Nernst equilibrium

Corrosion — Corrosion current density — Figure. Polarization curves of a metal/metal ion electrode and the H2/H+ electrode including the anodic and cathodic partial current curves, the Nernst equilibrium electrode potentials E(Me/Mez+) and (H2/H+), their exchange current densities / o,M> o,redox and related overpotentials Me) and 77(H), the rest potential r, the polarization n and the corrosion current density ic at open circuit conditions (E = Er) [i]... 

Nernst equation — A fundamental equation in -> electrochemistry derived by - Nernst at the end of the nineteenth century assuming an osmotic equilibrium between the metal and solution phases (- Nernst equilibrium). This equation describes the dependence of the equilibrium electrode - potential on the composition of the contacting phases. The Nernst equation can be derived from the - potential of the cell reaction (Ecen = AG/nF) where AG is the - Gibbs energy change of the - cell reaction, n is the charge number of the electrochemical cell reaction, and F is the - Faraday constant. 

Nernst equilibrium — It was - Nernst who first treated the thermodynamical - equilibrium for an -> electrode [i], and derived the - Nernst equation. Although the model used by Nernst was not appropriate (see below) the Nernst equation - albeit in a modified form and with a different interpretation - is still one of the fundamental equations of electrochemistry. In honor of Nernst when equilibrium is established at an electrode, i.e., between the two contacting phases of the electrode or at least at the interface (interfacial region), it is called Nernst equilibrium. In certain cases (see - reversibility) the Nernst equation can be applied also when current flows. If this situation prevails we speak of reversible or... 

Passivation — Metals usually dissolve in acidic electrolytes when their electrode potential becomes more positive than the value of the related -> Nernst equilibrium potential of the metal/metal ion electrode. The dissolution current density increases exponentially with... 

Express the equilibrium potential in terms of ion conductances for a membrane permeable to both Na+ and K+ ions with relative conductivities gNa and gK- [Hint assume that the flux of each ion is proportional to conductivity multiplied by the driving force, which can be expressed as the difference between the membrane potential and the Nernst equilibrium potential for a given ion.]... 

Synthetic OHB19/23/polyP complexes demonstrated weak but explicit selectivity for divalent over monovalent ions79 when incorporated in planar bilayers composed of synthetic di22 l PC, cholesterol (5 1 w/w) between unequal solutions of Ca2+ and Na+ at pH 7.4 (Figure 15C). The reversal potential was -20 mV the Nernst equilibrium potentials were the same as for the synthetic complexes above. The data indicate selectivity for Ca2+ over Na+ of about 4 1, a significant improvement over channels formed by the oligomers alone (see Figure 8C), but still much poorer discrimination than the >90 1 selectivity demonstrated by the natural and synthetic... 

An increase in the permeability of the PM to sodium ions (Na+) permits Na+ to enter the cell down a concentration gradient with a consequent increase in the positive charge within the cell that opposes Na+ entry. At equilibrium there is no further net entry and v m approximates to the Nernst equilibrium potential (t]txj for Na+ given by the following equation (noting that. c =tlie charge on the ion (+1)) ... 

Similarly, increasing the permeability of the membrane to K+ (PK+) will permit K+ to flow out of the cell down a concentration gradient, this efflux of positively charged K+ causing the inside of the cell to be more negative with respect to the outside and hence increasingly opposing further efflux. At equilibrium, when there is no further net efflux of K+, the v]tm approximates to the Nernst equilibrium potential (vjt -) for K+ ... 

In this case, the Nernst equilibrium potential, jvie/Me + represents the limit of the stability ranges of both 2D and 3D Me phases. At " =, 2D and 3D Me... 

The corrosion behavior of metals cannot be predicted from the position of their standard potentials in the electrochemical series because the potential of an electrode changes with the current density. If an electrode in which only one electrode process takes place is termed a working electrode and the resultant potential, a working potential, then the differences between working potential and the Nernst equilibrium potential is called an overpotential, that is caused by reaction restraints. In general, polarization is defined as the shift in potential of working electrodes within a corrosion element. In such an element, at least two electrode reactions occur whose overpotentials are superimposed, resulting in the polarization effect. 

Sodium currents were evoked by voltage pulses from a holding potential of —90 mV to a membrane potential, of —60 to +80 mV for 3 ms (Fig. lOA). All stimulation for sodium currents was followed by a similar pulse protocol in which the pulse amplitude was reduced to 1/4 and the holding potential was brought to —120 mV (P/4 protocol [42]). A plot of the peak current as a function of the membrane potential is shown in Fig. lOB. The peak current was estimated by a cubic fit of each record in a short interval around the peak with a third-order polynomial. The reversal potential, V, i.e., the applied voltage for which the peak current changes sign is estimated around 53 mV, which is very close to the Nernst equilibrium potential for sodium ions in these experimental conditions. 

Following the seminal paper of Butler (37) in 1924 on the kinetic basis of Nernst equilibrium potentials, an electrochemical rate equation was written by Erdey-Gruz and Volmer (14), for a net current-density i, in terms of components of i for the forward and backward directions of the process. They recognized that only some fraction (denoted by a or 3) of the electrical energy change riF associated with change of electrode potential, would exponentially modify the current, giving a potential-dependent rate-equation of the form ... 

The first condition for T > 0 is the Nernst equilibrium at potential 0 the second expresses that the sum of fluxes of R and P is zero the third that there is no flux in C into or out of the electrode and we have the ENC also holding at the electrode (or very close to it in solution). Note that it was necessary here to set the potential if to zero in the bulk region X = L), and to use the potential rather than the potential field E. In fact, one finds from the results that E is level near the outer limit X = L, so that E could have been used in the equation system, setting a zero dE/dX at X = L as a boundary condition. But this was not certain beforehand and could not safely be assumed. 

Here, represents the crossover effect of fiiel/oxidizer through the electrolyte to the opposite electrode or internal short circuits in the cell that is responsible for more departure of theoretical equilibrium open-circuit potential from Nernst equilibrium voltage. 

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