Nernst equation

In electrochemistry, the Nernst equation is an equation that relates the reduction potential of an electrochemical reaction (half-cell or full cell reaction) to the standard electrode potential, temperature, and activities (often approximated by concentrations) of the chemical species undergoing reduction and oxidation. It was named after Walther Nernst, a German physical chemist who formulated the equation.
Expression.
A quantitative relationship between cell potential and concentration of the ions
standard thermodynamics says that the actual 
Gibbs free energy is related to the free energy change under standard state by the relationship:
where is the reaction quotient.
The cell potential associated with the electrochemical reaction is defined as the decrease in Gibbs free energy per coulomb of charge transferred, which leads to the relationship formula_2. The constant (the Faraday constant) is a unit conversion factor , where is the Avogadro constant and is the fundamental electron charge. This immediately leads to the Nernst equation, which for an electrochemical half-cell is
For a complete electrochemical reaction (full cell), the equation can be written as
where
Similarly to equilibrium constants, activities are always measured with respect to the standard state (1 mol/L for solutes, 1 atm for gases). The activity of species X, , can be related to the physical concentrations via , where is the activity coefficient of species X. Because activity coefficients tend to unity at low concentrations, activities in the Nernst equation are frequently replaced by simple concentrations. Alternatively, defining the formal potential as: 
the half-cell Nernst equation may be written in terms of concentrations as:
and likewise for the full cell expression.
At room temperature (25 °C), the thermal voltage formula_7 is approximately 25.693 mV. The Nernst equation is frequently expressed in terms of base-10 logarithms ("i.e.", common logarithms) rather than natural logarithms, in which case it is written:
where "λ"=ln(10) and "λVT" =0.05916...V. The Nernst equation is used in physiology for finding the electric potential of a cell membrane with respect to one type of ion. It can be linked to the acid dissociation constant.
Applications in biology.
Nernst potential.
The Nernst equation has a physiological application when used to calculate the potential of an ion of charge across a membrane. This potential is determined using the concentration of the ion both inside and outside the cell:
When the membrane is in thermodynamic equilibrium (i.e., no net flux of ions), and if the cell is permeable to only one ion, then the membrane potential must be equal to the Nernst potential for that ion.
Goldman equation.
When the membrane is permeable to more than one ion, as is inevitably the case, the resting potential can be determined from the Goldman equation, which is a solution of G-H-K influx equation under the constraints that total current density driven by electrochemical force is zero:
where
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.A similar expression exists that includes (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, so r equals 1.5 in the formula below. The reason why we insert a factor r = 1.5 here is that current density "by electrochemical force" Je.c.(Na+)+Je.c.(K+) is no longer zero, but rather Je.c.(Na+)+1.5Je.c.(K+)=0 (as for both ions flux by electrochemical force is compensated by that by the pump, i.e. Je.c.=-Jpump), altering the constraints for applying GHK equation. The other variables are the same as above. The following example includes two ions: potassium (K+) and sodium (Na+). Chloride is assumed to be in equilibrium.
When chloride (Cl−) is taken into account, 
Derivation.
Using Boltzmann factor.
For simplicity, we will consider a solution of redox-active molecules that undergo a one-electron reversible reaction
and that have a standard potential of zero, and in which the activities are well represented by the concentrations (i.e. unit activity coefficient). The chemical potential of this solution is the difference between the energy barriers for taking electrons from and for giving electrons to the working electrode that is setting the solution's electrochemical potential.The ratio of oxidized to reduced molecules, , 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:
Taking the natural logarithm of both sides gives
If at  = 1, we need to add in this additional constant:
Dividing the equation by to convert from chemical potentials to electrode potentials, and remembering that , we obtain the Nernst equation for the one-electron process :
Using thermodynamics (chemical potential).
Quantities here are given per molecule, not per mole, and so Boltzmann constant and the electron charge are used instead of the gas constant and Faraday's constant . To convert to the molar quantities given in most chemistry textbooks, it is simply necessary to multiply by the Avogadro constant: and . The entropy of a molecule is defined as
where is the number of states available to the molecule. The number of states must vary linearly with the volume of the system (here an idealized system is considered for better understanding, so that activities are posited very close to the true concentrations. Fundamental statistical proof of the mentioned linearity goes beyond the scope of this section, but to see this is true it is simpler to consider usual isothermal process for an ideal gas where the change of entropy takes place. It follows from the definition of entropy and from the condition of constant temperature and quantity of gas that the change in the number of states must be proportional to the relative change in volume . In this sense there is no difference in statistical properties of ideal gas atoms compared with the dissolved species of a solution with activity coefficients equaling one: particles freely "hang around" filling the provided volume), which is inversely proportional to the concentration , so
we can also write the entropy as
The change in entropy from some state 1 to another state 2 is therefore
so that the entropy of state 2 is
If state 1 is at standard conditions, in which is unity (e.g., 1 atm or 1 M), it will merely cancel the units of . We can, therefore, write the entropy of an arbitrary molecule A as
where is the entropy at standard conditions and [A] denotes the concentration of A.The change in entropy for a reaction
We define the ratio in the last term as the reaction quotient:
where the numerator is a product of reaction product activities, , each raised to the power of a stoichiometric coefficient, , and the denominator is a similar product of reactant activities. All activities refer to a time . Under certain circumstances (see chemical equilibrium) each activity term such as may be replaced by a concentration term, [A].In an electrochemical cell, the cell potential is the chemical potential available from redox reactions (). is related to the Gibbs energy change only by a constant: 
, where is the number of electrons transferred and is the Faraday constant. There is a negative sign because a spontaneous reaction has a negative free energy and a positive potential . The Gibbs energy is related to the entropy by , where is the enthalpy and is the temperature of the system. Using these relations, we can now write the change in Gibbs energy,
and the cell potential,
This is the more general form of the Nernst equation. For the redox reaction , 
and we have:
The cell potential at standard conditions is often replaced by the formal potential , which includes some small corrections to the logarithm and is the potential that is actually measured in an electrochemical cell.
Relation to equilibrium.
At equilibrium, the electrochemical potential and therefore the reaction quotient attains the special value known as the equilibrium constant: . Therefore,
Or at standard temperature,
We have thus related the standard electrode potential and the equilibrium constant of a redox reaction.
Limitations.
In dilute solutions, the Nernst equation can be expressed directly in the 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.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 when there is current flow, and there are additional overpotential and resistive loss terms which contribute to the measured potential. 
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 may be no thermodynamic equilibrium necessary for Nernst equation to hold. The electrode is called unpoised in such case. Other effects tend to take control of the electrochemical behavior of the system, like the involvement of the solvated electron in electricity transfer and electrode equilibria, as analyzed by Alexander Frumkin and B. Damaskin, Sergio Trasatti, etc.
Time dependence of the potential.
The expression of time dependence has been established by Karaoglanoff.
Significance to related scientific domains.
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 1027 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. He cited a paper about Pd–Zr alloys.
The equation permits the extent of reaction between two redox systems to be calculated and can be used, for example, to decide whether a particular reaction will go to completion or not. At equilibrium the emfs of the two half cells are equal. This enables to be calculated hence the extent of the reaction.