Nernst equation


In electrochemistry, the Nernst equation is an equation that relates the reduction potential of an electrochemical reaction to the standard electrode potential, temperature, and activities 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
gibs 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. The 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, the equation can be written as
where
Similarly to equilibrium constants, activities are always measured with respect to the standard state. 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, the thermal voltage is approximately 25.693 mV. The Nernst equation is frequently expressed in terms of base-10 logarithms rather than natural logarithms, in which case it is written:
where λ=ln 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.

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, the membrane potential must be equal to the Nernst potential. However, in physiology, due to active 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, which is a solution of G-H-K influx equation under the constraints that total current density driven by electrochemical force is zero:
where
When chloride 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. 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 over the probability of being reduced, 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, 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, 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 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 each activity term such as may be replaced by a concentration term, .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. 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 is 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.

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.