Hydrogen-electrode concentration cell

It is highly desirable to establish a set of pH buffer solutions which can be used at temperatures above 100 °C. Thus far, little has been done to develop the necessary sets of the high-temperature buffer systems as primary standards. Only a 0.05 mol-kg potassium hydrogen phthalate solution has been adopted by lUPAC as an appropriate primary buffer system to be used at temperatures up to about 225 °C. However, the acid dissociation constants of many organic and inorganic buffers have been measured with the hydrogen-electrode concentration cell (see discussion below) and these results are currently available for developing the secondary pH standards to 250 °C. 

The available experimental systems for potentiometric measmements, as well as for measmements of the electrochemical reaction rates as a function of temperature are represented here. 

Hydrogen-Electrode Concentration Cell for Solubility Measurements 

It is immediately apparent from Equation (3.8) that knowledge of the desired quantity or dependent variable, is based on the measured value of f HEcc, the calculated value of Eo, and the known molality, ( fin)i, in the reference compartment, where the latter is corrected for loss of water vapor to the head space when a static cell is used. Therefore, the pH obtained from these measurements is in fact defined on the molality rather than the commonly adopted activity scale according to Equation (3.9) 

The conventional pH defined on the activity scale is related to pHm by the relationship pH = pH - log Yh+- The precision of these pH measmements is generally 0.01 for well-behaved electrolyte solutions. 

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An electrode concentration cell can be constructed using two electrodes which consist of a gas at different partial pressures. The following is an example of the hydrogen electrode concentration cell ... 

Ridley, M.K. et al., Poteirtiometric studies of the rutile-water interface Hydrogen-electrode concentration-cell versus glass-electrode titrations. Colloids Surf. A, 204, 295, 2002. 

Mesmer and Baes (P6) studied the dissociation equilibria potentiometrically over a wide temperature range. Using a hydrogen-electrode concentration cell measurements were made to SOO C. The equilibria were expressed as neutralization reactions ... 

Therefore the setup where the dissolution reaction is followed by potentiometric measurements of the pertinent H -concentration within the galvanic cell (glass electrode solution salt bridge reference electrode) is an optimal system for solubility measurements. This method is called potentiometric method of solubility measiuements ( Potentio in Table 1.1). An example of such measurements could be tiiat of solubility of metal oxides at temperatures up to 300°C performed in ORNL using simultaneously sampling method and the hydrogen-electrode concentration cell (HECC). The details of such measurements and experimental setup can be found in Chapter 3 in this Book. 

As has been pointed out, another interesting type of electrode-concentration cell is that comprising two hydrogen electrodes working at different pressures and remaining immersed in a solution of hydrochloric acid. The cell may be represented as ... 

The voltage of any half cell can be recorded against a standard hydrogen electrode (half cell). Table 2.1 gives the standard half cell potentials that are of interest to us as we evaluate corrosion problems. Half cell potentials are a function of concentration as well as the metal and the solution. A more concentrated solution is generally) more corrosive than a dilute one so a current will flow in a cell made up of a single metal in two different concentrations of the same solution. We can consider the corrosion of steel in concrete as a concentration cell. 

A particular concentration measure of acidity of aqueous solutions is pH which usually is regarded as the common logarithm of the reciprocal of the hydrogen-ion concentration (see Hydrogen-ION activity). More precisely, the potential difference of the hydrogen electrode in normal acid and in normal alkah solution (—0.828 V at 25°C) is divided into 14 equal parts or pH units each pH unit is 0.0591 V. Operationally, pH is defined by pH = pH(soln) + E/K, where E is the emf of the cell ... 

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... 

To measure the hydrogen ion concentration of a solution the glass electrode must be combined with a reference electrode, for which purpose the saturated calomel electrode is most commonly used, thus giving the cell ... 

If the concentrations in the cell are such that it is reported as having a positive emf (that is, the mercury/mercury(I) chloride electrode is found to be positive), then the reaction as written is spontaneous. If the concentrations were such that the emf were reported as negative (that is, the hydrogen electrode were found to be positive), then the reverse of the reaction that we have derived would be spontaneous. 

In the discussion of the Daniell cell, we indicated that this cell produces a voltage of 1.10 V. This voltage is really the difference in potential between the two half-cells. The cell potential (really the half-cell potentials) is dependent upon concentration and temperature, but initially we ll simply look at the half-cell potentials at the standard state of 298 K (25°C) and all components in their standard states (1M concentration of all solutions, 1 atm pressure for any gases and pure solid electrodes). Half-cell potentials appear in tables as the reduction potentials, that is, the potentials associated with the reduction reaction. We define the hydrogen half-reaction (2H+(aq) + 2e - H2(g)) as the standard and has been given a value of exactly 0.00 V. We measure all the other half-reactions relative to it some are positive and some are negative. Find the table of standard reduction potentials in your textbook. 

Sorensen is usually considered to be the first to have realized the importance of hydrogen ion concentration in cells and in the solutions in which the properties of cell components were to be studied. He is also credited with the introduction of the pH scale. Electrochemistry started at the end of the nineteenth century. By 1909, Sorensen had introduced a series of dyes whose color changes were related to the pH of the solution, which was determined by the H+ electrode. The dyes were salts of weak acids or weak bases. He also devised simple methods for preparing phosphate buffer solutions covering the pH range 6-8. Eventually buffers and indicators were provided covering virtually the whole pH range. 

It is possible to select a cell that contains a weak acid in solution whose potential depends on the ion concentrations in the solution and hence on the dissociation constant of the acid. As an example, we will consider acetic acid in a cell that contains a hydrogen electrode and a silver-silver chloride electrode ... 

SHE, standard hydrogen electrode The electrode used as a standard against which aU other half-cell potentials are measured. The following reaction occurs at the platinum electrode when immersed in an acidic solution and cormected to the other half of an electrochemical cell 2H (aq) -H 2e —> H2(g). The half- cell potential of this reaction at 25°C, 1 atm and 1 m concentrations of aU solutes is agreed, by convention, to be OV... 

E is the standard equilibrium potential, i. e. the potential corresponding to unit activity and RTF. The dissolution reaction leads to the development of an electrical double layer at the iron-solution interface. The potential difference of the Fe/Fe " half cell cannot be measured directly, but if the iron electrode is coupled with a reference electrode (usually the standard hydrogen electrode, SHE), a relative potential difference, E, can be measured. This potential is termed the single potential of the Fe/Fe electrode on the scale of the standard hydrogen couple H2/H, the standard potential of which is taken as zero. The value of the equilibrium potential of an electrochemical cell depends upon the concentrations of the species involved. 

Accurate and precise measures of pH can be made with a pH meter. Typical pH meters usually contain a glass electrode and reference electrode arranged similar to an electrochemical cell. We discuss electrochemical cells in Chapter 14. For now, though, consider a pH meter as essentially a modified voltmeter in which the voltage measured is directly proportional to the hydrogen ion concentration of the solution. Simple pH meters are capable of measuring pH within 0.1 pH units, while more sophisticated instruments are precise to within 0.001 pH units. 

Chapter 2) apply. The standard reference half-cell is reaction 15.6, the standard hydrogen electrode (SHE), and the standard conditions are those listed in Section 2.3, although for our purposes the molar concentration scale (mol L 1) can generally be used without significant loss of precision. We will simplify matters further, for illustrative purposes, by equating activities with molar concentrations our numerical results will therefore be only approximate, except where concentrations are very low. A thermodynamically acceptable treatment would require the calculation or measurement of ionic activities or, at the very least, maintenance of constant ionic strength, as outlined in Section 2.2. 

Many half-reactions of interest to biochemists involve protons. As in the definition of AG °, biochemists define the standard state for oxidation-reduction reactions as pH 7 and express reduction potential as E °, the standard reduction potential at pH 7. The standard reduction potentials given in Table 13-7 and used throughout this book are values for E ° and are therefore valid only for systems at neutral pH Each value represents the potential difference when the conjugate redox pair, at 1 m concentrations and pH 7, is connected with the standard (pH 0) hydrogen electrode. Notice in Table 13-7 that when the conjugate pair 2ET/H2 at pH 7 is connected with the standard hydrogen electrode (pH 0), electrons tend to flow from the pH 7 cell to the standard (pH 0) cell the measured E ° for the 2ET/H2 pair is -0.414 V... 

The standard reduction potential would be observed if the half-cell of interest (with unit activities) were connected to a standard hydrogen electrode, as it is in Figure 14-7. It is nearly impossible to construct such a cell, because we have no way to adjust concentrations and ionic strength to give unit activities. In reality, activities less than unity are used in each half-cell, and the Nemst equation is employed to extract the value of E° from the cell voltage.12 In the hydrogen electrode, standard buffers with known pH (Table 15-3) are used to obtain known activities of H+. 

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