Electrolysis experiment

An electrolysis experiment is performed to determine the value of the Faraday constant (number of coulombs per mole of electrons). In this experiment, 28.8 g of gold is plated out from a AuCN solution by running an electrolytic cell for two hours with a current of 2.00 A. What is the experimental value obtained for the Faraday constant ... 

Charles Hall was inspired by his chemistry professor at Oberlin College, who observed that whoever perfected an inexpensive way of producing aluminum would become rich and famous. After his graduation. Hall set to work in his home laboratory, trying to electrolyze various compounds of aluminum. He was aided by his sister Julia, who had studied chemistry and shared Charles interests. Julia helped to prepare chemicals and witnessed many of the electrolysis experiments. After only eight months of work. Hall had successfully produced globules of the metal. Meanwhile, Heroult was developing the identical process in France. 

Figure 3.43 Schematic illustration of the transition time t in a constant-current (chronopotentin-metric) electrolysis experiment (see text for details).

The solution used for all experimentation was 5 M NaCl, made up using triply distilled water. Electrolysis experiments were carried out in a one-compartment cell at either room temperature or at 90°C. All electrochemical experiments were carried out using a two-compartment cell, with the RE compartment connected to the WE and CE compartments via a Luggin capillary. 

From an exhaustive potentiostatic electrolysis, the product(s) formed at the selected electrode potential can be isolated. Preparative and analytical techniques are available to determine the composition of the product mixture and the structure of its components. Mechanistic reasoning will often allow defining the reaction steps. Even more information about the reaction can be gained from electrolysis experiments at various defined potentials, for example, after each peak in the cyclic voltammogram of the substrate. 

A typical Kolbe electrolysis experiment calls for the constant current electrolysis of... 

Industrially, the silver is recovered from either the wash water, or the bleach fix separately or from a mixture of the two using electrolysis employing a stainless steel cathode cylinder and an anode of stainless steel mesh. A typical wash solution composition contains silver (4 g L ), sodium thiosulphate (220 g L ), sodium bisulphite (22 g L ) and sodium ferric EDTA (4 g L ). At Coventry we have used a scaled down version of the industrial process employing 250 mL samples [46]. Electrolysis experiments were performed at ambient temperature with both wash and bleach fix solutions and in which the potential applied to the cathode and the speed of rotation of the cathode were varied. The sonic energy (30 W) was supplied by a 38 kHz bath. The results are given in Tab. 6.9. The table shows that the recovery of silver on sonication of the wash or bleach fix solutions is much improved especially if the electrode is rotated while ultrasound is applied. Yields with bleach fix (which contains ferric ions) are less since Fe and Ag compete for discharge (Eqs. 6.13 and 6.14). 

In a recent study, the analysis of the reaction products at the outlet of the anode compartment of a DEFC fitted with a Pt/C anode showed that only AA, AAL and CO2 could be detected by HPLC [24]. Depending on the electrode potential, AAL, AA, CO2 and traces of CH4 are observed, the main products being AAL and AA (Table 1.2), AA being considered as a final product because it is not oxidized under smooth conditions. Long-term electrolysis experiments on a Pt catalyst show that AAL is detected at potentials as low as 0.35 V vs RHE, whereas no AA was detected in this potential range. 

This mechanism explains also the higher efficiency of Pt-Sn in forming AA compared with Pt at low potentials ( <0.35 V vs RHE), as was shown by electrolysis experiments. Indeed, adsorbed OH species on Sn atoms can be used to oxidize adsorbed CO species to CO2 or to oxidize adsorbed -COCH3 species to CH3COOH, according to the bifunctional mechanism [14]. 

Costogue EN, Yasui RK (1977) Performance data for a terrestrial solar photovoltaic/water electrolysis experiment. Sol Energy 19 205-210... 

Fulleride anions are often more soluble, especially in more polar solvents, than the parent fullerenes. For example, in bulk electrolysis experiments with tetra-n-butylammonium perchlorate (TBACIO4) as supporting electrolyte, carried out in acetonitrile where Cjq is completely insoluble, fairly concentrated, dark red-brown solutions of 50 can be obtained [81]. Upon reoxidation, a quantitative deposition of a neutral Cjq film on the surface of a gold/quartz crystal working electrode takes place. This Cjq film can be stepwise reductively doped with TBA, leading to (Cjo )... 

A student prepares an aqueous solution of NaCl and performs an electrolysis experiment. The products at the cathode would be... 

Problem 21.7 (a) How does one repress the ionization of an a-amino acid (b) The pH at which [anion] = [cation] is called the isoelectric point. In an electrolysis experiment, do a-amino acids migrate at the isoelectric p>oint ... 

The electrochemical and chemical behavior of rotaxane 7 + was analyzed by CV and controlled potential electrolysis experiments.34,35 From the CV measurements at different scan rates (from 0.005 to 2 V/s) both on the copper(I) and on the copper(II) species, it could be inferred that the chemical steps (motions of the ring from the phenanthroline to the terpyridine and vice versa) are slow on the timescale of the experiments. As the two redox couples involved in these systems are separated by 0.7 V, the concentrations of the species in each environment (tetra- or pentacoor-dination) are directly deduced from the peak intensities of the redox signals. In Fig. 14.13 are displayed some voltammograms (curves a-e) obtained on different oxidation states of the rotaxane 7 and at different times. 

A sample of 25 ml prepared for an electrolysis experiment has a zinc concentration of approximately 2 x 10 8 M which leads to the passage of a current of 1.5 nA. Calculate the time necessary to deposit 3% of the Zn present. 

The fuel cell was discovered during electrolysis experiments with water. It is the reverse process which produces the electricity. Write a balanced chemical equation to represent the overall reaction taking place in a fuel cell. 

The electrochemical behavior of 14+ is particularly clean and interesting, since only the 4- and the 5-coordinate geometries can be obtained on translating the metal-complexed ring from the phen site to the terpy site)841 The electrochemically induced molecular motions (square scheme1851), similar to those represented in Figure 10 but now involving stopped compounds, can be monitored by cyclic voltammetry (CV) and controlled potential electrolysis experiments)851... 

Figure 1 Experimental set-up used for single-cell electrolysis experiments...

When the low-flow rate single-cell electrolysis cell was used the cathode and anode were fabricated by painting one side of a 5 cm2 piece of EC-TP1-060 Toray Carbon Fiber Paper with Pt on XC-72R catalyst mixture. In high-flow rate single-cell electrolysis experiments the cathode was a... 

Figure 2 Photograph of the single-cell commercial fuel cell used for low flow CuCI/HCI electrolysis experiments...

The equilibrium potential of the single-cell electrolysis cell, as determined from polarisation curves, increases during single-cell electrolysis experiments. 

Copper ions, presumably copper(II), enter the catholyte solution from the anolyte solution by crossing the membrane during electrolysis experiments. 

Figure 8 Outlet gas composition as a function of current density for co-electrolysis experiments, 10-cell stack...

The electrolysis of the studied systems was carried out in the same cell as voltammetry measurements under the mode of either constant current or voltage. In the constant current mode, the applied current density was in the range of 0.01 0.2 A/ sm2 with reference to the surface area of the cathode before starting the electrolysis. Semi-immersed glassy carbon plate electrodes (cathode area - 5 sm2, anode area - 10 sm2) were used while electrolysis experiments. A powder product was either settled down onto the crucible bottom or assembled on the cathode in the view of electrolytic pear . The deposit was separated from salts by successive leaching with hot water. Thereafter, the precipitate was washed with distilled water by decantation method several times and dried to a constant mass at 100 - 150 °C. The electrolysis products were analyzed by chemical and X-ray phase analyses, methods of electron diffraction and electronic microscopy (transmission and scanning). 

An explosion (appropriately enough, in a military laboratory) has actually been reported to have occurred during an electrolysis experiment in which a perchlorate was used as the supporting electrolyte (Titus, 1971). Whatever the chemistry involved, this accident merits attention from large-scale practitioners of organic electrolysis. 

Coal and many coal-derived liquids contain polycyclic aromatic structures, whose molecular equivalents form radical cations at anodes and radical anions at cathodes. ESR-electrolysis experiments support this (14). Chemically, radical cations form by action of H2SO4 (15,19), acidic media containing oxidizing agents (15,20,21,22), Lewis acid media (18,23-35) halogens (36), iodine and AgC104 (37,38), and metal salts (39,40). They also form by photoionization (41,42,43) and on such solid catalytic surfaces as gamma-alumina (44), silica-alumina (45), and zeolites (46). Radical anions form in the presence of active metals (76). 

Panizza, M. and Cerisola, G. (2004b) Influence of anode material on the electrochemical oxidation of 2-naphthol Part 2. Bulk electrolysis experiments. Electrochim. Acta 49, 3221-3226. 

Figure 7.6 was obtained by carrying out electrolysis experiments at extremely low chloride concentrations. Both curves show a tendency of chlorine formation and destruction in terms of a spectrophotometrical DPD signal. Even if there are some uncertainties with respect to the DPD method (see Sect. 7.3.3.7) these results support the second theory (2). 

In fact, active chlorine consumption was observed in disturbed experiments with higher C102 concentration. Chemical formation from active chlorine and chlorite is also favoured due to the reaction behaviour after switching off the electrolysis experiments (Fig 7.11). It can often be seen from the UV spectra that the hypochlorite peak decreases and the C102 peak slightly increases. Increase of C102 concentra-... 

Panizza, M. and Cerisola, G. (2004) Influence of anode material on the electrochemical oxidation of 2-naphthol. Part 2. Bulk electrolysis experiments. Electrochim. Acta 49, 3221-3226 Panizza, M., Delucchi M., and Cerisola, G. (2005) Electrochemical degradation of anionic surfactants, J. Appl. Electrochem. 35,357-361... 

Fleischmann et al. [549] studied the electro-oxidation of a series of amines and alcohols at Cu, Co, and Ag anodes in conjunction with the previously described work for Ni anodes in base. In cyclic voltammetry experiments, conducted at low to moderate sweep rates, organic oxidation waves were observed superimposed on the peaks associated with the surface transitions, Ni(II) - Ni(III), Co(II) -> Co(III), Ag(I) - Ag(II), and Cu(II) - Cu(III). These observations are in accord with an electrogenerated higher oxide species chemically oxidizing the organic compound in a manner similar to eqns. (112) (114). For alcohol oxidation, the rate constants decreased in the order kCn > km > kAg > kCo. Fleischmann et al. [549] observed that the rate of anodic oxidations increases across the first row of the transition metals series. These authors observed that the products of their electrolysis experiments were essentially identical to those obtained in heterogeneous reactions with the corresponding bulk oxides. 

In order to prevent contamination with products formed at the counter electrode or reaction of these products with the electroactive solution of interest, a salt bridge may be used in conjunction with the counter electrode, e.g. this is essential in bulk electrolysis experiments. 

Fortunately, the low enol content in simple ketone systems does not necessarily impose an obstacle to generating the corresponding enol radical cations in solution. As outlined in Sect. 2 the selective oxidation of the enol tautomer even in the presence of a vast excess of the ketone opens up an indirect, but quantitative access to enol radical cation intermediates for all systems, if an appropriate oxidant has been chosen. The first, albeit indirect evidence for this selective oxidation step stems from kinetic studies by Henry [109] and Littler [110-112] and will be discussed in more detail in Sect. 3.3. Direct evidence for a specific oxidation of enols was provided by Orliac-Le Moing and Simonet [108]. Using voltammetry at a rotating disc electrode they were able to establish a linear correlation between the anodic current and the enol content for various a-cyano ketones 11. In electrolysis experiments the corresponding 1,4-diketones 13 were obtained in high current yield (ca. 90%). 

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