If the electrode potential of copper is +0.34V and that of zinc is -0.76V, calculate the e.m.f of the cell.įrom the above question, zinc will always act as an anode in the presence of copper, and the values of the standard electrode potentials provided are the SRPs, so we apply equation (i) or (ii) (iii), orĮ°cell = E°oxidation - E°reduction. However, if the standard electrode potentials are given as standard oxidation potentials (SOPs), i.e, based on the reverse reactions of the half-cell reactions in the table, then equations (iii) and (iv) below are applied.Į°cell = E°anode - E°cathode. (ii)Įquations (i) and (ii) are used if the standard electrode potentials of the electrode are provided as standard reduction potentials (as in the table above). (i), orĮ°cell = E°reduction - E°oxidation. Recall that the e.m.f of the cell is given as:Į°cell = E°cathode - E°anode. The left-hand side of the salt-bridge represents the oxidation half-reaction, while the right-hand side represents the reduction half-reaction. Where the single slash (/) represents the interface between the electrode and its electrolyte, and the double slash (//) represents the salt-bridge. The cell can be represented symbolically by:Īn electrochemical cell can be generally represented by:Īnode(s)/Anode(aq)//Cathode(aq)/Cathode(s) Recall the electrochemical cell made up of a Zn2+(aq)/Zn(s) system and a Cu2+(aq)/Cu(s) system, where the zinc electrode is the anode and the copper electrode, the cathode. It is the potential difference between the two electrodes when no current is flowing. The e.m.f of a cell, E°cell, is a measure of the driving force with which current flows out of the cell. This implies that those elements lower in the series (especially those with positive SRP values) will make good cathodes, while those above (especially the ones with negative SRP values) can act as good anodes in electrochemical calls.įor the non-metallic series, the oxidizing ability or the tendency of atoms to gain electrons increases downward, while the reducing ability or the tendency of ions to lose electrons increases upward. If we take a closer look at the metallic series, it will be observed that the oxidizing ability or tendency of ions to gain electrons increases downward and decreases upward while the reducing ability or the tendency of atoms to lose electrons decreases downward and increases upward. Similarly, the electrode potential of the reversed reaction will become the standard oxidation potential with an opposite sign. The half-cell reactions represented above are the reduction half-reactions and the reverse reactions will give us the oxidation half-reactions. The above table is the electrochemical or electromotive series of some elements showing their standard reduction potentials (SRP), E° in volts. Half-Cell Reactions Std Reduction Potential, E° (V) m.f & free energy and the relationship between e. Here, we will focus on the calculations involving electrode potentials, which include calculations of the electromotive force (e.m.f) of electrochemical cells, the relationship between e. In our last post, we looked at the overview of electrode potentials, where we discussed metal ions/metal systems or half-cells, standard electrode potential and electrochemical cells in depth.
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