## Electrochemical/galvanic cells
+Spontaneous reaction
+
1. Find two half reactions involved (between electrode and solution)
2. Higher equation will proceed left to right
3. Lower equation will proceed right to left
- **Alkaline fuel cell** - ce{KOH(aq)} electrolyte, ce{H2(g)} from tank, ce{O2(g)} from atmosphere, water out of hydrogen side
- **Acid fuel cell** - ce{H3PO4} electrolyte, water out of oxygen side
+## Electrolysis reactions
+
+- Occurs when electricity passes through ionic compound or electrolyte solution
+- Opposite of reactions in electrochemical (galvanic) cells
+- Non-spontaneous
+- Electrical energy $\rightarrow$ chemical energy
+- Anode +ve; cathode -ve
+- Cathode & anode swap relative to galvanic cell
+- Molten e.g. ce{Na+(l)} vs electrolyte e.g. ce{Na+(aq)}
+- Molten - use echem series to determine probability of electrolyssis of solution or water
+- Min voltage = e_oxidising - e_reducing
+
+### Factors affecting electrolysis
+
+- concentration of electrolyte
+- - e- values must be "close" for electrolysis rxn to prevail over \ce{H2O}
+- nature of electrodes
+
+### Electroplating
+
+- For each ion going into plated object, an ion is replaced from the anode (+ve)
+- Electrolyte balances charges
+
+### Coulomb's law
+
+$$Q=It$$
+
+### Faraday's first law
+
+$$m \propto Q$$
+
+i.e. mass produced at cathode is proportional to charge supplied
+
+*Faraday* - charge on 1 mol of electrons = 96500 C.
+
+### Faraday's second law
+
+> To produce one mole of substance by electrolysis, a whole number of mole of electrons is needed