How can we apply standard conditions in electrochemical cells?
Standard conditions are defined for convenience of the mathematical form of the reaction quotient, but not necessarily in order to make experiments easy. The standard hydrogen electrode is a case in point. The half reaction that involves the reaction of H+ at 1 M concnetration. That high concnetration corresponds to a super acid. Then the reaction
The standard hydrogen electrode (SHE) is an essential feature of redox chemistry since it defines the zero of the reduction potential scale. As was the case for enthalpy and for free energy, the reduction potential is not an absolute quantity. It is a relative quantity and therefore we must define the zero of potential. The convention is to use the hydrogen electrode under standard conditions of 1 M of concentration and 1 bar of pressure. Those numbers might give you pause since this means that we would run the reaction under conditions where the pH = 0 and there is a high pressure of H2 gas, which is explosive if the least bit of O2 were to leak into the reaction vessel. THis is, of course, less than desirable from an experimental perspective since we normally do not enjoy working around potentially explosive vessels with strong acide in them. However, the Nernst equation comes to the resucue since we can work under conditions that are not as extreme as the standard conditions and then correct the value using the concentratoin and pressure dependence of the Nernst equation. We will consider this procedure in the next segment. The details of the SHE itself as discussed in this video.
The wide use of solvents with varying pH values requires us to carefully consider the effect of non-standard [H+] concnetration on the electrochemical potnetial. SInce the [HM+} concnetration appears in the reaction quotient in the Nernst equation (i.e. in a natural logatirhm) we can use the following correction to write a compact form of the pH dependence of the electrode potnetial.
2.303 log10 [H+} = ln [H+}
This expression permits us to write the correction to cell potential as -0.0591 pH
There are two main applications of electrochemical cells. If the cell has a positive cell potnetial then the electrode reactions are spontanceous and the cell will gneerate a voltage. As current flows from one electrode to the other work can be done. This configuration is known as a Galvanic cell or a Voltaic cell. If the electrode reactions are not spontaneous then the cell is used to drive chemical processes using an input of electride energy from some other source. This type of cell is called an electrolytic cell. It can be used to synthesize compounds that would other otherwise be very difficult to obtain using standard reaction chemistry. An example is the electrolysis of NaCl to generate Na metal and Cl2 gas. THe products are use industrial materials that would be very difficult to obtain in a chemical reaction, but can easily be generated at an electrode provided there is an external voltage that is sufficiently large to drive the process. In this video we condier these definitions as well as the significance of the anode and cathode in each type of cell.
