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| In [[electrochemistry]], the '''Nernst equation''' is an equation that relates the [[reduction potential]] of a [[half-cell]] (or the total [[voltage]] ([[electromotive force]]) of the [[electrochemical cell|full cell]]) at any point in time to the [[standard electrode potential]], [[Thermodynamic temperature|temperature]], [[Thermodynamic activity|activity]], and [[reaction quotient]] of the underlying reactions and species used. When the reaction quotient is equal to the equilibrium constant of the reaction for a given temperature, i.e. when the concentration of species are at their equilibrium values, the Nernst equation gives the equilibrium voltage of the half-cell (or the full cell), which is zero; at equilibrium, Q=K, ΔG=0, and therefore, E=0. It is named after the German physical chemist who first formulated it, [[Walther Nernst]].<ref name="isbn0-8412-1572-3">{{cite book |author=Orna, Mary Virginia; Stock, John |title=Electrochemistry, past and present |publisher=American Chemical Society |location=Columbus, Ohio |year=1989 |pages= |isbn=0-8412-1572-3 |oclc= 19124885|doi= |accessdate=}}</ref><ref name=Wahl2005>{{Cite journal | last = Wahl | year = 2005 | title = A Short History of Electrochemistry | journal = Galvanotechtnik | volume = 96 | issue = 8 | pages = 1820–1828 }}</ref>
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| The Nernst equation gives a formula that relates the numerical values of the [[concentration gradient]] to the [[electric gradient]] that balances it. For example, if a concentration gradient is established by dissolving KCl in half of a divided vessel that is full of H<sub>2</sub>O, and then a membrane permeable to K<sup>+</sup> ions is introduced between the two halves, then after a relaxation period, an equilibrium situation arises where the chemical concentration gradient, which at first causes ions to move from the region of high concentration to the region of low concentration, is exactly balanced by an electrical gradient that opposes the movement of charge.
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| ==Expression==
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| The two (ultimately equivalent) equations for these two cases (half-cell, full cell) are as follows:
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| :<math>
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| E_\text{red} = E^{\ominus}_\text{red} + \frac{RT}{zF} \ln\frac{a_\text{Ox}}{a_\text{Red}}
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| </math> (half-cell reduction potential for <math>\text{Ox} + z e^- = \text{Red}</math>)
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| :<math>
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| E_\text{cell} = E^{\ominus}_\text{cell} - \frac{RT}{zF} \ln Q_r
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| </math> (total cell potential)
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| where
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| *{{math|''E''<sub>red</sub>}} is the half-cell [[reduction potential]] at the temperature of interest
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| *{{math|''E''<sup><s>o</s></sup><sub>red</sub>}} is the [[standard electrode potential|''standard'' half-cell reduction potential]]
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| *{{math|''E''<sub>cell</sub>}} is the cell potential ([[electromotive force]]) at the temperature of interest
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| *{{math|''E''<sup><s>o</s></sup><sub>cell</sub>}} is the ''standard'' cell potential
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| *{{mvar|R}} is the [[universal gas constant]]: {{math|''R'' {{=}} 8.314 472(15) J K<sup>−1</sup> mol<sup>−1</sup>}}
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| *{{mvar|T}} is the [[absolute temperature]]
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| *{{mvar|a}} is the chemical [[activity (chemistry)|activity]] for the relevant species, where {{math|''a''<sub>Red</sub>}} is the [[reductant]] and {{math|''a''<sub>Ox</sub>}} is the [[oxidant]]. {{math|''a''<sub>X</sub> {{=}} ''γ''<sub>X</sub>''c''<sub>X</sub>}}, where {{math|''γ''<sub>X</sub>}} is the [[activity coefficient]] of species X. (Since activity coefficients tend to unity at low concentrations, activities in the Nernst equation are frequently replaced by simple concentrations.)
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| *{{mvar|F}} is the [[Faraday constant]], the number of [[coulomb]]s per [[mole (unit)|mole]] of electrons: {{math|''F'' {{=}} 9.648 533 99(24)×10<sup>4</sup> C mol<sup>−1</sup>}}
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| *{{mvar|z}} is the number of moles of [[electron]]s transferred in the cell reaction or [[half-reaction]]
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| *{{math|''Q''<sub>r</sub>}} is the [[reaction quotient]].
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| At room temperature (25 °C), {{math|''RT''/''F''}} may be treated like a constant and replaced by 25.693 mV for cells.
| | '''MathML''' |
| | :<math forcemathmode="mathml">E=mc^2</math> |
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| The Nernst equation is frequently expressed in terms of base 10 [[logarithms]] (''i.e.'', [[common logarithm]]s) rather than [[natural logarithms]], in which case it is written, ''for a cell at 25 °C'':
| | <!--'''PNG''' (currently default in production) |
| | :<math forcemathmode="png">E=mc^2</math> |
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| :<math> | | '''source''' |
| E = E^0 - \frac{0.05916\mbox{ V}}{z} \log_{10}\frac{a_\text{Ox}}{a_\text{Red}}. | | :<math forcemathmode="source">E=mc^2</math> --> |
| </math> | |
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| The Nernst equation is used in [[physiology]] for finding the [[electric potential]] of a [[cell membrane]] with respect to one type of [[ion]].
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| ==Nernst potential== | | ==Demos== |
| {{main|Reversal potential}}
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| The Nernst equation has a physiological application when used to calculate the potential of an ion of charge ''z'' across a membrane. This potential is determined using the concentration of the ion both inside and outside the cell:
| | Here are some [https://commons.wikimedia.org/w/index.php?title=Special:ListFiles/Frederic.wang demos]: |
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| :<math>E = \frac{R T}{z F} \ln\frac{[\text{ion outside cell}]}{[\text{ion inside cell}]} = 2.3026\frac{R T}{z F} \log_{10}\frac{[\text{ion outside cell}]}{[\text{ion inside cell}]}.</math>
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| When the membrane is in [[thermodynamic equilibrium]] (i.e., no net flux of ions), the [[membrane potential]] must be equal to the Nernst potential. However, in physiology, due to active [[Na+/K+-ATPase|ion pumps]], the inside and outside of a cell are not in equilibrium. In this case, the [[resting potential]] can be determined from the [[Goldman equation]]:
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| <math>E_{m} = \frac{RT}{F} \ln{ \left( \frac{ \sum_{i}^{N} P_{M^{+}_{i}}[M^{+}_{i}]_\mathrm{out} + \sum_{j}^{M} P_{A^{-}_{j}}[A^{-}_{j}]_\mathrm{in}}{ \sum_{i}^{N} P_{M^{+}_{i}}[M^{+}_{i}]_\mathrm{in} + \sum_{j}^{M} P_{A^{-}_{j}}[A^{-}_{j}]_\mathrm{out}} \right) }</math>
| | ==Test pages == |
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| *<math>E_{m}</math> = The membrane potential (in [[volt]]s, equivalent to [[joule]]s per [[coulomb]])
| | To test the '''MathML''', '''PNG''', and '''source''' rendering modes, please go to one of the following test pages: |
| *<math>P_\mathrm{ion}</math> = the permeability for that ion (in meters per second)
| | *[[Displaystyle]] |
| *<math>[ion]_\mathrm{out}</math> = the extracellular concentration of that ion (in [[Mole (unit)|moles]] per cubic meter, to match the other [[SI]] units, though the units strictly don't matter, as the ion concentration terms become a dimensionless ratio) | | *[[MathAxisAlignment]] |
| *<math>[ion]_\mathrm{in}</math> = the intracellular concentration of that ion (in moles per cubic meter) | | *[[Styling]] |
| *<math>R</math> = The [[ideal gas constant]] (joules per [[kelvin]] per mole) | | *[[Linebreaking]] |
| *<math>T</math> = The temperature in [[kelvin]] | | *[[Unique Ids]] |
| *<math>F</math> = [[Faraday constant|Faraday's constant]] (coulombs per mole) | | *[[Help:Formula]] |
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| The potential across the cell membrane that exactly opposes net diffusion of a particular ion through the membrane is called the Nernst potential for that ion. As seen above, the magnitude of the Nernst potential is determined by the ratio of the concentrations of that specific ion on the two sides of the membrane. The greater this ratio the greater the tendency for the ion to diffuse in one direction, and therefore the greater the Nernst potential required to prevent the diffusion.
| | *[[Inputtypes|Inputtypes (private Wikis only)]] |
| | | *[[Url2Image|Url2Image (private Wikis only)]] |
| A similar expression exists that includes r (the absolute value of the transport ratio). This takes transporters with unequal exchanges into account. See: [[Sodium-Potassium Pump]] where the transport ratio would be 2/3. The other variables are the same as above. The following example includes two ions: Potassium (K<sup>+</sup>) and sodium (Na<sup>+</sup>). Chloride is assumed to be in equilibrium.
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| <math>V_{m} = \frac{RT}{F} \ln{ \left( \frac{ rP_{K}[K]_{o} + P_{Na}[Na]_{o}}{ rP_{K}[K]_{i} + P_{Na}[Na]_{i}} \right) }</math>
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| When Chloride (Cl<sup>−</sup>) is taken into account, its part is flipped to account for the negative charge.
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| <math>V_{m} = \frac{RT}{F} \ln{ \left( \frac{ P_{K}[K]_{o} + P_{Na}[Na]_{o} + P_{Cl}[Cl]_{i}}{ P_{K}[K]_{i} + P_{Na}[Na]_{i} + P_{Cl}[Cl]_{o}} \right) }</math>
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| ==Derivation==
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| ===Using Boltzmann factors===
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| For simplicity, we will consider a solution of redox-active molecules that undergo a one-electron reversible reaction
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| :<math>\text{Ox} + e^- \rightleftharpoons \text{Red}\,</math>
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| and that have a standard potential of zero. The [[chemical potential]] <math>\mu_c</math> of this solution is the difference between the energy barriers for taking electrons from and for giving electrons to the [[Cyclic voltammetry|working electrode]] that is setting the solution's [[electrochemical potential]].
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| The ratio of oxidized to reduced molecules, [Ox]/[Red], is equivalent to the probability of being oxidized (giving electrons) over the probability of being reduced (taking electrons), which we can write in terms of the [[Boltzmann factor]] for these processes:
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| :<math>
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| \frac{[\mathrm{Red}]}{[\mathrm{Ox}]}
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| = \frac{\exp \left(-[\mbox{barrier for losing an electron}]/kT\right)}
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| {\exp \left(-[\mbox{barrier for gaining an electron}]/kT\right)}
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| = \exp \left(\mu_c / kT \right).
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| </math>
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| Taking the natural logarithm of both sides gives
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| :<math>
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| \mu_c = kT \ln \frac{[\mathrm{Red}]}{[\mathrm{Ox}]}.
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| </math>
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| If <math>\mu_c \ne 0</math> at [Ox]/[Red] = 1, we need to add in this additional
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| constant:
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| :<math>
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| \mu_c = \mu_c^0 + kT \ln \frac{[\mathrm{Red}]}{[\mathrm{Ox}]}.
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| </math>
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| Dividing the equation by {{mvar|e}} to convert from chemical potentials to electrode potentials, and remembering that {{math|''k''/''e'' {{=}} ''R/F''}},<ref>{{math|''R''{{=}}''N<sub>A</sub>k''}}; see [[gas constant]]<br>{{math|''F''{{=}}''eN<sub>A</sub>''}}; see [[Faraday constant]]</ref> we obtain the Nernst equation for the one-electron process
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| <math>\mathrm{Ox} + e^- \rightarrow \mathrm{Red}</math>:
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| :<math>
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| \begin{align}
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| E &= E^0 + \frac{kT}{e} \ln \frac{[\mathrm{Red}]}{[\mathrm{Ox}]} \\
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| &= E^0 - \frac{RT}{F} \ln \frac{[\mathrm{Ox}]}{[\mathrm{Red}]}.
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| \end{align}
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| </math>
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| ===Using thermodynamics (chemical potential)===
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| Quantities here are given per molecule, not per mole,
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| and so [[Boltzmann constant]] ''k'' and the electron charge ''e'' are used
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| instead of the gas constant ''R'' and Faraday's constant ''F''. To convert
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| to the molar quantities given in most chemistry textbooks, it is simply
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| necessary to multiply by Avogadro's number: <math>R = kN_A</math> and
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| <math>F = eN_A</math>.
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| The entropy of a molecule is defined as
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| :<math>
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| S \ \stackrel{\mathrm{def}}{=}\ k \ln \Omega,
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| </math>
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| where <math>\Omega</math> is the number of states available to the molecule.
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| The number of states must vary linearly with the volume ''V'' of the
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| system, which is inversely proportional to the concentration ''c'', so
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| we can also write the entropy as
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| :<math>
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| S = k\ln \ (\mathrm{constant}\times V) = -k\ln \ (\mathrm{constant}\times c).
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| </math>
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| The change in entropy from some state 1 to another state 2 is therefore
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| :<math>
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| \Delta S = S_2 - S_1 = - k \ln \frac{c_2}{c_1},
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| </math>
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| so that the entropy of state 2 is
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| :<math>
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| S_2 = S_1 - k \ln \frac{c_2}{c_1}.
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| </math>
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| If state 1 is at standard conditions, in which <math>c_1</math> is unity (e.g., | |
| 1 atm or 1 M), it will merely cancel the units of <math>c_2</math>. We can, therefore,
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| write the entropy of an arbitrary molecule ''A'' as
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| :<math>
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| S(A) = S^0(A) - k \ln [A], \,
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| </math>
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| where <math>S^0</math> is the entropy at standard conditions and [''A''] denotes the
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| concentration of ''A''.
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| The change in entropy for a reaction
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| :<math>
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| aA + bB \rightarrow yY + zZ
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| </math>
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| is then given by
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| :<math>
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| \Delta S_\mathrm{rxn} = [yS(Y) + zS(Z)] - [aS(A) + bS(B)]
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| = \Delta S^0_\mathrm{rxn} - k \ln \frac{[Y]^y [Z]^z}{[A]^a [B]^b}.
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| </math>
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| We define the ratio in the last term as the [[reaction quotient]]:
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| :<math> | |
| Q = \frac{\prod_j a_j^{\nu_j}}{\prod_i a_i^{\nu_i}} \approx \frac{[Z]^z [Y]^y}{[A]^a [B]^b}.
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| </math>
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| where the numerator is a product of reaction product activities, ''a<sub> j</sub>'', each raised to the power of a [[stoichiometric coefficient]], ''ν<sub> j</sub>'', and the denominator is a similar product of reactant activities. All activities refer to a time ''t''. Under certain circumstances (see [[chemical equilibrium]]) each activity term such as <math>a_j^{\nu_j}</math> may be replaced by a concentration term, [''A''].
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| In an electrochemical cell, the cell potential ''E'' is the
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| chemical potential available from redox reactions (<math>E = \mu_c/e</math>).
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| ''E'' is related to
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| the [[Gibbs free energy|Gibbs energy]] change <math>\Delta G</math> only by a constant:
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| <math>\Delta G = -nFE</math>, where ''n'' is the number of electrons transferred and <math>F</math> is the Faraday constant.
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| There is a negative sign because a spontaneous reaction has a negative free energy <math>\Delta G</math> and a positive potential ''E''.
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| The Gibbs energy is related to the entropy by <math>G = H - TS</math>, where ''H'' is
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| the enthalpy and ''T'' is the temperature of the system. Using these
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| relations, we can now write the change in
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| Gibbs energy,
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| :<math>
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| \Delta G = \Delta H - T \Delta S = \Delta G^0 + kT \ln Q, \,
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| </math>
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| and the cell potential,
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| :<math>
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| E = E^0 - \frac{kT}{ne} \ln Q.
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| </math>
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| This is the more general form of the Nernst equation.
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| For the redox reaction
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| <math>\mathrm{Ox} + ne^- \rightarrow \mathrm{Red},</math>
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| <math>Q = \frac{[\mathrm{Red}]}{[\mathrm{Ox}]}</math>, and we have:
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| :<math>
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| \begin{align}
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| E &= E^0 - \frac{kT}{ne} \ln \frac{[\mathrm{Red}]}{[\mathrm{Ox}]} \\
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| &= E^0 - \frac{RT}{nF} \ln \frac{[\mathrm{Red}]}{[\mathrm{Ox}]} \\ | |
| &= E^0 - \frac{RT}{nF} \ln Q.
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| \end{align}
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| </math>
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| The cell potential at standard conditions <math>E^0</math> is often
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| replaced by the formal potential <math>E^{0'}</math>, which includes some small
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| corrections to the logarithm and is the potential that is actually measured
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| in an electrochemical cell.
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| ==Relation to equilibrium==
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| At equilibrium, ''E'' = 0 and ''Q'' = ''K''. Therefore
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| :<math>
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| \begin{align}
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| 0 &= E^o - \frac{RT}{nF} \ln K\\
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| \ln K &= \frac{nFE^o}{RT}
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| \end{align}
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| </math>
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| Or at [[Standard conditions for temperature and pressure|standard temperature]],
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| :<math>\log_{10} K = \frac{nE^o}{59.2\text{ mV}} \quad\text{at }T = 298 \text{ K}.</math>
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| We have thus related the [[standard electrode potential]] and the [[equilibrium constant]] of a redox reaction.
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| ==Limitations==
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| In dilute solutions, the Nernst equation can be expressed directly in terms of concentrations (since activity coefficients are close to unity). But at higher concentrations, the true activities of the ions must be used. This complicates the use of the Nernst equation, since estimation of non-ideal activities of ions generally requires experimental measurements.
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| The Nernst equation also only applies when there is no net current flow through the electrode. The activity of ions at the electrode surface changes [[electrochemical kinetics|when there is current flow]], and there are additional [[overpotential]] and resistive loss terms which contribute to the measured potential.
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| At very low concentrations of the potential-determining ions, the potential predicted by Nernst equation approaches toward ±∞. This is physically meaningless because, under such conditions, the [[exchange current density]] becomes very low, and there is no thermodynamic equilibrium necessary for Nernst equation to hold. The electrode is called to be unpoised in such case. Other effects tend to take control of the electrochemical behavior of the system.
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| ===Time dependence of the potential===
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| The expression of time dependence has been established by Karaoglanoff.<ref>{{citation |first=Z. |last=Karaoglanoff |authorlink=Zakhari Karaoglanoff |title=Über Oxydations- und Reduktionsvorgänge bei der Elektrolyse von Eisensaltzlösungen |language=German |trans_title=On Oxidation and Reduction Processes in the Electrolysis of Iron Salt Solutions |journal=Zeitschrift fuer Elektrochemie |volume=12 |issue=1 |pages=5–16 |date=January 1906 |doi=10.1002/bbpc.19060120105 |url=http://onlinelibrary.wiley.com/doi/10.1002/bbpc.19060120105/abstract}}</ref><ref>{{citation |title=Electrochemical Dictionary |editor1-first=Allen J. |editor1-last=Bard |editor2-first=György |editor2-last=Inzelt |editor3-first=Fritz |editor3-last=Scholz |contribution=Karaoglanoff equation |pages=527–528 |url=http://books.google.com/books?id=4TBWg3dIyKQC&pg=PA527&lpg=PA527&hl=en&f=false |publisher=Springer}}</ref><ref>{{citation |title=Introduction to Polarography and Allied Techniques |first=Kamala |last=Zutshi |page=127–128 |url=http://books.google.com/books?id=WmiaCVH-MEIC&pg=PA127&lpg=PA127&hl=en&f=false |isbn= |year=2008}}</ref><ref>The Journal of Physical Chemistry, Volume 10, p 316. http://books.google.com/books?id=zCMSAAAAIAAJ&pg=PA316&lpg=PA316&hl=en&f=false</ref>
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| ==Significance to related scientific domains==
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| The equation has been involved in the scientific controversy involving [[cold fusion]]. The discoverers of cold fusion, Fleischmann and Pons, calculated that a palladium cathode immersed in a heavy water electrolysis cell could achieve up to 10<sup>27</sup> atmospheres of pressure on the surface of the cathode, enough pressure to cause spontaneous nuclear fusion. In reality, only 10,000-20,000 atmospheres were achieved. [[John R. Huizenga]] claimed their original calculation was affected by a misinterpretation of Nernst equation.<ref>{{cite journal |separator=, | last=Huizenga | first=John R. | authorlink=John R. Huizenga | title=Cold Fusion: The Scientific Fiasco of the Century | edition=2 | location=Oxford and New York | publisher=Oxford University Press | year=1993 | pages=33, 47 | isbn=0-19-855817-1 }}</ref> He cited a paper about Pd-Zr alloys.<ref>[http://jes.ecsdl.org/content/136/3/630.abstract J. Y. Huot, et al J. Electroche. Soc 136 631 (1989)]</ref>
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| ==See also==
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| *[[Concentration cell]]
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| *[[Electrode potential]]
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| *[[Galvanic cell]]
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| *[[Goldman equation]]
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| *[[Membrane potential]]
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| *[[Nernst–Planck equation]]
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| ==References==
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| {{Reflist}}
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| ==External links==
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| * [http://www.nernstgoldman.physiology.arizona.edu/ Nernst/Goldman Equation Simulator]
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| * [http://www.physiologyweb.com/calculators/nernst_potential_calculator.html Nernst Equation Calculator]
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| * [http://thevirtualheart.org/GHKindex.html Interactive Nernst/Goldman Java Applet]
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| * [http://www.doitpoms.ac.uk/tlplib/pourbaix/index.php DoITPoMS Teaching and Learning Package- "The Nernst Equation and Pourbaix Diagrams"]
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| [[Category:Electrochemical equations]]
| |
