The Nernst Equation permits the determination of cell potential under non-conventional conditions. It relates the measured cell potential to the reaction quotient and allows the specific determicountry of equilibrium constants (consisting of solubility constants).

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Introduction

The Nernst Equation is acquired from the Gibbs complimentary energy under standard conditions.

(DeltaG) is also concerned (E) under basic conditions (traditional or not) via

with

(n) is the number of electrons transferred in the reactivity (from balanced reaction), (F) is the Faraday constant (96,500 C/mol), and (E) is potential difference.

Under conventional problems, Equation ef2 is then

Hence, when (E^o) is positive, the reactivity is spontaneous and when (E^o) is negative, the reactivity is non-spontaneous. From thermodynamics, the Gibbs energy adjust under non-conventional problems have the right to be related to the Gibbs energy adjust under traditional Equations via

Substituting (DeltaG = -nFE) and also (DeltaG^o = -nFE^o) into Equation ef4, we have:

<-nFE = -nFE^o + RT ln Q label5>

Divide both sides of the Equation above by (-nF), we have

Equation ef6 can be recreated in the create of (log_10):

At traditional temperature T = 298 K, the (frac2.303 RTF) term amounts to 0.0592 V and Equation efGeneralized Nernst Equation have the right to be rewritten:

298 K>

The Equation above shows that the electrical potential of a cell relies upon the reaction quotient (Q) of the reactivity. As the redox reactivity proceeds, reactants are consumed, and therefore concentration of reactants decreases. Conversely, the commodities concentration increases because of the boosted in products formation. As this happens, cell potential slowly decreases until the reaction is at equilibrium, at which (DeltaG = 0). At equilibrium, the reactivity quotient (Q = K_eq). Also, at equilibrium, (DeltaG = 0) and also (DeltaG = -nFE), so (E = 0).

As such, substituting (Q = K_eq) and (E = 0) right into the Nernst Equation, we have:

<0 = E^o - dfracRTnF ln K_eq label7>

At room temperature, Equation ef7 simplifies right into (notice natural log was converted to log base 10):

<0 = E^o - dfrac0.0592, Vn log_10 K_eq label8>

This deserve to be rearranged into:

The Equation above suggests that the equilibrium consistent (K_eq) is proportional to the conventional potential of the reactivity. Specifically, when:

(K > 1, E^o > 0), reactivity favors commodities formation. (K

This outcome fits Le Châtlier"s Principle, which says that as soon as a device at equilibrium experiences a readjust, the mechanism will certainly minimize that adjust by moving the equilibrium in the oppowebsite direction.


Example (PageIndex1)

The (E^o_cell = +1.10 ; V) for the Zn-Cu redox reaction:

What is the equilibrium continuous for this reversible reaction?

Solution

Under standard problems, ( = = 1.0, M) and also T = 298 K. As the reaction proceeds, () decreases as () increases. Lets say after one minute, ( = 0.05, M) while ( = 1.95, M). According to the Nernst Equation, the cell potential after 1 minute is:

As you have the right to see, the initial cell potential is (E = 1.10, V), after 1 minute, the potential drops to 1.05 V. This is after 95% of the reactants have actually been consumed. As the reactivity proceeds to progress, even more (Cu^2+) will be consumed and even more (Zn^2+) will certainly be created (at a 1:1 ratio). As an outcome, the cell potential proceeds to decrease and also once the cell potential drops down to 0, the concentration of reactants and also commodities stops altering.

This is as soon as the reactivity is at equilibrium. From from Equation 9, the (K_eq) deserve to be calculated from

<eginalign log K_eq & = dfrac2 imes 1.10, V0.0592,V\ & = 37.2 endalign>

This make feeling from a Le Châtlier"s Principle, since the reactivity strongly favors the products over the reactants to cause a large (E^o_cell) of 1.103 V. Hence, the cell is greatly out of equilibrium under traditional problems. Reactions that are simply weakly out of equilibrium will certainly have actually smaller (E^o_cell) values (neglecting a readjust in (n) of course).


References

Atkins, Peter and de Paula, Julio. Physical muzic-ivan.infoistry for the Life Sciences. New York: W.H. Freemale and also Company kind of. p. 214-222. Sherwood, Lauralee. Person Physiology 6th edition. Thompson Corp. 2007. p. 77

Outside Links

Feiner, A.-S.; McEvoy, A. J. "The Nernst Equation." J. muzic-ivan.info. Educ. 1994, 71, 493. Thompkid, Martin L.; Kateley, Laura J. "The Nernst Equation: Determicountry of Equilibrium Constants for Complex Ions of Silver." J. muzic-ivan.info. Educ.

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