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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. Although the overall potential of a cell can be measured, there is no simple way to accurately measure the individual electrode/electrolyte potentials in isolation. The electric potential also varies with temperature, concentration and pressure. The standard electrode potential (abbreviated E^o) is the measure of individual potential of any electrode at standard ambient conditions (temperature 298K, solutes at 1 M and gases at 1 bar). 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.

Contents 1 Calculation of Standard Electrode Potential 2 The Standard Reduction Potential table 3 Non-standard condition 4 See also 5 External links 6 References

[edit] Calculation of Standard Electrode Potential

To overcome the difficulty of measuring individual potential of an electrode, an electrode of unknown reduction potential at anode can be paired with a standard cathode whose unknown potential has been defined to be zero. Such a zero potential standard, called the standard hydrogen electrode, consists of a platinized platinum electrode in contact with 1 M H⁺ ions and bathed by hydrogen gas at 1 atm. Given that the potential of cathode (reduction half-reaction) is standardized at zero, the potential at the anode (oxidation half-reaction at the unknown electrode) can be measured.

For example, to measure the standard reduction potential of zinc metal electrode, an electrochemical cell can be built with the anode compartment containing zinc metal electrode in 1 M ZnSO₄ solution. The anode oxidation half-reaction is then:

Zn_(s) → Zn²⁺_(aq) + 2e^-

The electrons produced at the anode flow through the wire and at the cathode, the reduction half-reaction occurs as follows:

2H⁺_(aq) + 2e^- → H_2(g)

Where concentration of H⁺ is 1 M and pressure of H₂ gas is at 1 atm. The line notation for this cell is:

Zn_(s) | Zn²⁺_(aq),(1M) || 2H⁺_(aq) | H_2(g),(1atm)

Since the reduction half-reaction has a potential of zero, the reading of the voltmeter E^o_cell (or EMF) corresponds to the potential of the zinc metal because:

E^o</sup>_cell(0.76V)= E^o_{2H⁺(aq) → H2(g)}(0V) + E^o_{Zn(s) → Zn²⁺(aq)}(0.76V)

where the superscript ^o denotes the standard states that are employed. Through this way, the standard reduction potentials for common electrodes can be measured. (See Table of standard electrode potentials)

Since the conventional values are given as reduction potentials, in combining two half-reactions in a redox reaction, the sign of the known reduction potential for the metal electrode being oxidized must be reversed. Also, since the number of electrons lost must equal the number gained, the half-reactions must be balanced by multiplying with an integer. However, the value of standard reduction potential is independent of the number of times the reaction occurs. Thus the value of E^o_cell is not changed when the half-reactions are multiplied by an integer.

[edit] The Standard Reduction Potential table

Main article: Reduction potential

Since the values are given in their ability to be reduced, the bigger the standard reduction potentials, the easier they are to be reduced, in other words, they are simply better oxidizing agents. For example, F₂ has 2.87V and Li⁺ has -3.05V. F₂ reduces easily and is therefore a good oxidizing agent. In contrast, Li⁺ would rather undergo oxidation (hence a good reducing agent). Thus Zn²⁺ whose standard reduction potential is -0.76V can be oxidized by any other electrode whose standard reduction potential is greater than -0.76V (eg. H⁺(0V), Cu²⁺(0.16V), F₂(2.87V)) and can be reduced by any electrode with standard reduction potential less than -0.76V (eg. H₂(-2.23V), Na⁺(-2.71V), Li⁺(-3.05V)).

In a galvanic cell, where a spontaneous redox reaction drives the cell to produce an electric potential, Gibbs free energy ΔG^o must be negative.

(If E^o_cell> 0, we have a spontaneous process (galvanic cell))

(If E^o_cell< 0, we have a nonspontaneous process (electrolytic cell))

ΔG^o_cell is related to E^o_cell by the equation:

ΔG^o_cell = -nFE^o_cell

where n is number of moles of electrons per mole of products

F is the Faraday constant, ~96,485 C/mol

Thus in order to have a spontaneous reaction (-ΔG^o), E^o_cell must be positive, where

E^o</sup>_cell= E^o_anode + E^o_cathode

where E^o_anode is the standard potential at the anode (reverse the sign of the standard reduction potential value for the electrode) and E^o_cathode is the standard potential at the cathode as given in the table of standard electrode potential.

[edit] Non-standard condition

The standard electrode potentials are measured at standard conditions. Yet in most cases, real cells may operate under non-standard conditions. Given the standard potential of the cell, the non-standard cell potential can be calculated using the Nernst equation:

E_cell= E^{o</sub> - <math>\frac{RT}{nF}ln\frac{[red]}{[oxd]}</math>}

[edit] See also

• Nernst equation
• Electrochemical cell
• Electrochemical potential
• Concentration cell
• Table of standard electrode potentials
• Galvanic series

[edit] External links

• Standard Hydrogen Potential
• Redox Equilibria
• Chemistry of Batteries
• Electrochemical Cells

[edit] References

Zumdahl, Steven S.; Zumdahl, Susan A. (2000). Chemistry (5th ed.), Houghton Mifflin Company. ISBN 0-395-98583-8

Atkins, Peter; Jones, Loretta. (2005). Chemical Principles(3rd ed.), W.H. Freeman and Company. ISBN 0-7167-5701-X

Yanbing Z, Manon M, Couture, Derrick R, Kolling, Antony R, Lindsay D. Eltis, James A and Judy Hirst (2003). Biochemistry, 42, 12400-12408de:Standardpotenzial it:Potenziale standard di riduzione ru:Стандартный электродный потенциал