Electrode iodine

The indicator electrode must be reversible to one or the other of the ions which is being precipitated. Thus in the titration of a potassium iodide solution with standard silver nitrate solution, the electrode must be either a silver electrode or a platinum electrode in the presence of a little iodine (best introduced by adding a little of a freshly prepared alcoholic solution of iodine), i.e. an iodine electrode (reversible to I-). The exercise recommended is the standardisation of silver nitrate solution with pure sodium chloride. 

In this equation sJ> j2, j- represents the standard oxidation potential of the iodine electrode (sJ.tija j- — — 0.535 V) and aj- the activity of iodide ions in the solution. The negative value of this potential means that in combination with a hydrogen electrode a spontaneous reduction process will take place at the iodine electrode during which iodine will pass into the solution in ionic form. 

The value of the electromotive force arrived at is positive which suggests that the reaction, proceeding in the direction of the written equation, is a spontaneous one with the negative pole of the cell at the iodine electrode. 

Figure 12 Cell scheme for the use of the iodine electrode in molten Agl [108] 1, graphite electrode 2, silver electrode 3, thermocouple Ptl0%Rh-Pt 4, electrically heated tube, through which iodine vapors come from the generator 5, heated exit tube for iodine vapor 6, tube for introducing iodine into the cell and removing it 7, iodine collector.

The iodine electrode was also used by Sternberg, Adorian and Galasiu [111] to study the cell... 

Using the same iodine electrode, Sternberg et al. [115] studied the cells... 

The iodine level in solution determines the potential at the iodine electrode ... 

Affinity between Silver and Iodine.—When I compared the heat evolution and affinity in this case, greater differences appeared at first than were to be expected, on our Theorem, from the variation of the specific heats. I therefore set Herr Ulrich Fischer (74) the detailed examination of the case he showed conclusively that Thomsen s value for the heat of formation (13,800 cals.) was very considerably in error. From the temperature coefficient of the E.M.F. of the silver-iodine electrode, Fischer found 15,170 and, by... 

From Thomsen s determinations of the heat of formation of silver iodide it may be calculated, by means of the approximation formula, that iodine should have a measurable dissociation pressure at moderately high temperatures. Naumann (n) found, however, that this was too small to be measured even at 6oo°. Analogous differences occurred in the theoretical calculation of the E.M.F. of the silver-iodine electrode. U. Fischer then found the heat of formation by three independent methods as 15,200, 14,800, and 15,000, whereas Thomsen had given 13,800 so the differences between theory and observation are reconciled (cf. further, p. 113). 

Silver-silver bromide and silver-silver iodide electrodes prepared in a similar manner behave as reversible bromine and iodine electrodes, respectively. 

Chlorine has a lower electrode potential and electronegativity than fluorine but will displace bromine and iodine from aqueous solutions of bromide and iodide ions respectively ... 

Iodine has the lowest standard electrode potential of any of the common halogens (E = +0.54 V) and is consequently the least powerful oxidising agent. Indeed, the iodide ion can be oxidised to iodine by many reagents including air which will oxidise an acidified solution of iodide ions. However, iodine will oxidise arsenate(lll) to arsenate(V) in alkaline solution (the presence of sodium carbonate makes the solution sufficiently alkaline) but the reaction is reversible, for example by removal of iodine. 

The pH must be kept at 7.0—7.2 for this method to be quantitative and to give a stable end poiut. This condition is easily met by addition of soHd sodium bicarbonate to neutralize the HI formed. With starch as iudicator and an appropriate standardized iodine solution, this method is appHcable to both concentrated and dilute (to ca 50 ppm) hydraziue solutious. The iodiue solutiou is best standardized usiug mouohydraziuium sulfate or sodium thiosulfate. Using an iodide-selective electrode, low levels down to the ppb range are detectable (see Electro analytical techniques) (141,142). Potassium iodate (143,144), bromate (145), and permanganate (146) have also been employed as oxidants. 

Methods for iodine deterrnination in foods using colorimetry (95,96), ion-selective electrodes (94,97), micro acid digestion methods (98), and gas chromatography (99) suffer some limitations such as potential interferences, possibHity of contamination, and loss during analysis. More recendy neutron activation analysis, which is probably the most sensitive analytical technique for determining iodine, has also been used (100—102). 

C. HIO is prepared by oxidation of iodine with perchloric acid, nitric acid, or hydrogen peroxide or oxidation of iodine in aqueous suspension to iodic acid by silver nitrate. Iodic acid is also formed by anodic oxidation at a platinum electrode of iodine dissolved in hydrochloric acid (113,114). 

They form a monolayer that is rich in defects, but no second monolayer is observed. The interpretation of these results is not straightforward from a chemical point of view both the electrodeposition of low-valent Ge Iy species and the formation of Au-Ge or even Au Ge h compounds are possible. A similar result is obtained if the electrodeposition is performed from GeGl4. There, 250 20 pm high islands are also observed on the electrode surface. They can be oxidized reversibly and disappear completely from the surface. With Gel4 the oxidation is more complicated, because the electrode potential for the gold step oxidation is too close to that of the island electrodissolution, so that the two processes can hardly be distinguished. The gold step oxidation already occurs at -i-lO mV vs. the former open circuit potential, at h-485 mV the oxidation of iodide to iodine starts. 

Discussion. One of the most useful titrations involving iodine is that originally developed by Winkler18 to determine the amount of oxygen in samples of water. The dissolved oxygen content is not only important with respect to the species of aquatic life which can survive in the water, but is also a measure of its ability to oxidise organic impurities in the water (see also Section 10.103). Despite the advent of the oxygen-selective electrode (Section 16.36) direct titrations on water samples are still used extensively.19... 

Dilute solutions of sodium thiosulphate (e.g. 0.001 M) may be titrated with dilute iodine solutions (e.g. 0.005M) at zero applied voltage. For satisfactory results, the thiosulphate solution should be present in a supporting electrolyte which is 0.1 M in potassium chloride and 0.004 M in potassium iodide. Under these conditions no diffusion current is detected until after the equivalence point when excess of iodine is reduced at the electrode a reversed L-type of titration graph is obtained. 

Dilute solutions of iodine, e.g. 0.0001 M, may be titrated similarly with standard thiosulphate. The supporting electrolyte consists of 1.0 M hydrochloric acid and 0.004M potassium iodide. No external e.m.f. is required when an S.C.E. is employed as reference electrode. 

Procedure. Place 25.0 mL of the thiosulphate solution in the titration cell. Set the applied voltage to zero with respect to the S.C.E. after connecting the rotating platinum micro-electrode to the polarising unit. Adjust the range of the micro-ammeter. Titrate with the standard 0.005 M iodine solution in the usual manner. 

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. 

The end point of the reaction is conveniently determined electrometrically using the dead-stop end point procedure. If a small e.m.f. is applied across two platinum electrodes immersed in the reaction mixture a current will flow as long as free iodine is present, to remove hydrogen and depolarise the cathode. When the last trace of iodine has reacted the current will decrease to zero or very close to zero. Conversely, the technique may be combined with a direct titration of the sample with the Karl Fischer reagent here the current in the electrode circuit suddenly increases at the first appearance of unused iodine in the solution. 

The Karl Fischer procedure has now been simplified and the accuracy improved by modification to a coulometric method (Chapter 14). In this procedure the sample under test is added to a pyridine-methanol solution containing sulphur dioxide and a soluble iodide. Upon electrolysis, iodine is liberated at the anode and reactions (a) and (b) then follow the end point is detected by a pair of electrodes which function as a biamperometric detection system and indicate the presence of free iodine. Since one mole of iodine reacts with one mole of water it follows that 1 mg of water is equivalent to 10.71 coulombs. 

The platinum electrode was modified with an iodine adlayer of the specified symmetry. ) AC = acetate anion. 

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