The Standard Hydrogen Electrode

If the wire between both beakers is connected a reaction begins, just as if all these constituents were mixed together in a single beaker. But as they are separated. 

These are the half-reactions or half-cell reactions. The total or whole-cell reaction is the sum of these, or 

As these reactions proceed, electrons flow through the wire from the Zn to the Pt electrode. This would produce a charge imbalance in both beakers if it weren t for the salt bridge, which diffuses Cl ions into the right hand beaker and into the left, and completes the electrical circuit. 

Let s say that we set up this cell, balance the potential with our potentiometer, and read a cell voltage of volts. We could next suppose that this voltage is made up of the sum of the two half-cell voltages, i.e., that 

Because there is no way to measure the properties of individual half-cells, we are in the same position as we were in Chapter 17 with respect to single ion activities.. We must choose a convention in this case we must choose one half-cell to reference all the others to. For example if we chose the hydrogen half-cell, we could say then that the voltage of the zinc half cell was 

The standard hydrogen electrode is the ultimate reference for Eh (and pH) measurements. This electrode is formed by bubbling hydrogen gas at 1 bar pressure over a platinum electrode in a 1 N HCl solution (cf. Bates 1964). The electrode reaction is that of the H2 gas-H ion couple 

Shown in Fig. 11.1 is a hydrogen electrode used in an apparatus to measure the standard potential of the redox couple 

4e = 2H2O is at thermodynamic equilibrium and dis.solved oxygen is at or above a detection limit of 5 fig/L. 

The measured Eh, which depends on the concentration of in the right-hand cell, is described by 

In operation this electrode dips into a solution of hydrogen ions of constant activity while hydrogen gas passes over its surface. In Fig. 5.6 are shown two types of hydrogen electrode the one has the electrolyte solution enclosed and protected from possible air contamination, the other may be dipped into a solution whose hydrogen ion concentration is to be determined. The former is more desirable for the determination of electrode potentials, the latter is suited for following changes in hydrogen ion concentrations as in titrations. 

By applying the Nemst equation to the hydrogen electrode equilibrium it is seen that the potential of a hydrogen electrode is given by 

When both the activities of hydrogen and hydrogen ion are unity E = which is arbitrarily given the value zero at all temperatures. Potentials of all other electrodes may then be given values relative to this standard. 

---

Figure 3-1 Voltage Measurements on a Silver-Silver Chloride, Hydrogen Cell at 298.15 K. The contribution of the Standard Hydrogen Electrode is taken as zero by convention.

Standard Hydrogen Electrode The standard hydrogen electrode (SHE) is rarely used for routine analytical work, but is important because it is the reference electrode used to establish standard-state potentials for other half-reactions. The SHE consists of a Pt electrode immersed in a solution in which the hydrogen ion activity is 1.00 and in which H2 gas is bubbled at a pressure of 1 atm (Figure 11.7). A conventional salt bridge connects the SHE to the indicator half-cell. The shorthand notation for the standard hydrogen electrode is... 

Corrective action should be initiated when value is > — 0.23 V against the standard hydrogen electrode (SHE). Plant-specific values should be estabUshed for protection of stainless steels and nickel-based critical components. 

Laboratory experiments have shown that IGSCC can be mitigated if the electrochemical potential (ECP) could be decreased to —0.230 V on the standard hydrogen electrode (SHE) scale in water with a conductivity of 0.3 ]lS/cm (22). This has also been demonstrated in operating plants. Equipment has been developed to monitor ECP in the recirculation line and in strategic places such as the core top and core bottom, in the reactor vessel during power operation. 

Fig. 6. Band edge positions of several semiconductors ia contact with an aqueous electrolyte at pH 1 ia relation to the redox (electrode) potential regions (vs the standard hydrogen electrode) for the oxidation of organic functional groups (26,27).

Other Coordination Complexes. Because carbonate and bicarbonate are commonly found under environmental conditions in water, and because carbonate complexes Pu readily in most oxidation states, Pu carbonato complexes have been studied extensively. The reduction potentials vs the standard hydrogen electrode of Pu(VI)/(V) shifts from 0.916 to 0.33 V and the Pu(IV)/(III) potential shifts from 1.48 to -0.50 V in 1 Tf carbonate. These shifts indicate strong carbonate complexation. Electrochemistry, reaction kinetics, and spectroscopy of plutonium carbonates in solution have been reviewed (113). The solubiUty of Pu(IV) in aqueous carbonate solutions has been measured, and the stabiUty constants of hydroxycarbonato complexes have been calculated (Fig. 6b) (90). 

It must not be assumed that the protection potential is numerically equal to the equilibrium potential for the iron/ferrous-ion electrode (E ). The standard equilibrium potential (E ) for iron/ferrous-ion is -0-440V (vs. the standard hydrogen electrode). If the interfacial ferrous ion concentration when corrosion ceases is approximately 10 g ions/1 then, according to the Nernst equation, the equilibrium potential (E ) is given by ... 

Since the single potential of a metal cannot be measured it is necessary to use a suitable reference elecrode such as the Hg/Hg2Cl2/KCl electrode or the Ag/AgCl/KCl electrode, and although potentials are frequently expressed with reference to the standard hydrogen electrode (S.H.E.) the use of this electrode in practice is confined to fundamental studies rather than testing. 

It is apparent that since the electrode potential of a metal/solution interface can only be evaluated from the e.m.f. of a cell, the reference electrode used for that purpose must be specified precisely, e.g. the criterion for the cathodic protection of steel is —0-85 V (vs. Cu/CuSOg, sat.), but this can be expressed as a potential with respect to the standard hydrogen electrode (S.H.E.), i.e. -0-55 V (vs. S.H.E.) or with respect to any other reference electrode. Potentials of reference electrodes are given in Table 21.7. 

Table 21.6 Standard elearode potentials against the standard hydrogen electrode for inorganic systems at 25°Ct...

Electrode Potential (E) the difference in electrical potential between an electrode and the electrolyte with which it is in contact. It is best given with reference to the standard hydrogen electrode (S.H.E.), when it is equal in magnitude to the e.m.f. of a cell consisting of the electrode and the S.H.E. (with any liquid-junction potential eliminated). When in such a cell the electrode is the cathode, its electrode potential is positive when the electrode is the anode, its electrode potential is negative. When the species undergoing the reaction are in their standard states, E =, the stan-... 

Hydrogen Electrode an electrode at which the equilibrium (aq.) + jHj, is established. By definition, at unit activity of hydrogen ions and unit fugacity of hydrogen gas the potential of the standard hydrogen electrode h+/y//2 =... 

When the activity of the ion M"+ is equal to unity (approximately true for a 1M solution), the electrode potential E is equal to the standard potential Ee. Some important standard electrode potentials referred to the standard hydrogen electrode at 25 °C (in aqueous solution) are collected in Table 2.5.5... 

It may be noted that the standard hydrogen electrode is rather difficult to manipulate. In practice, electrode potentials on the hydrogen scale are usually... 

The most widely used reference electrode, due to its ease of preparation and constancy of potential, is the calomel electrode. A calomel half-cell is one in which mercury and calomel [mercury(I) chloride] are covered with potassium chloride solution of definite concentration this may be 0.1 M, 1M, or saturated. These electrodes are referred to as the decimolar, the molar and the saturated calomel electrode (S.C.E.) and have the potentials, relative to the standard hydrogen electrode at 25 °C, of 0.3358,0.2824 and 0.2444 volt. Of these electrodes the S.C.E. is most commonly used, largely because of the suppressive effect of saturated potassium chloride solution on liquid junction potentials. However, this electrode suffers from the drawback that its potential varies rapidly with alteration in temperature owing to changes in the solubility of potassium chloride, and restoration of a stable potential may be slow owing to the disturbance of the calomel-potassium chloride equilibrium. The potentials of the decimolar and molar electrodes are less affected by change in temperature and are to be preferred in cases where accurate values of electrode potentials are required. The electrode reaction is... 

In general, the baser the metal, the lower (more negative) the electrical potential at the anode and the higher the potential rate of corrosion. Carbon steel and low-alloy steels (which are widely used in boiler plants) have a relatively low potential with respect to the standard hydrogen electrode and can therefore be expected to corrode readily unless active prevention measures are taken. Copper and brasses have a relatively higher potential. 

Electrode potentials are customarily tabulated on the standard hydrogen electrode (SHE) scale (although the SHE is never actually used experimentally because it is inconvenient in many respects). Therefore, conversion of potentials into the UHV scale requires the determination of E°(H+/H2) vs. UHV. According to the concepts developed above, such a potential would measure the energy of electrons in the Pt wire of the hydrogen electrode, modified by the contact with the solution. 

Experimental Values of the Potential of the Standard Hydrogen Electrode in the UHV Scale... 

Figure 7.12 shows the relationship between the standard oxygen electrode (soe) scale of solid state electrochemistry, the corresponding standard hydrogen electrode (she) scale of solid state electrochemistry, the standard hydrogen electrode (she) scale of aqueous electrochemistry, and the physical absolute electrode scale. The first two scales refer to a standard temperature of 673.15 K, the third to 298.15 K. In constructing Figure 7.12 we have used the she aqueous electrochemical scale as presented by Trasatti.14... 

Under open-circuit conditions the catalyst potential UwR=Urhe takes values of the order 0.4-0.85 V, that is -0.35 to +0.1 V on the standard hydrogen electrode scale (she), depending on the hydrogen to oxygen ratio. 

In redox couple notation, E°(HJ"/H2) = 0 at all temperatures. A hydrogen electrode in its standard state, with hydrogen gas at 1 bar and the hydrogen ions present at 1 mol-L 1 (strictly, unit activity), is called a standard hydrogen electrode (SHE). The standard hydrogen electrode is then used to define the standard potentials of all other electrodes ... 

When the two electrodes are connected, current flows from M to X in the external circuit. When the electrode corresponding to half-reaction 1 is connected to the standard hydrogen electrode (SHE), current flows from M to SHE. (a) What are the signs of ° of the two half-reactions (b) What is the standard cell potential for the cell constructed from these two electrodes ... 

All values taken from literature sources are given on the standard hydrogen electrode scale. If necessary, conversions were based on standard values for the involved reference electrodes. If a single report contains several values measured with different concentrations of the electrolyte under investigation, all values are included in the table in order to allow insight into the influence of the concentration. 

All reported values are converted with respect to the standard hydrogen electrode for the solvent rmder investigation if not stated otherwise. 

---