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 58 591. Find two half reactions involved (between electrode and solution) 602. Higher equation will proceed left to right 613. Lower equation will proceed right to left 62 63emf for each cell is calculated as $E^0(\text{red}) - E^0(\text{ox})$ 64 65For a *spontaneous* (primary/fuel cell) reaction to occur, species on left must be in electrical contact with species on lower right 66 67### Primary cells 68 69Used for low-current electronic devices. Fixed quantity of reactants. 70 71- **Zinc-carbon dry cell** - carbon rod cathode and zinc anode (case) in ammonium chloride/zinc chloride electrolyte 72- **Alkaline cell** - steel cathode (case) and steel/brass rod anode in potassium hydroxide electrolyte 73- **Silver-zinc cell** - zinc anode, graphite/silver-oxide electrolyte, potassium hydroxide electrolyte 74- **Lithium cell** - magnesium oxide anode, nickel/steel cathode (case), lithium, electrolyte. Lithium is low on electrochemical series enables higher voltage 75 76### Fuel cells 77 78Used for vehicles/long-lasting applications. ce{H2} may be replenished. 79 80- **Alkaline fuel cell** - ce{KOH(aq)} electrolyte, ce{H2(g)} from tank, ce{O2(g)} from atmosphere, water out of hydrogen side 81- **Acid fuel cell** - ce{H3PO4} electrolyte, water out of oxygen side 82 83## Electrolysis reactions 84 85- Opposite of reactions in electrochemical (galvanic) cells 86- Non-spontaneous 87- Electrical energy $\rightarrow$ chemical energy 88 89### Electroplating 90 91- For each ion going into plated object, an ion is replaced from the anode (+ve) 92- Electrolyte balances charges 93 94### Coulomb's law 95 96$$Q=It$$ 97 98### Faraday's first law 99 100$$m \propto Q$$ 101 102i.e. mass produced at cathode is proportional to charge supplied 103 104*Faraday* - charge on 1 mol of electrons = 96500 C. 105 106### Faraday's second law 107 108> To produce one mole of substance by electrolysis, a whole number of mole of electrons is needed