Standard electrode potential
In electrochemistry, it is defined as the measure of the individual potential of reversible electrode at standard state with ions at an effective concentration of 1mol dm-3 at the pressure of 1 atm The basis for an electrochemical cell, such as the galvanic cell, is always a redox reaction which can be broken down into two half-reactions: oxidation at anode (loss of electron) and reduction at cathode (gain of electron). Electricity is generated due to electric potential difference between two electrodes. This potential difference is created as a result of the difference between individual potentials of the two metal electrodes with respect to the electrolyte. (Reversible electrode is an electrode that owes its potential to changes of a reversible nature, in contrast to electrodes used in electroplating which are destroyed during their use.)
Although the overall potential of a cell can be measured, there is no simple way to accurately measure the electrode/electrolyte potentials in isolation. The electric potential also varies with temperature, concentration and pressure. Since the oxidation potential of a half-reaction is the negative of the reduction potential in a redox reaction, it is sufficient to calculate either one of the potentials. Therefore, standard electrode potential is commonly written as standard reduction potential.
The electrode potential cannot be obtained empirically. The galvanic cell potential results from a pair of electrodes. Thus, only one empirical value is available in a pair of electrodes and it is not possible to determine the value for each electrode in the pair using the empirically obtained galvanic cell potential. A reference electrode, standard hydrogen electrode (SHE), for which the potential is defined or agreed upon by convention, needed to be established. In this case SHE is set to 0.00 V and any electrode, for which the electrode potential is not yet known, can be paired with SHE—to form a galvanic cell—and the galvanic cell potential gives the unknown electrode's potential. Using this process, any electrode with an unknown potential can be paired with either the SHE or another electrode for which the potential has already been derived and that unknown value can be established.
Since the electrode potentials are conventionally defined as reduction potentials, the sign of the potential for the metal electrode being oxidized must be reversed when calculating the overall cell potential. Note that the electrode potentials are independent of the number of electrons transferred —they are expressed in volts, which measure energy per electron transferred—and so the two electrode potentials can be simply combined to give the overall cell potential even if different numbers of electrons are involved in the two electrode reactions.
Standard reduction potential table
The larger the value of the standard reduction potentials, the easier it is for the element to be reduced (accept electrons); in other words, they are better oxidizing agents. For example, F2 has 2.87 V and Li+ has −3.05 V. F reduces easily and is therefore a good oxidizing agent. In contrast, Li(s) would rather undergo oxidation (hence a good reducing agent). Thus Zn2+ whose standard reduction potential is −0.76 V can be oxidized by any other electrode whose standard reduction potential is greater than −0.76 V (e.g. H+(0 V), Cu2+(0.34 V), F2(2.87 V)) and can be reduced by any electrode with standard reduction potential less than −0.76 V (e.g. H2(−2.23 V), Na+(−2.71 V), Li+(−3.05 V)).
- ΔG°cell = −nFE°cell
- If E°cell > 0, then the process is spontaneous (galvanic cell)
- If E°cell < 0, then the process is nonspontaneous (electrolytic cell)
Thus in order to have a spontaneous reaction (ΔG° < 0), E°cell must be positive, where:
- E°cell = E°cathode − E°anode
where E°anode is the standard potential at the anode and E°cathode is the standard potential at the cathode as given in the table of standard electrode potential.
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