Half-cells potentials

A problem that has fascinated surface chemists is whether, through suitable measurements, one can determine absolute half-cell potentials. If some one standard half-cell potential can be determined on an absolute basis, then all others are known through the table of standard potentials. Thus, if we know E for... 

Test method for porosity in gold platings on metal substrates by gas exposures Test method for half-cell potentials of uncoated reinforcing steel in concrete Method for detection of copper corrosion from petroleum products by the copper strip tarnish test... 

This shows that the voltage of a given cell may be thought of as being made up of two parts, one part characteristic of one of the half-reactions and one part characteristic of the other halfreaction. Chemists call these two parts half-cell potentials, a term that emphasizes the relation between voltage and potential energy. The halfcell potentials are symbolized °. 

We would like to measure the contribution each half-reaction makes to the voltage of a cell. Yet every cell involves two half-reactions and every cell voltage measures a difference between their half-cell potentials. We can never isolate one half-reaction to measure its E°. An easy escape is to assign an arbitrary value to the potential of some selected half-reaction. Then we can combine all other half-reactions in turn with this reference half-reaction and find values for them relative to our reference. The handiest arbitrary value to assign is zero and chemists have decided to give it to the half-reaction... 

Since concentration variations have measurable effects on the cell voltage, a measured voltage cannot be interpreted unless the cell concentrations are specified. Because of this, chemists introduce the idea of standard-state. The standard state for gases is taken as a pressure of one atmosphere at 25°C the standard state for ions is taken as a concentration of 1 M and the standard state of pure substances is taken as the pure substances themselves as they exist at 25°C. The half-cell potential associated with a halfreaction taking place between substances in their standard states is called ° (the superscript zero means standard state). We can rewrite equation (37) to include the specifications of the standard states ... 

Chemists have determined a large number of these half-cell potentials. The magnitude of the voltage is a quantitative measure of the tendency of that half-reaction to release electrons in comparison to the H2-2H+ half-reaction. If the sign is positive, the half-reaction has greater tendency to release electrons than does the H2-2H+ half-... 

The half-reactions, listed in order of decreasing half-cell potentials, are in the same order as in Table 12-1, which was dictated by laboratory experience. 

In Exercise 12-2 you placed the Ni-Ni+2 half-reaction into Table 12-1. Check your placement by examining the half-cell potential of this half-reaction in Table 12-11. 

By the half-cell potentials, we conclude the Zn-Zn+2 half-reaction has the greater tendency to release electrons. It will tend to transfer an electron to silver ion, forcing (54) in the reverse direction. Hence we obtain the net reaction by subtracting (54) from (52). But remember that this subtraction must be in the proportion that causes no net gain or loss of electrons. If two electrons are lost per atom of zinc oxidized in (52), then we must double half-reaction (54) so that two electrons will be consumed. 

If a chemist wishes to know whether zinc can be oxidized if it is placed in contact with a solution of nickel sulfate, the values of E° help him decide. The half-cell potential for Zn-Zn+2 is +0.76 volt, which is greater than that for Ni-Ni+2 (which is +0.25 volt). The difference, +0.51 volt, is positive, indicating that zinc has a greater tendency to lose electrons than does nickel. Therefore, zinc can transfer electrons to Ni+2. The chemist predicts zinc will react with Ni+2, zinc being oxidized and nickel being reduced. 

Suppose the question is whether silver will be oxidized if it is immersed in copper sulfate. The half-cell potential for Ag-Ag+ is —0.80 volt and that for Cu-Cu+2 is —0.34 volt. The first value, —0.80 volt, is more negative than the second, —0.34 volt. The difference, then, is still negative —0.80 — (—0.34) = —0.46 volt. The negative answer shows that Ag-Ag+ has less tendency to lose electrons than does Cu-Cu+2. The reaction will not tend to proceed spontaneously. Silver will not be oxidized to an appreciable extent in copper sulfate. 

The three values of ° are easily calculated from half-cell potentials. Then, we can predict with confidence that reaction (65) will not occur to an appreciable extent if solid copper is immersed in dilute acid. The negative value of ° (—0.34 volt) indicates that equilibrium in (65) strongly favors the reactants, not the products. 

How can this trend in half-cell potentials and oxidizing abilities be explained Let us imagine that reaction (7), as an example, is carried out in a hypothetical series of... 

H+], calculation of, 192, see also Hydrogen ion Haber, Fritz, 151 Haber process, 140, 150 Hafnium, oxidation number, 414 Haldane, J. B. S., 436 Half-cell potentials effect of concentration, 213 measuring, 210 standard, 210 table of, 211, 452 Half-cell reactions, 201 Half-life, 416 Half-reaction, 201 balancing, 218 potentials, 452 Halides... 

E° values have been measured for many reactions and tabulated as standard half-cell potentials. Table 9.3 summarizes half-cell potentials as standard reduction potentials for a select set of reactions.aa In the tabulations, E° for... 

W.N. Hansen, and G.J. Hansen, Reference States for Absolute Half-Cell-Potentials, Physical Review Letters 59(9), 1049-1052 (1987). 

The elemental reaction used to describe a redox reaction is the half reaction, usually written as a reduction, as in the following case for the reduction of oxygen atoms in O2 (oxidation state 0) to H2O (oxidation state —2). The half-cell potential, E°, is given in volts after the reaction ... 

The half-cell potential and the half-cell free energy change are related by the following relationship for reversible conditions ... 

The Nemst equation describes the dependence of the half-cell potential on concentration ... 

If the equilibrium half-cell potentials for two redox reactions are different, electrons will be transferred from the reduced species in the... 

The Zn /Zn reduction potential is more negative than the H3 O /H2 reduction potential (-0.76 V vs. 0 V), so zinc is the anode in this cell. Zinc is oxidized and hydronium ions are reduced, causing electrons to flow from the more negative zinc electrode to the less negative SHE. Again, we reverse the direction of the half-reaction with the more negative potential and find E by subtracting the half-cell potentials ... 

Atmospheric O2 has a partial pressure of 0.20 bar, and atmospheric water vapor is saturated with carbon dioxide. This dissolved CO2 forms carbonic acid, which generates a hydronium ion concentration of about 2.0 X 10 M. The Nemst equation allows calculation of the half-cell potential for the reduction of 02(g) under these... 

Figure 5. Typical half-cell potentials vs. current density curves tor an Flj-Oj fuel cell and aDMFC.

Potentiometric methods are based on the measurement of the potential of an electrochemical cell consisting of two electrodes immersed in a solution. Since the cell potential is measured under the condition of zero cmrent, usually with a pH/mV meter, potentiometry is an equilibrium method. One electrode, the indicator electrode, is chosen to respond to a particular species in solution whose activity or concentration is to be measured. The other electrode is a reference electrode whose half-cell potential is invariant. 

The half-cell potentials of the two reference electrodes are constant sample solution conditions can often be controlled so that E,j is effectively constant and the composition of the internal solution can be maintained so that (ai)i , ai is fixed. Consequently Eq. (3) can be simplified to give... 

In Table 7-1 the relative tendencies of certain elements to react were listed qualitatively. We can give a quantitative measure of relative tendency to react, called standard reduction potential, as shown in Table 14-2. In this table, the standard half-cell potential for each half-reaction, as a reduction, is tabulated in order with the highest potential first. If we turn these half-reactions around, we change the signs of the potentials and we get oxidation potentials. We thus have half-reactions including both elementary metals and elementary nonmetals in the same table, as well as many half-reactions that do... 

We can combine the half-cell potentials for any two half-reactions in the table to get a complete cell potential. The chemical reaction may proceed spontaneously if the complete cell potential is positive. Otherwise, the opposite reaction may proceed spontaneously. We combine half-cells by adding the chemical reactions and by adding the corresponding half-cell potentials. We must first get the correct chemical reactions and corresponding half-cell potentials for the half-reactions, as follows ... 

If you multiply the coefficients of the chemical equation by any number, leave the half-cell potential unchanged e is intensive it does not depend on the number of moles of chemical involved. 

Then add the copper half-cell reduction to the zinc half-cell oxidation and add the half-cell potentials ... 

What is the meaning of a positive sign for (a) a cell potential (b) a half-cell potential ... 

Table 8.1 Standard half-cell potentials of some oxidant species in aqueous solution at 25 °C.

A schematic diagram of a typical pH electrode system is shown in Fig. 10.1. The cell potential, i.e. the electromotive force, is measured between a pH electrode and a reference electrode in a test solution. The pH electrode responds to the activity or concentration of hydrogen ions in the solution. The reference electrode has a very stable half-cell potential. The most commonly used reference electrodes for potentiometry are the silver/silver chloride electrodes (Ag/AgCl) and the saturated calomel electrodes (SCE). 

A We obtain the two balanced half-equations and the half-cell potentials from Table 21-1. Oxidation Fe2+ (aq)-+Fe3+ (aq) + e x2 -E° = -0.771V... 

A We write down the oxidation half-equation with the method of Chapter 5, and obtain the reduction half-equation from Table 21-1, along with the reduction half-cell potential. 

A First we write down the two half-equations, obtain the half-cell potential for each, and then calculate c0ell. From that value, we determine AG°... 

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