical reactions
1.2. Anode is negative; oxidation occurs
1.2.1. The more reactive metal is located here
1.2.2. Mg>Al>Zn>Fe>Pb>Cu>Ag
1.2.3. Ex: Zn(s) --> Zn+2 (aq) + e-
1.2.4. The E cell is less positive
1.3. Cathode is positive; reduction occurs
1.3.1. The less reactive metal is located here
1.3.2. Ex: Cu+2(aq) +2e- --> Cu(s)
1.3.3. The E cell is more positive
1.3.3.1. The more Positive, the more readily is reduced (able to accept electrons)
1.4. The electrons flow from the anode (-) to the cathode (+) which creates a current and can be read with a voltmeter.
1.5. A salt bridge is required to complete the circuit. The build-up of anions that occur at the cathode can move to the anode. The cations at the anode can move to the anode. This helps maintain the potential difference.
2. Electrolytic cells
2.1. Non-spontaneous so electricity is added to make the redox reaction occur
2.2. Anode is positive and oxidation occurs here
2.3. Cathode is negative an reduction occurs here
3. Spontaneity
3.1. Voltaic cell will always run in a direction that gives a + E cell; if -E cell, then reverse reaction is spontaneous
3.2. delta G= -nFE
3.2.1. If +E cell, -G= spontaneous
3.2.2. If -E cell, +G= non-spontaneous
3.2.3. E cell & G=0 = equilibrium
3.2.4. F= Faradays' constant which is the charge carried by 1 mol of e- = 96,500 C mol-1
3.2.4.1. E is J=C x V
3.2.4.1.1. G= Gibb's free energy in J