<\/p>\n\n\n\n
Expression of Nernst Equation<\/h2>\n\n\n\n
Nernst equation is an equation relating the capacity of an atom\/ion to take up one or more electrons (reduction potential) measured at any conditions to that measured at standard conditions (standard reduction potentials) of 298K and one molar or one atmospheric pressure.<\/p>\n\n\n\n
Nernst Equation for Single Electrode Potential<\/h3>\n\n\n\n
Ecell<\/sub> = E0<\/sup> \u2013 [RT\/nF] ln Q<\/p>\n\n\n\n Where,<\/p>\n\n\n\n The calculation of single electrode reduction potential (Ered<\/sub>) from the standard single electrode reduction potential (E\u00b0red<\/sub>) for an atom\/ion is given by the Nernst equation.<\/p>\n\n\n\n <\/p>\n\n\n\n For a reduction reaction, Nernst equation for a single electrode reduction potential for a reduction reaction<\/p>\n\n\n\n Mn+ <\/sup>+ ne\u2013 <\/sup>\u2192 <\/strong>nM is;<\/p>\n\n\n\n Ered<\/sub> = EMn+<\/sup>\/M = Eo<\/sup>Mn+<\/sup>\/M \u2013 [2.303RT\/nF] log [1\/[Mn+<\/sup>]]<\/p>\n\n\n\n Where,<\/p>\n\n\n\n For measurements carried out 298K, the Nernst equation can be expressed as follows.<\/p>\n\n\n\n E = E0<\/sup> \u2013 0.0592\/n log10<\/sub> Q<\/p>\n\n\n\n Therefore, as per the Nernst equation, the overall potential of an\u00a0electrochemical cell\u00a0is dependent on the reaction quotient.<\/p>\n\n\n\n Consider a metal in contact with its own salt aqueous solution. Reactions of metal losing an electron to become an ion and the ion gaining electron to return to the atomic state are equally feasible and are in an equilibrium state.<\/p>\n\n\n\n Mn+ <\/sup>+ ne\u2013 <\/sup>\u2192 <\/strong>nM<\/p>\n\n\n\n In the reduction reaction, \u2018n\u2019 moles of an electron is taken up by the ion against a reduction potential of Ered.<\/sub><\/p>\n\n\n\n 1.<\/strong> The work done in the movement of electron<\/p>\n\n\n\n Wred<\/sub> = nFEred<\/sub><\/p>\n\n\n\n Where,<\/p>\n\n\n\n 2.<\/strong>\u00a0Change in the\u00a0Gibbs free energy\u00a0is an indication of the spontaneity and it is also equal to the maximum useful work (other than volume expansion) done in a process.<\/p>\n\n\n\n Combining work done and Gibbs free energy change:<\/p>\n\n\n\n Wred<\/sub> = nFEred<\/sub> = \u2013 \u2206G or \u2206G = \u2013 nFEred<\/sub><\/p>\n\n\n\n 3.<\/strong> Change in the free energy at standard conditions of 298K and one molar \/one atmospheric pressure conditions is \u2206G\u00b0. From the above relation, it can be written that<\/p>\n\n\n\n \u2206G\u00b0 = \u2013 nFE\u00b0red<\/sub><\/p>\n\n\n\n Where,<\/p>\n\n\n\n 4.<\/strong>\u00a0During the reaction, concentration keeps changing and the potential also will decrease with the\u00a0rate of reaction.<\/p>\n\n\n\n To get the maximum work or maximum free energy change, the concentrations have to be maintained the same. This is possible only by carrying out the reaction under a reversible equilibrium condition.<\/p>\n\n\n\n For a reversible equilibrium reaction, vant Hoff isotherm says:<\/p>\n\n\n\n \u2206G = \u2206G\u00b0 + RT ln K<\/p>\n\n\n\n Where,<\/p>\n\n\n\n 5.<\/strong>\u00a0Substituting for free energy changes in ant Hoff equation,<\/p>\n\n\n\n <\/p>\n\n\n\n When the reactants and the products of the electrochemical cell reach\u00a0equilibrium, the value of \u0394G becomes 0. At this point, the reaction quotient and the equilibrium constant (Kc<\/sub>) are the same. Since \u0394G = -nFE, the cell potential at equilibrium is also 0.<\/p>\n\n\n\n Substituting the values of Q and E into the Nernst equation, the following equation is obtained.<\/p>\n\n\n\n 0 = E0<\/sup>cell<\/sub> \u2013 (RT\/nF) ln Kc<\/sub><\/p>\n\n\n\n The relationship between the Nernst equation, the equilibrium constant, and Gibbs energy change is illustrated below.<\/p>\n\n\n\n Nernst equation vs Equilibrium constant vs Gibbs energy change<\/strong><\/p>\n\n\n\n Converting the\u00a0natural logarithm into base-10 logarithm\u00a0and substituting T=298K (standard temperature), the equation is transformed as follows.<\/p>\n\n\n\n E0<\/sup>cell<\/sub> = (0.0592V\/n) log Kc<\/sub><\/strong><\/p>\n\n\n\n Rearranging this equation, the following equation can be obtained.<\/p>\n\n\n\n log Kc<\/sub> = (nE0<\/sup>cell<\/sub>)\/0.0592V<\/strong><\/p>\n\n\n\n Thus, the relationship between the standard cell potential and the\u00a0equilibrium constant\u00a0is obtained. When Kc<\/sub>\u00a0is greater than 1, the value of E0<\/sup>cell<\/sub>\u00a0will be greater than 0, implying that the equilibrium favours the forward reaction. Similarly, when Kc<\/sub>\u00a0is less than 1, E0<\/sup>cell<\/sub>\u00a0will hold a negative value which suggests that the reverse reaction will be favoured.<\/p>\n\n\n\n The Nernst equation can be used to calculate:<\/p>\n\n\n\n The activity of an ion in a very dilute solution is close to infinity and can, therefore, be expressed in terms of the ion concentration. However, for solutions having very high concentrations, the ion concentration is not equal to the ion activity. In order to use the Nernst equation in such cases, experimental measurements must be conducted to obtain the true activity of the ion.<\/p>\n\n\n\n Another shortcoming of this equation is that it cannot be used to measure cell potential when there is a current flowing through the electrode. This is because the flow of current affects the activity of the ions on the surface of the electrode. Also, additional factors such as resistive loss and overpotential must be considered when there is a current flowing through the electrode.<\/p>\n","protected":false},"excerpt":{"rendered":" What is Nernst Equation? The Nernst equation provides a relation between the cell potential of an electrochemical cell, the standard cell potential, temperature, and the reaction quotient. Even under non-standard conditions, the cell potentials of electrochemical cells can be determined with the help of the Nernst equation. The\u00a0Nernst equation is often used to calculate the […]<\/p>\n","protected":false},"author":1,"featured_media":0,"comment_status":"open","ping_status":"open","sticky":false,"template":"","format":"standard","meta":{"footnotes":""},"categories":[636],"tags":[],"_links":{"self":[{"href":"https:\/\/mudassirbackup.infinitycodestudio.com\/index.php\/wp-json\/wp\/v2\/posts\/5602"}],"collection":[{"href":"https:\/\/mudassirbackup.infinitycodestudio.com\/index.php\/wp-json\/wp\/v2\/posts"}],"about":[{"href":"https:\/\/mudassirbackup.infinitycodestudio.com\/index.php\/wp-json\/wp\/v2\/types\/post"}],"author":[{"embeddable":true,"href":"https:\/\/mudassirbackup.infinitycodestudio.com\/index.php\/wp-json\/wp\/v2\/users\/1"}],"replies":[{"embeddable":true,"href":"https:\/\/mudassirbackup.infinitycodestudio.com\/index.php\/wp-json\/wp\/v2\/comments?post=5602"}],"version-history":[{"count":0,"href":"https:\/\/mudassirbackup.infinitycodestudio.com\/index.php\/wp-json\/wp\/v2\/posts\/5602\/revisions"}],"wp:attachment":[{"href":"https:\/\/mudassirbackup.infinitycodestudio.com\/index.php\/wp-json\/wp\/v2\/media?parent=5602"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/mudassirbackup.infinitycodestudio.com\/index.php\/wp-json\/wp\/v2\/categories?post=5602"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/mudassirbackup.infinitycodestudio.com\/index.php\/wp-json\/wp\/v2\/tags?post=5602"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}Nernst Equation at 25o<\/sup>C<\/h3>\n\n\n\n
Derivation of Nernst Equation<\/h2>\n\n\n\n
Determining Equilibrium Constant with Nernst Equation<\/h2>\n\n\n\n
![\"\"](\"https:\/\/mdr.foobrdigital.com\/wp-content\/uploads\/2022\/07\/Screenshot-2020-11-04-at-16.png\")
<\/figure>\n\n\n\nNernst Equation Applications<\/h2>\n\n\n\n
Limitations of Nernst Equation<\/h3>\n\n\n\n