Emf determination

Which is the more powerful oxidizing agent under standard conditions, an acidified aqueous permanganate solution or an acidified aqueous dichromate solution Specify the cell for the spontaneous reaction of the two couples by writing a cell diagram that under standard conditions has a positive emf. Determine the standard emf of the cell and write the net ionic equation for the spontaneous cell reaction. 

Experimentally, the home-made Ag AgCl electrode and another reference electrode (e.g. an SCE) are immersed in solutions of varying [Cr ] and the emf determined for each concentration, e.g. in the range 0.001-0.1 mol dm. Hydrochloric acid is a suitable source of chloride ion for this experiment. Respective values of EAgcuAg tire then obtained from each emf via the use of equation (3.3). 

The main goal will be the determination of the apparent charge of the micelle in the binary aqueous solvents. Contradictory findings are found in the literature when the apparent charge is deduced from emf determinations (12,13). We shall discuss these results and compare them with those obtained using a conductivity approach. 

Experimentally, the emf can be measured either by compensating the circuit voltage (classical technique which became rare nowadays) or by using a voltmeter of very high internal resistance. The accuracy of emf determination of about 1 xV can be achieved in precise measurements, whereas common devices provide the accuracy of about 1 mV. The potential unit named Volt, which is used in the modern literature (particularly, below), is the so-called absolute (abs) Volt it differs slightly from the international (int) Volt value. The ratio abs/int is 1.00033. To determine the sign of emf, a conventional rule is adopted, which states that the left electrode should be considered as the reference one. 

Thus, it appears imperative to complement equilibrium measurements, such as cell emf determinations, by transient electrochemical investigations (as well as nonelectrochemical or spectroelectrochemical techniques, where possible), in order to definitively establish the nature of the process being examined. 

Some EMF-determined standard Gibbs-energy functions (A,G° eq. 2) of formation of RCUO2 from oxides ... 

EMF-determined standard Gibbs-energy fimctions (A G eq. 2) of formation of (La, j,Sr )2Cu04 from (1 - )La203 +2xSrO + CuO+ according to Idemoto et al. (1993) units have been omitted... 

Salt bridges are useful for interconnecting two half-cells to form a reference cell, as shown in Figure 5.9, or for interconnecting a hydrogen reference half-cell with another half-cell for emf determination, as shown in Figure 5.8. 

Afly j I — jother methods include emf determinations, taken up in the next chapter. 

Of all the techniques, it is those of Group 1 that are likely to give the most realistic data, simply because they measure transport of charged species only. They are not the easiest experimental techniques to perform on polymeric systems and this probably explains why so few studies have been undertaken. The experimental difficulties associated with the Tubandt-Hittorf method are in maintaining nonadherent thin-film compartments. One way is to use crosslinked films [79], while an alternative has been to use a redesigned Hittorf cell [80]. Although very succesful experimentally, the latter has analytical problems. Likewise, emf measurements can be performed with relative ease [81, 82] it is the necessary determination of activity coefficients that is difficult. 

The determination of Ka requires a measurement using a technique such as electrical conductivity, absorption of light, or, in our case, the emf of an electrochemical cell. When K.x is determined, ArG° is obtained from equation (9.116), and tables giving AtG° must take into account this value. For example, = 1.75 x 10 5 at 298.15 K and A,-G° = 27.15 kj-mol-1 for the reaction1 ... 

R. Gomer, and G. Tryson, An experimental determination of absolute half-cell emfs and single ion free energies of solvation, J. Chem. Phys. 66, 4413-4424 (1977). 

SOLUTION Use Eq. 1 to determine a reaction Gibbs free energy—a thermodynamic quantity—from a cell emf—an electrical quantity. From the chemical equation for the cell reaction (reaction A), we see that n = 2 mol. 

SOLUTION We can determine the standard potential of an electrode by measuring the emf of a standard cell in which the other electrode has a known standard potential and applying Kq. 3. 

The value of Kip is the same as that listed in Table 11.4. Many of the solubility products listed in tables were determined from emf measurements and calculations like this one. 

STRATEGY First, write the balanced equation for the cell reaction and the corresponding expression for Q, and note the value of n. Then determine E° from the standard potentials in Table 12.1 or Appendix 2B. Determine the value of Q for the stated conditions. Calculate the emf by substituting these values into the Nernst equation, Eq. 6. At 25.00°C, RT/1 = 0.025 693 V. 

Therefore, by measuring the cell emf, E, we can determine the pH. The value of E is established by calibrating the cell, which requires measuring E for a solution of known pH. 

Determine the standard potential of an electrode from a cell emf (Example 12.5). 

For each reaction that is spontaneous under standard conditions, write a cell diagram, determine the standard cell emf, and calculate AG° for the reaction ... 

Note that a number of complicating factors have been left out for clarity For instance, in the EMF equation, activities instead of concentrations should be used. Activities are related to concentrations by a multiplicative activity coefficient that itself is sensitive to the concentrations of all ions in the solution. The reference electrode necessary to close the circuit also generates a (diffusion) potential that is a complex function of activities and ion mobilities. Furthermore, the slope S of the electrode function is an experimentally determined parameter subject to error. The essential point, though, is that the DVM-clipped voltages appear in the exponent and that cheap equipment extracts a heavy price in terms of accuracy and precision (viz. quantization noise such an instrument typically displays the result in a 1 mV, 0.1 mV, 0.01 mV, or 0.001 mV format a two-decimal instrument clips a 345.678. .. mV result to 345.67 mV, that is it does not round up ... 78 to ... 8 ). 

For obtaining internal or external mobilities, the corresponding transport numbers are usually measured. There are several methods for determining transport numbers in molten salts that is, the Kleimn method (countercurrent electromigration method or column method), the Hittorf method (disk method), the zone electromigration method (layer method), the emf method, and the moving boundary method. These are described in a comprehensive review. ... 

The EMF values of galvanic cells and the electrode potentials are usually determined isothermally, when all parts of the cell, particularly the two electrode-electrolyte interfaces, are at the same temperature. The EMF values will change when this temperature is varied. According to the well-known thermodynamic Gibbs-Helmholtz equation, which for electrochemical systems can be written as... 

Direct measurements of solute activity are based on studies of the equilibria in which a given substance is involved. The parameters of these equilibria (the distribution coefficients, equilibrium constants, and EMF of galvanic cells) are determined at different concentrations. Then these data are extrapolated to very low concentrations, where the activity coincides with concentration and the activity coefficient becomes unity. 

Comparison of the calculated and observed changes in the EMF is shown in Table 1. It can be seen that the calculated changes in the phase boundary potential of membranes with 1.0 mM 1-3 in contact with 0.1 and 0.01 M aqueous KCl or RbCl were in good agreement with the corresponding observed values. Such an agreement indicates that it is reasonable to apply the present surface model to explain that the phase boundary potential is, in fact, determined by the amount of the primary cation permeated into or released out of the membrane side of the interface. 

To determine the influence of ionic sites on the charge separation at the membrane interface, we have measured in this study SHG with ionophore-free and ionophore-incorporated liquid membranes in absence and presence of ionic sites. The dependence of the SHG intensity on the activity of the primary ion in the aqueous solution is presented and compared to the corresponding EMF. 

---