And it is a valuable supplementarytext for undergraduate- and graduate-level chemistry students. Milazzo and S. Hempel Furthermore, when we consulted published EMF data, we found that the compilations often paid no attention to disparate conditions and varying pressures and electrolyte concentrations. To achieve a useful compilation, we resorted to Nernst equation thermodynamic calculations to reconcile disparate data. While it is impossible to determine the electrical potential of a single electrode, we can assign an electrode the value of zero and then use it as a reference.
Use this information to calculate the missing electrode potential for the half-equation shown in the table. Each potential value in this table may be interpreted as the interfacial potential difference of the test electrode relative to the standard hydrogen electrode. Therefore, standard electrode potential is commonly written as standard reduction potential. The electrode potential cannot be obtained empirically.
The galvanic cell potential results from a pair of electrodes. Thus, only one empirical value is available in a pair of electrodes and it is not possible to determine the value for each electrode in the pair using the empirically Abstract: Calculations of the standard electrode potential of the silver-silver bromide electrode indicate that many values listed in textbooks are incorrect.
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These data allow us to compare the oxidative and reductive strengths of a variety of substances. The strongest oxidant in the table is F 2 , with a standard electrode potential of 2. This high value is consistent with the high electronegativity of fluorine and tells us that fluorine has a stronger tendency to accept electrons it is a stronger oxidant than any other element.
This fact might be surprising because cesium, not lithium, is the least electronegative element. The apparent anomaly can be explained by the fact that electrode potentials are measured in aqueous solution, where intermolecular interactions are important, whereas ionization potentials and electron affinities are measured in the gas phase.
Lithium metal is therefore the strongest reductant most easily oxidized of the alkali metals in aqueous solution. Species that lie below H 2 are stronger oxidizing agents.
Any species on the left side of a half-reaction will spontaneously oxidize any species on the right side of another half-reaction that lies below it in the table. Conversely, any species on the right side of a half-reaction will spontaneously reduce any species on the left side of another half-reaction that lies above it in the table. The black tarnish that forms on silver objects is primarily Ag 2 S. The half-reaction for reversing the tarnishing process is as follows:.
Given: reduction half-reaction, standard electrode potential, and list of possible reductants. Asked for: reductants for Ag 2 S, strongest reductant, and potential reducing agent for removing tarnish. Determine which species is the strongest reductant. Given: redox reaction and list of standard electrode potentials Table P2. B Adding the two half-reactions gives the overall reaction:.
The standard cell potential is quite negative, so the reaction will not occur spontaneously as written. B The two half-reactions and their corresponding potentials are as follows. The standard potential for the reaction is positive, indicating that under standard conditions, it will occur spontaneously as written.
Hydrogen peroxide will reduce MnO 2 , and oxygen gas will evolve from the solution. To answer these questions requires a more quantitative understanding of the relationship between electrochemical cell potential and chemical thermodynamics. When using a galvanic cell to measure the concentration of a substance, we are generally interested in the potential of only one of the electrodes of the cell, the so-called indicator electrode , whose potential is related to the concentration of the substance being measured.
To ensure that any change in the measured potential of the cell is due to only the substance being analyzed, the potential of the other electrode, the reference electrode , must be constant. You are already familiar with one example of a reference electrode: the SHE. The potential of a reference electrode must be unaffected by the properties of the solution, and if possible, it should be physically isolated from the solution of interest. To measure the potential of a solution, we select a reference electrode and an appropriate indicator electrode.
Whether reduction or oxidation of the substance being analyzed occurs depends on the potential of the half-reaction for the substance of interest the sample and the potential of the reference electrode.
The potential of any reference electrode should not be affected by the properties of the solution to be analyzed, and it should also be physically isolated. There are many possible choices of reference electrode other than the SHE. The SHE requires a constant flow of highly flammable hydrogen gas, which makes it inconvenient to use. Consequently, two other electrodes are commonly chosen as reference electrodes.
One is the silver—silver chloride electrode , which consists of a silver wire coated with a very thin layer of AgCl that is dipped into a chloride ion solution with a fixed concentration. The cell diagram and reduction half-reaction are as follows:.
If a saturated solution of KCl is used as the chloride solution, the potential of the silver—silver chloride electrode is 0. That is, 0. A second common reference electrode is the saturated calomel electrode SCE , which has the same general form as the silver—silver chloride electrode.
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