Cells Electrochemical cells can be used as a commercial source of electrical energy Cells can be non-rechargeable (irreversible), rechargeable and fuel cells. You should be able to work out Ecell for given half reactions. Example primary non rechargeable cells Dry Cell Zn2+(aq) + 2e- Zn(s) E = – 0.76 V 2MnO2 (s) + 2NH4+ (aq) + 2 e- → Mn2O3 (s) + 2NH3 (aq) + H2O(l) E = 0.75 V More negative half equation will oxidise Overall reaction 2MnO2 + 2NH4++ Zn → Mn2O3 + 2NH3 + H2O + Zn2+ Ecell =+1.51V Example secondary rechargeable cells PbSO4 + 2e- Pb + SO4 2- PbO2 + SO4 2- + 4H+ + 2e- PbSO4 + 2H2O E= -0.356V E= +1.685 V Lead acid Cell Overall reaction PbO2 +Pb + 2SO4 2- + 4H+ 2 PbSO4 + 2H2O Ecell= +2.04V The forward reaction occurs on discharge giving out charge. Charging causes the reaction to reverse Reversible cells only work if the product stays attached to the electrode and does not disperse 6 Example primary Lithium –manganese dioxide cell- non rechargeable Li (s)| Li+ aq | | Li+ aq | MnO2 (s) , LiMnO2(s) | Pt (s) Li+ aq +e- Li (s) Li+ aq + MnO2 (s) +e- LiMnO2(s) Ecell =+2.91V E = – 3.04 V E = – 0.13 V (Mn will reduce changing oxidation state from +4 to +3) Overall reaction Li + MnO2 → LiMnO2 More negative half equation will oxidise You do not need to learn the details of most of these cells. Relevant cell information will be given. You should be able to convert between standard electrode potential half cells, full cell reactions and cell diagrams and be able to calculate potentials from given data. Conventional cell diagram Ecell = E red- Eox = -0.13 – – 3.04 = 2.91 V. Example secondary Nickel–cadmium cells are used to power electrical equipment such as drills and shavers. They are rechargeable cells. The electrode reactions are shown below. NiO(OH) + H2O + e- Ni(OH)2 + OH– E = +0.52 V (Ni will reduce changing oxidation state from 3 to 2) Cd(OH)2 + 2e- Cd + 2OH– E = –0.88 V (Cd will oxidise changing oxidation state from 0 to 2) Overall reaction discharge 2NiO(OH) + Cd + 2H2O 2Ni(OH)2 + Cd(OH)2 E= +1.40V Ecell = E red- Eox = +0.52 – – 0.88 = + 1.40 V The overall reaction would be reversed in the recharging state 2Ni(OH)2 + Cd(OH)2 2NiO(OH) + Cd + 2H2O Li | Li+ || Li+ , CoO2 | LiCoO2 | Pt Example secondary Lithium ion cells are used to power cameras and mobile phones. Li+ + CoO2 + e- Li+ [CoO2 ] – E=+0.6V Li+ + e- Li E=-3.0V Overall discharge Li + CoO2 LiCoO2 E=3.6V reaction (Co will reduce changing oxidation state from 4 to 3 Conventional cell diagram The reagents in the cell are absorbed onto powdered graphite that acts as a support medium. The support medium allows the ions to react in the absence of a solvent such as water. Water would not be good as a solvent as it would react with the lithium metal. The overall reaction would be reversed in the recharging state
126.96.36.199 Commercial applications of electrochemical cells (A-level only)
Electrochemical cells can be used as a commercial source of electrical energy.
The simplified electrode reactions in a lithium cell:
Positive electrode: Li+ + CoO2 + e– → Li+[CoO2 ] – Negative electrode: Li → Li+ + e–
Cells can be non-rechargeable (irreversible), rechargeable or fuel cells.
Students should be able to:
• use given electrode data to deduce the reactions occurring in non-rechargeable and rechargeable cells
• deduce the EMF of a cell
• explain how the electrode reactions can be used to generate an electric current.