Relationship between cell potential and concentration | Physics Forums
I The Nernst Equation Reported cell potentials are typically . (K) € n is the number of moles of electrons transferred between oxidising and. The Effect of Concentration on Cell Potential: The Nernst Equation to base logarithms give the relationship of the actual cell potential (E cell), . As the reaction proceeds, the difference between the concentrations of Ag+. To understand the relationship between cell potential and the equilibrium The relationship between voltage and concentration is one of the factors that must.
- Relationship between cell potential and concentration
- 24.4: The Nernst Equation
Similarly, any gases that take part in an electrode reaction are at an effective pressure known as the fugacity of 1 atm. If these concentrations or pressures have other values, the cell potential will change in a manner that can be predicted from the principles you already know.
The Nernst Equation - Chemistry LibreTexts
Are you in danger of being electrocuted? You need not worry; without any electron transfer, there is no charge to zap you with. More to the point, however, the system is so far from equilibrium for example, there are not enough ions to populate the electric double layer that the Nernst equation doesn't really give meaningful results.
Such an electrode is said to be unpoised. What ionic concentration is needed to poise an electrode? I don't really know, but I would be suspicious of anything much below 10—6 M. The Nernst Equation works only in Dilute Ionic Solutions Ions of opposite charge tend to associate into loosely-bound ion pairs in more concentrated solutions, thus reducing the number of ions that are free to donate or accept electrons at an electrode.
For this reason, the Nernst equation cannot accurately predict half-cell potentials for solutions in which the total ionic concentration exceeds about 10—3 M. How the cell potential really depends on concentration! The Nernst equation accurately predicts cell potentials only when the equilibrium quotient term Q is expressed in activities.
Ionic activities depart increasingly from concentrations when the latter exceed 10—4 to 10—5 M, depending on the sizes and charges of the ions. If the Nernst equation is applied to more concentrated solutions, the terms in the reaction quotient Q must be expressed in "effective concentrations" or activities of the electroactive ionic species.
The resulting Es can then be used to convert concentrations into activities for use in other calculations involving equilibrium constants. Cell potentials and pH: Relate cell potentials to Gibbs energy changes Use the Nernst equation to determine cell potentials at nonstandard conditions Perform calculations that involve converting between cell potentials, free energy changes, and equilibrium constants The Nernst Equation enables the determination of cell potential under non-standard conditions.
It relates the measured cell potential to the reaction quotient and allows the accurate determination of equilibrium constants including solubility constants. The Effect of Concentration on Cell Potential: The Nernst equation is arguably the most important relationship in electrochemistry. We can therefore determine the spontaneous direction of any redox reaction under any conditions, as long as we have tabulated values for the relevant standard electrode potentials.
Determine the number of electrons transferred during the redox process. Then use the Nernst equation to find the cell potential under the nonstandard conditions.
The overall reaction involves the net transfer of two electrons: Recall that the overall reaction for this cell is as follows: Suppose that the cell initially contains 1. Thus the value of Q will increase further, leading to a further decrease in Ecell. When the concentrations in the two compartments are the opposite of the initial concentrations i. So it needs to decrease the concentration of zinc two plus ions in solution.
Electrochemical Cells and Thermodynamics
It can do that if zinc two plus ions come out of solution. So if they gain electrons to form solid zinc. So that's a reduction.
So reduction occurs on the more concentrated side, so let's write that, reduction, right here, so this would be zinc two plus ions. So this would be gaining two electrons, to form solid zinc. So overall, what is happening overall here, so let's draw a line, so we have solid zinc on both sides.
We can cancel that out, we have two electrons on both sides.
11.4: Dependence of Cell Potential on Concentration
So on the left side, we would have zinc two plus, at initial concentration of 1. And this is going to zinc two plus at. How do we find the voltage of our concentration cell? Remember, from the last few videos that the Nernst equation allows us to calculate the potential of the cell.
So let's get some more room down here, and let's write down the Nernst equation. The cell potential, which is what we're trying to find, E, is equal to the standard cell potential E zero, minus. So this is one form of the Nernst equation from the last few videos. Let's think about Q, so what would Q be for our concentration cell? So Q would be equal to the concentration of zinc two plus, this would be the concentration of zinc two plus, on the less concentrated side, so this is the concentration on the less concentrated side, over the concentration of zinc two plus on the more concentrated side.
So over the concentration on the more concentrated side. So right now, that would be. So this is what Q is equal to.
Next, let's think about the standard cell potential, so the standard cell potential E zero. What's the standard cell potential here?
Well remember, the standard cell potential is the potential under standard conditions, so one molar concentration of zinc two plus. So let's write down the reduction half-reaction, so zinc two plus, this would be at one molar, so this is a reduction half-reaction, so gaining two electrons to give us solid zinc.
If you look at a table of standard reduction potentials, the standard reduction potential for this half-reaction, is negative.
For the oxidation half-reaction, we need to show solid zinc turning into zinc two plus ions, and this would need to be a one molar concentration of zinc two plus ions, because we're talking about standard cell potential, standard conditions.
This is oxidation, so losing two electrons. The standard oxidation potential would be just the negative of the standard reduction potential.
So the standard oxidation potential is positive.