1# Electrochemistry 2 3## Oxidation states 4 5Indicates charge (ionisation) of an element 6 7**Oxidation** - loss of e- (at anode) 8**Reduction** - gain of e- (at cathode) 9 10Main group elements (i.e. group 2) - generally one oxidation state: 11 12| elements | valence config | oxidation state | 13| ------------- | -------------- | --------------- | 14| alkali metals | $s^1$ | +1 | 15| alkali earths | $s^2$ | +2 | 16| aluminium | $s^2 p^1$ | +2 | 17| nitrogen | $s^2 p^3$ | -3 | 18| oxygen | $s^2 p^4$ | -2 | 19| halogens | $s^2 p^5$ | -1 | 20| noble gases | $s^2,\>s^2p^5$ | n/a | 21 22Transition metals (d shell) may have several oxidation states. 23 24Common oxidation numbers: 25 26| elements | common ox. state | exceptions | 27| ----------------- | ---------------- | ------------------- | 28| main group metals | valency | | 29| hydrogen | +1 | metal hydrides (-1) | 30| oxygen | -2 | ce{H2O2} (-1) | 31| halogens | -1 | | 32 33 34### Rules for oxidation states 35 36- oxidation states >3 may only exist in compounds 37- oxidation number of free element is 0 38- oxidation number of simple ion is the charge of the ion 39- sum of oxidation numbers in polyatomic ion is the charge of the ion 40- sum of oxidation numbers of a neutral compound is zero 41 42## Electrochemical series 43 44- Top is most likely to be reduced 45- Strongest reductants are bottom right 46- Strongest oxidants are top left 47- Strong oxidants have weak conjugate reductants 48- $E^0$ values are measured relative to ce{H2} / ce{H^+} = 0V 49 50## Conjugate redox pairs 51 52Oxidant and conjugate *reduced form* 53e.g. ce{Cu^2+} / ce{Cu}, $\quad$ ce{Zn^2+} / ce{Zn} 54 55Usually one member of pair is used as electrode (except for *inert electrodes*, e.g. platinum) 56 57## Electrochemical/galvanic cells 58 59Spontaneous reaction 60 611. Find two half reactions involved (between electrode and solution) 622. Higher equation will proceed left to right 633. Lower equation will proceed right to left 64 65emf for each cell is calculated as $E^0(\text{red}) - E^0(\text{ox})$ 66 67For a *spontaneous* (primary/fuel cell) reaction to occur, species on left must be in electrical contact with species on lower right 68 69### Primary cells 70 71Used for low-current electronic devices. Fixed quantity of reactants. 72 73- **Zinc-carbon dry cell** - carbon rod cathode and zinc anode (case) in ammonium chloride/zinc chloride electrolyte 74- **Alkaline cell** - steel cathode (case) and steel/brass rod anode in potassium hydroxide electrolyte 75- **Silver-zinc cell** - zinc anode, graphite/silver-oxide electrolyte, potassium hydroxide electrolyte 76- **Lithium cell** - magnesium oxide anode, nickel/steel cathode (case), lithium, electrolyte. Lithium is low on electrochemical series enables higher voltage 77 78### Fuel cells 79 80Used for vehicles/long-lasting applications. ce{H2} may be replenished. 81 82- **Alkaline fuel cell** - ce{KOH(aq)} electrolyte, ce{H2(g)} from tank, ce{O2(g)} from atmosphere, water out of hydrogen side 83- **Acid fuel cell** - ce{H3PO4} electrolyte, water out of oxygen side 84 85## Electrolysis reactions 86 87- Occurs when electricity passes through ionic compound or electrolyte solution 88- Opposite of reactions in electrochemical (galvanic) cells 89- Non-spontaneous 90- Electrical energy $\rightarrow$ chemical energy 91- Anode +ve; cathode -ve 92- Cathode & anode swap relative to galvanic cell 93- Molten e.g. ce{Na+(l)} vs electrolyte e.g. ce{Na+(aq)} 94- Molten - use echem series to determine probability of electrolyssis of solution or water 95- Min voltage = e_oxidising - e_reducing 96 97### Factors affecting electrolysis 98 99- concentration of electrolyte 100- - e- values must be "close" for electrolysis rxn to prevail over \ce{H2O} 101- nature of electrodes 102 103### Electroplating 104 105- For each ion going into plated object, an ion is replaced from the anode (+ve) 106- Electrolyte balances charges 107 108### Coulomb's law 109 110$$Q=It$$ 111 112### Faraday's first law 113 114$$m \propto Q$$ 115 116i.e. mass produced at cathode is proportional to charge supplied 117 118*Faraday* - charge on 1 mol of electrons = 96500 C. 119 120### Faraday's second law 121 122> To produce one mole of substance by electrolysis, a whole number of mole of electrons is needed