Electrochemistry

1. Electrochemistry Redox reactions The half-reactions for Sn2+(aq) + 2Fe3+(aq)  Sn4+(aq) + 2Fe2+(aq) are Oxidation: Sn2+(aq)  Sn4+(aq) +2e- Reduction: 2Fe3+(aq) + 2e-  2Fe2+(aq) In oxidation: electrons are products. In reduction: electrons are reactants.

2. Voltaic or Galvanic Cells • The energy released in a spontaneous redox reaction can be used to perform electrical work. • Voltaic or galvanic cells are devices in which electron transfer occurs via an external circuit. • Voltaic cells use spontaneous reactions. • If a strip of Zn is placed in a solution of CuSO4, Cu is deposited on the Zn and the Zn dissolves by forming Zn2+. • Zn is spontaneously oxidized to Zn2+ by Cu2+. • The Cu2+ is spontaneously reduced to Cu0 by Zn.

3. Galvanic cells consist of • Anode: Zn(s)  Zn2+(aq) + 2e2 • Cathode: Cu2+(aq) + 2e- Cu(s) • Salt bridge (used to complete the electrical circuit): cations move from anode to cathode, anions move from cathode to anode. • The two solid metals are the electrodes (cathode and anode). • As oxidation occurs, Zn is converted to Zn2+ and 2e-. The electrons flow towards the anode where they are used in the reduction reaction.

4. A shorthand convention exists for describing batteries. For the Cu/Zn battery, it would be described as follows: Zn(s)|Zn2+(aq)|| Cu2+(aq)|Cu(s) The ANODE... The CATHODE... -supplies electrons to the external circuit (wire) -accepts electrons from the external circuit (wire) -is the negative pole of the battery -is the positive pole of the battery -is the site of OXIDATION -is the site of REDUCTION -is written on the left hand side if the convention -is written on the right hand side if the is followed convention is followed -is the half-cell with the lowest electrode potential -is the half-cell with the highest electrode potential

5. Cell EMF • Electromotive force (emf) is the force required to push electrons through the external circuit. • Cell potential: Ecell is the emf of a cell (in Volts). • For 1M solutions at 25 C (standard conditions), the standard emf (standard cell potential) is called Ecell. Standard Reduction Potentials • Convenient tabulation of electrochemical data. • Standard reduction potentials, Ered are measured relative to the standard hydrogen electrode (SHE).

6. Standard reduction potentials To predict the reactivity of oxidants or reductants we need to measure the potential of each half-reaction.  impossible!!....for every oxidation we have a reduction Define a standard half-cell of potential = 0V against which all other half-cell reduction potentials are measured. Each component in these standard cells having unit concentration. By convention: Standard (or Normal) Hydrogen Electrode is used Pt(s) | H2(g) | H+(a) || Ag+(a) | Ag(s) |_______________| NHE H+(aq) + e- H2(g)E0=0V

7. The cell potential E = E+ - E- (written as reductions) Standard Reduction Potentials • Consider Zn(s)  Zn2+(aq) + 2e-. We measure Ecell relative to the SHE (cathode): Ecell = E+(cathode) - E-(anode) 0.76 V = 0 V - Ered(anode). • Therefore, Ered(anode) = -0.76 V. • Standard reduction potentials must be written as reduction reactions: Zn2+(aq) + 2e- Zn(s), Ered = -0.76 V. • Since Ered = -0.76 V we conclude that the reduction of Zn2+ in the presence of the SHE is not spontaneous.

8. half-rxns oxidant reductant E0 (V) stronger oxidant F2(g) + 2e- <=> 2F- 2.890 Ce4+ + e- <=> Ce3+ 1.720 Ag+ + e- <=> Ag(s) 0.799 Fe3+ + e- <=> Fe2+ 0.771 O2 + 2H+ + 2e- <=> H2O2 0.695 Cu2+ + 2e- <=> Cu(s) 0.339 2H+ + 2e- <=> H2(g) 0.000 Cd2+ + 2e- <=> Cd(s) -0.402 Zn2+ + 2e- <=> Zn(s) -0.762 K+ + e- <=> K(s) -2.936 Li+ +e- <=> Li(s) -3.040 stronger reducer Electrochemical series DG = -nFE Therefore if Ecell is positive, the reaction is spontaneous

9. Applications of Galvanic Cells Potentiometry and Ion Selective Electrodes the measure of the cell potential to yield chemical information (conc., activity, charge ....) A difference in the activity of an ion on either side of a selective membrane results in a thermodynamic potential difference being created across that membrane

10. The glass pH electrode

11. Corrosion Fe2+ +2e Fe E0=-0.44V 2H+ + 2e H2 E0=0V 2H2O + O2 =4e4OH- E0=1.23V Iron is oxidized in water or humid conditions to give rust. Inhibit this by coating with another material (Zn for example that forms a protective oxide on the iron), or by providing a sacrificial anode (b).

12. Batteries-providing electricity from chemistry The Lead Acid Storage Battery was developed in the late 1800's and has remained the most common and durable of the batterytechnologies (in vehicles).

13. Lead-acid batteries When the battery is used as a voltage supply, electrons flow from the Pb metal to the Pb(IV)oxide. The reactions aren't quite thereverse of the formation reactions, because now the sulfate ions in the solution begin to play a role. The two reactions are: PbO2 + 4H+ + 2e + SO4-2 PbSO4 + 2H2O Pb + SO4-2 PbSO4 + 2e The overall reaction if we combine the hydrogen ions and the sulfate: PbO2 + Pb + 2H2SO4  2 PbSO4 + 2 H2O Lead sulfate is fairly insoluble so that as soon as Pb(II) ions are formed by either reaction, the ions immediately precipitate as leadsulfate. The beauty is that this lead sulfate stays attached to the grids so that it is there for recharging of the battery.

14. Other batteries Primary battery-non rechargeable Longer shelf-life Rechargeable Offer higher efficiencies compared to burning fuels

15. Images of batteries Leclanche Alkaline Fuel Cell

16. Electrolysis Use Faraday’s Laws to evaluate the number of moles of a substance oxidised or reduced by passage of charge (current over a given period of time = I.t) through an electrode Faraday: Q (charge) = nF N=number of moles of electrons F=constant of 96500 Coulomb/mole Example (try it): What current is needed to deposit 0.500g of chromium from a solution containing Cr3+ over a one hour period (MW for Cr=52)? (Ans=0.77A)

17. Applications of Electrolytic Cells Aluminium refining:The major ore ofaluminium is bauxite,Al2O3. Anhydrous Al2O3 melts at over 2000°C. This is too high to permit its use as a molten medium forelectrolytic formation of free aluminium. The electrolytic process commercially used to produce aluminiumis known as the Hall process, named after its inventor, Charles M. Hall. Al2O3 is dissolved inmolten cryolite, Na3AlF6, which has a melting point of 1012oC and is an effective conductor of electriccurrent. Graphite rods are employed as anodesand are consumed in the electrolysis process. The cell electrolytic reaction is: 2Al2O3 + 3C  4Al(l) + 3CO2(g)

18. Electrolysis of brine Chlorine and sodium hydroxide are both manufactured by electrolysis of brine (aqueous sodium chloride) using inert electrodes. Chlorine isevolved at the anode, Cl-1/2Cl2 + e Hydrogen is evolved at the cathode: H+ + e 1/2H2 The removal of chloride ions and hydrogen ions leaves sodium ions and hydroxide ions in solution. Chlorine is used to disinfect municipal water supplies and water in swimming pools. It is used to manufacture household bleaches anddisinfectants. It is used to manufacture plastics (e.g. PVC), pesticides,anaesthetics, CFCs etc. Sodium hydroxide is used in the manufacture of synthetic fibres, soapsand detergents.

19. Electroplating In all aspects of our lives we are surrounded byproducts with electroplated surfaces. Whether weare looking atasilver-plated watch through gold-plated glasses,watching television, using the washing machine,getting into a car or boarding a plane:electroplating plays animportant part in all of thesesituations. The objective is to prevent corrosion and wear,produce hardness andconductivity, and give products an attractiveappearance. The principle: thin metallic layers with specific properties are deposited on base materialsincluding steel, brass, aluminium, plastic and die-cast parts. Silver electroplating was the first large scaleuse of electrolysis for coating base metal objects with a higher value decorative finish.