Beaker cells

Tank Cells. A direct extension of laboratory beaker cells is represented in the use of plate electrodes immersed into a lined, rectangular tank, which may be fitted with a cover for gas collection or vapor control. The tank cell, which is usually undivided, is used in batch or semibatch operations. The tank cell has the attraction of being both simple to design and usually inexpensive. However, it is not the most suitable for large-scale operation or where forced convection is needed. Rotating cylinders or rotating disks have been used to overcome mass-transfer problems in tank cells. An example for electroorganic synthesis is available (46). 

Less than 10% of the reports on pyridine electrochemistry deal with anodic reactions. The mechanisms of these reactions are rarely known and, as a result, yields or current efficiencies have not always been optimized. Many of the anodic reactions were studied in beaker cells, which are simply not good models for modern flow cells moreover, uncontrolled power supplies were often used. Consequently, anode overpolarization caused ring degradation in many cases. 

Bulk Liquid Membranes (BLM). This is the simplest type of liquid membrane (2-8) and is utilized for fundamental studies of certain aspects of liquid membrane transport processes. In one such process, a beaker-in-a-beaker cell (Figure 1) consists of inner and outer compartments which contain the aqueous feed (F) and strip (S) solutions, respectively. The inner beaker contains the stripping solution and is surroimded by the feed solution. Both aqueous solutions contact the upper organic layer, which is the liquid membrane. Mass transfer takes place from the feed solution through the liquid membrane and into the strip solution. Bartsch et aL studied the transport of alkali metal cations across bulk liquid membranes in which a crown ether carboxylic acid in the organic layer served as the carrier (2,3). 

SiNWs in beaker cell [c] voltage profile and (d) galvanostatic cycling results for SiNWs in coffee bag cell using C/20 rate in 1 M LiPFg/ECiDEC electrolyte [Ref. 65). 

This method involves very simple and inexpensive equipment that could be set up m any laboratory [9, 10] The equipment consists of a 250-mL beaker (used as an external half-cell), two platinum foil electrodes, a glass tube with asbestos fiber sealed m the bottom (used as an internal half-cell), a microburet, a stirrer, and a portable potentiometer The asbestos fiber may be substituted with a membrane This method has been used to determine the fluoride ion concentration in many binary and complex fluondes and has been applied to unbuffered solutions from Willard-Winter distillation, to lon-exchange eluant, and to pyrohydrolysis distil lates obtained from oxygen-flask or tube combustions The solution concentrations range from 0 1 to 5 X 10 M This method is based on complexing by fluonde ions of one of the oxidation states of the redox couple, and the potential difference measured is that between the two half-cells Initially, each cell contains the same ratio of cerium(IV) and cerium(tll) ions... 

It consists of a porous cell which foims the cathode chamber and contains 20 grams nitrobenzene and 160 grains 2 5 per cent, caustic soda solution. The two aie kept well mixed throughout the operation by a rapidly revolving stinei. The cathode is a cylinder of nickel gauze (12 cms. x S 5 cms. = too sq. cms.). The anode chamber is the outer glass vessel or beaker. 

During the operation of the cell (or during the direct interaction of zinc metal and cupric ions in a beaker) the zinc is oxidised to Zn and corrodes, and the Daniell cell has been widely used to illustrate the electrochemical mechanism of corrosion. This analogy between the Daniell cell and a corrosion cell is perhaps unfortunate, since it tends to create the impression that corrosion occurs only when two dissimilar metals are placed in contact and that the electrodes are always physically separable. Furthermore, although reduction of Cu (aq.) does occur in certain corrosion reactions it is of less importance than reduction of HjO ions or dissolved oxygen. 

Let s begin our investigation of an electrochemical cell by assembling one. Fill a beaker with a dilute solution of silver nitrate (about 0.1 M will do) and another beaker with dilute copper sulfate. Put a silver rod in the AgN03 solution and a copper rod in the CuSO< solution. With a wire, connect the silver rod to one terminal of an... 

The overall reaction describes what goes on in the entire electrochemical cell. In half of the cell, the right beaker, reaction (7) occurs. In the other half of the cell, the left beaker, reaction (2) occurs. Hence, reactions (7) and (2) are called half-cell reactions or half-reactions. 

The two half-reactions are written separately. In our electrochemical cell the half-reactions occur in separate beakers. As the name implies, there must be two such reactions. 

These ideas, developed for an electrochemical cell, have great importance in chemistry because they are also applicable to chemical reactions that occur in a single beaker. Without an electric circuit or an opportunity for electric current to flow, the chemical changes that occur in a cell can be duplicated in a single solution. It is reasonable to apply the same explanation. 

The moles of silver deposited per mole of copper dissolved are the same whether reaction (J) is carried out in an electrochemical cell or in a single beaker, as in Experiment 7. If, in the cell, electrons are transferred from copper metal (forming Cu+2) to silver ion (forming metallic silver), then electrons must have been transferred from copper metal to silver ion in Experiment 7. 

Thus, Experiment 7 involved the same oxidation-reduction reaction but the electron transfer must have occurred locally between individual copper atoms (in the metal) and individual silver ions (in the solution near the metal surface). This local transfer replaces the wire middleman in the cell, which carries electrons from one beaker (where they are released by copper) to the other (where they are accepted by silver ions). 

Suppose water is added to each of the beakers containing copper sulfate in the two electrochemical cells shown in Figure 12-4 (p. 204). What change will occur in the voltage in each cdl Explain. 

Procedure. Place 50 mL of the supporting electrolyte in the beaker and add some of the same solution to the tube carrying the silver electrode so that the liquid level in this tube is just above the beaker. Pass nitrogen into the solution until the pH is 7.0. Pipette 10.00 mL of either 0.01 M or 0.001 M hydrochloric acid into the cell. Continue the passage of nitrogen. Proceed with the titration as described under (a) above. 

Prepare 250 mL of 0.02 M potassium dichromate solution and an equal volume of ca 0.1 M ammonium iron(II) sulphate solution the latter must contain sufficient dilute sulphuric acid to produce a clear solution, and the exact weight of ammonium iron(II) sulphate employed should be noted. Place 25 mL of the ammonium iron(II) sulphate solution in the beaker, add 25 mL of ca 2.5M sulphuric acid and 50 mL of water. Charge the burette with the 0.02 M potassium dichromate solution, and add a capillary extension tube. Use a bright platinum electrode as indicator electrode and an S.C.E. reference electrode. Set the stirrer in motion. Proceed with the titration as directed in Experiment 1. After each addition of the dichromate solution measure the e.m.f. of the cell. Determine the end point (1) from the potential-volume curve and (2) by the derivative method. Calculate the molarity of the ammonium iron(II) sulphate solution, and compare this with the value calculated from the actual weight of solid employed in preparing the solution. 

Procedure. Pipette 25.0 mL of the thiosulphate solution into the titration cell e.g. a 150mL Pyrex beaker. Insert two similar platinum wire or foil electrodes into the cell and connect to the apparatus of Fig. 16.17. Apply 0.10 volt across the electrodes. Adjust the range of the micro-ammeter to obtain full-scale deflection for a current of 10-25 milliamperes. Stir the solution with a magnetic stirrer. Add the iodine solution from a 5 mL semimicro burette slowly in the usual manner and read the current (galvanometer deflection) after each addition of the titrant. When the current begins to increase, stop the addition then add the titrant by small increments of 0.05 or 0.10 mL. Plot the titration graph, evaluate the end point, and calculate the concentration of the thiosulphate solution. It will be found that the current is fairly constant until the end point is approached and increases rapidly beyond. 

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