V O L U M E 24, N O . 1, J A N U A R Y 1 9 5 2 tubes t o sag and requires frequent replacement of the furnace insulation. T h e method, nevertheless, is cheaper and more reliable than those previously used for chromium alloys containing titanium and zirconium, which hold sulfur very tenaciously. Sulfur may be precipitated from high-chrome solutions as the sulfate and dried, the impure precipitate placed in t h e furnace, and the sulfur burned off. Ore tailings containing as much as 8.5%sulfur and 2% arsenic have been run and results obtained in excellent agreement with those obtained by the gravimetric method. With ferrotungsten and iron ores, an accelerator has not been found necessary; for ferrochrome, 30-mesh tin is required. Open hearth slags are run successfully using a lom--sulfur steel standard as a bedding material. The exothermic reaction of the steel with oxygen appears to improve recovery of the sulfur as sulfur dioxide. The consensus of the panel and audience was that an empirical factor based on a sulfur recovery in the order of 80 to 90% was necessary a t present in analyzing steels routinely on a 24-hour-aday basis. T h e experience of some persons was that this factor varied from day to day; others do not encounter this day-to-day variation but find a variation from one sulfur range to anotherLe., the factor for an 0.030 to 0.040% sulfur range is different from that for a 0 10 to 0.15% sulfur range. Apparently, there remain details of the procedure not yet well enough understood t o be standardized. It has not been established with certainty, for example, whether recovery of all the sulfur as sulfur dioxide depends on use of high temperatures or
205 on prevention of condensation of iron oxide in the cooler portion of the tube. Methods that will operate regardless of the form i n which the sulfur is evolved seemingly have not been completely investigated. Exploration of these and other aspects of t h e procedure remains a challenge t o the analytical chemist. I n t h e meantime, the combustion method is being used empirically but successfully in many different laboratories and by different operators. LITERATURE CITED
Allen, Alfred, J . Iron Steel Inst., 7, 480-90 (1879). I b i d . , 8, 181-91 (1880). Baker, Graham, and Stuart, TVm., J . Chem. SOC.(London), 17, 390-98 (1864). Beeghly, H.F.,ISD. ENG.CHEM.,ANAL.ED..14, 137-40 (1942). Committee on Analytical Reagents, AMERICAN CHEMICAL SOCIETY, “Reagent Chemicals,” Washington, D. C., AMERICAN CHEMICAL SOCIETY, 1951. Gotta, A., and Seehof, H. S., 2. anal. Chem., 1924, 216-26 (1942). Jordan, L., and Eckman, J. R., Natl. Bur. Standards, Sci. Paper 563 (1927). Kempf, H.,and Abresch, K., Arch. Eisenhiitfenw., 14, 255-9 (1940). Parnas. J. K., 2. anal. Chem.. 114, 261-75 (1938). Parnas, J. K., and Wagner, R., Biochem. Z., 125, 253-6 (1921). Pohl, H., Melallwirtschuft, 23, 347-9 (1949). Seuthe, A., Stahl n . Eisen, 52, 445-6 (1932). lVoodx\-ard, T. S.,and Wolthorn, H. J., “Sampling and Analysis of Carbon and Alloy Steels,” Appendix I, Kew York, Reinhold Publishing Corp., 1938. RECEIVED November 28, 1951.
Coulometric Titration of Manganese in the Submicrogram Range K, DONALD COOKE’, C. N. REILLEY, AND N. HOWELL FURMAN Princeton Cniversity, Princeton, AY.J . Recent w-orlr on a sensitive electrometric end-point procedure indicated that the range of titrimetric analysis could be considerably extended. The use of this end point, in conjunction with coulometric generation of ferrous ion, allowed permanganate solutions to be titrated in the parts per trillion range. Solutions of 0.001 to 0.0003 microgram of manganese per milliliter could be titrated with an accuracy of about 5’3,. Kith more concentrated solutions the
C
OULOLIETRIC methods have offered the possibility of
estending titrimetric procedures t o a concentration range unattainable by conventional titration. Currents of the order of 60 electrons a second can be measured using modern electrometer tubes ( 3 ) . Hence, i t seems that there is no practical minimum t o the magnitude of currents t h a t can be used t o geneiate reagents coulometrically. The liniitations of such procedures lie in the detection of the end point and possibly even the chemistry of the systems involved. -4recently developed end-point procedure has been applied t o the titration of ferrous ion with electrically generated ceric sulfate (21, and the titration of vanadium in uranium u i t h electrically generated ferrous sulfate (4). I n the present v-ork, a further investigation of the end-point method has been undertaken, and the method extended into the submicrogram range. A potentiometric-type procedure has been described which uses a n amperometric indication of the potential of the solution ( 2 ) . 1
Present address, Department of Chemistry, Cornell University, Ithaoa,
h’. Y.
error was considerably less. Coulometric titrations using generation currents of a few microamperes have been found feasible. To obtain a satisfactory end point in titrations at such dilutions, diffusion currents at a large stationary electrode were measured to 30 micromicroamperes. The use of these techniques has made possible the titration of solutions at greater dilutions than w-as possible by tcchniques previously in use.
-1reference potential is chosen near the point of steepest slope on the potentiometric titration curve in a fashion similar t o that of Muller ( 6 ) . The solution is then titrated t o the particular reference potential chosen. An ordinary amperometric circuit using a large stationary platinum electrode shows when the reference potential is reached during the titration. The desired reference potential is impressed upon the indicator electrode by the voltage supply shown in Figure 1. When no current f l o w in the indicator circuit, i t must be a s u m e d that the potential impressed upon the electrode and t,he potential of the solution are the same. At any other solution potential, a diffusion current will be noted in the indicator circuit. Thus, the titrat,ion is carried out unt,i) there is zero diffusion current, which indicates that the potential of the solution has reached the chosen reference potential. By pretitrating the reagents used t o the reference potential before addition of the unknown, much more accurate results could be obtained, and the method could be applied t o solutions of much greater dilution. This also made the exact choice of a reference potential less important ( 2 ) . Another advantage of this proce-
ANALYTICAL CHEMISTRY
206 dure was that the effects of impurities in the reagents were minimized. In fact, the reagents could be made up in t a p water and the correct results obtained in the submicrogram range. This is advantageous because it eliminates many of the tedious precautions necessary in conventional trace analysis. Using a sensitive galvanometer in the indicator circuit, very large deflections could be obtained for relatively small amounts of material. For example, 1 microgram of manganese as permanganate ion gave a galvanometer deflection equivalent to 700 meters when an indicator electrode of 3 sq. em. was used. This gave a sensitivity of about 10-6 microgram, or 1 micromicrogram.
justed to a potential of 0 . i 5 volt (by the generation of ferrous), a rapid drift of the indicating galvanometer was noted. This drift was in the direction of an increase in voltage or oxidizing power, and was believed caused by the air oxidation of the ferrous ions in solution. When the solution was adjusted to a higher voltage, the drift disappeared, because the solution became more difficult to oxidize. By choosing a voltage above the paint a t which the solution would react with air, it was possible to carry out the titrations without excluding air. Pure ferrous sulfate solutions in air, a t these dilutions, would probably be extremely unstable. In view of the above facts, a potential of about 0.87 volt was chosen. The buffering of the permanganate with manganous ion was found to be advantageous. Steadier diffusion currents and more accurate results were obtained by the addition of a large excess of divalent manganese. As a large escess of both ferric ion and manganous ion is present, the addition of the permanganate sample to this poised solution changes the poten- 46- 19sv. tial only slightly. The addition of a sample of 0.01 microgram changed the potential of the solution about lo-' volt. Honever, as a very senTIMER I S - 4 4 MCO. COARSE sitive galvanometer was used, this slight change I I in potential was amplified to a satisfactory deIlOV. flection. Because of the small magnitude of the INDICATOR C I R C U I T GENERATOR CIRCUIT voltages measured, the voltage impressed upon the indicator electrode should be kept constant Figure 1. Wiring Diagram to within a few microvolts during the titration period. As comparatively high diffusion currents were obtained using When the solution of ferric ion is adjusted to 0.87 volt, a rather the large stationary electrode, the amount of material electrolyzed large amount of ferrous ion is present in solution compared to the a t the indicator electrode was at times enough t o cause low resize of the samples of permanganate added. When the permansults. This effect was decreased by using smaller electrodes ganate is added it is believed that reaction with the ferrous ion with the larger samples. present immediately takes place. The titration then consists of In macrotitrations of potassium permanganate with ferrous merely adjusting the F e + + + / F e + + ratio to its original value. sulfate, the potential of the solution a t the end point is about 1.04 This possibility might be advantageous in the titration of solutions which are unstable in the titration medium, or require an volts versus the hydrogen electrode. I n titrations in the microgram range v, hen this voltage was used as an end-point reference, escess of reagent for rapid reaction. low results were obtained. According t o the Nernst equation Thermostating the titration cell was helpful in the submicrothis nould be the case if the correct end point voltage was lower gram range. It is believed that the difficulty necessitating than 1.04 volts. It was found that any reference potential below temperature control is not caused by the dependence of the diffusion coefficients on temperature, but the variance of the poten0.89 volt gave the correct titration value if the solution was pretitrated to the chosen value before addition of the unknoiin as tial of the solution with temperature according to the Nernst previously outlined. equation. -4s potential changes of a few microvolts are probably Although any reference voltage below 0.89 volt may be used, other factors influence the choice of this potential. The sensitivity of the end point varies considerably with the voltage applied. This is to be expected, as the slope of the potentiometric curve is different a t different voltages. On the steep portion of the potentiometric curve, a greater voltage change and hence a greater sensitivity are realized for the addition of a quantity of permanaganate. On the flat portions of the curve where the solution is ne11 poised, only small voltage changes can be realized. This change in sensitivity with potential is shown in Table I, in which the galvanometer deflection per microgram of manganese is shorn for the 0 1 sq em electrode.
m
h-+ -
25
Table I. Reference Potential 0.88 0.86
0.80 0.75
Sensitivity Change Sensitivity. h l i n . / r 11,500 4,200 3,300 930
One other factor influencing the choice of a reference potential is the stability of the solutions when adjusted to the chosen voltage. For example, when the solution of ferric sulfate n-as ad-
s
0
I
2
3
SECONDS
Figure 2.
Titration Curve
4
V O L U M E 2 4 , NO. 1, J A N U A R Y 1 9 5 2 Table 11.
207
Coulemetric Titration of Manganese
R
M n Taken, y
M n Found, y
1.201 1.091 1.328 1.227 1.196 0.712 0.501 0.516 0.543 0.484
1,189 1.096 1.310 1.229 1.216 0.721 0.404 0.515 0.536 0 492
1.0 0.5 1.4 0.2 1.5 1.3 1.3 0.2 1.1 1.7
0.197 0.171 0.166 0.146 0.199 0.215 0.173 0.192 0.153
0.187 0.171 0.157 0.158 0.213 0.176 0.205 0.157
5.0 0.0 5 .6 6.4 2.5 1.0 1.8 6.9 2.6
0.0786 0.0419 0.0384 0,0483 0.0511 0,0369 0,0009 0.0076 0.0033
0.0779 0.0429 0,0380 0.0488 0.0505 0.0373 0,0092 0.0084 0.0030
0.9 2.4 1.1 1.0 1 2 1.11 7.1 9.2 9.1
0.204
Error.
Approximate Generating Current, fia. 50
monium sulfate plug, and containing 0.1 N sulfuric acid, A small platinum wire was used as the anode. The indicator circuit consisted of a platinum-iridium foil electrode, and a saturated lead amalgam-lead sulfate reference electrode, as shown in Figure 1 Three indicator electrodes of different areas (3. 0.8, and 0.1 sq. cm.2) were used, so that a diffusion current of desirable magnitude could be obtained The solution was stirred by means of a magnetic stirrer. ELECTRIC4L CIRCUIT
10
5
significant, this factor could cause gross errors unless the temperature \vas controlled. However, only temperature changes during the course of the titration would cause difficulty. =it the beginning of this work, a mercurous sulfate half-cell was used as a reference electrode. This electrode was found to have disadvantages in submicrogram titrations. Because small currents were continually drawn through the cell, rather large fluctuations would occur in the e.m.f. of the cell. The second disadvantage waa the comparatively high solubility of mercurous sulfate. Slow diffusion of mercurous ions through the agar plug caused difficulty in the permanganate titration. To avoid these difficulties, a saturated lead amalgam-lead sulfate electrode in 0.5 S sulfuric acid was used. This half-cell maintained a rather constant e.m.f. (-0.2; volt) and no osidizible ions were introduced into the titration vessel. For this reason, the lead amalgam-lead sulfate electrode is recommended for microelectrometric titration. Care should be taken in the preparation of this electrode; especially, the platinum mire lead under the mercury should be dry ( 1 ) . In order to obtain known samples of 1 microgram or less, solutions down to 10-4 I V potassium permanganate were prepared daily by diluting a stock 0.1 N solution. For samples less than 0.1 microgram, quantities as small as 1 mg. of N solution were measured into the solution to be titrated by means of a weight buret constructed from a 1-ml. hypodermic syringe. More dilute standard solutions were not prepared becauseof two possible sources of error: the apparent presence of a small amount ( 10-o or lo-‘ AV)of reducing agent in the distilled water, and the fact that by storing the solution a t higher concentrations, the possibility of significant adsorption on the surface of the glass vessels was minimized. APPARATUS
The titration cell was similar to one previously described ( 2 ) except it was equipped with a a a t e r jacket, to maintain the temperature to wit,hin 0.1’ C. The capacity of the cell was about 7 ml. The generator cathode was of platinum-iridium foil, 1 sq. cm. in area. The anode compartment of the generator circuit consisted of glass tubing 3 mm. in diameter, isolated by a 4% agar-2% am-
Because the generating currents used in this work were less than 50 microamperes, a constant voltage supply was used as a constant current source (6). Five small-capacity 45-volt batteries were used as the generating current source, as shown in Figure 1. The current was measured by taking the ZR drop across a standard resistance with a potentiometer. The time was measured with a synchronous motor clock (Standard Time Co.. S-60). The galvanometer used was a Leeds & Xorthmp HS type XTith a sensitivity of 10-6 microampere per millimeter, equipped with an Ayrton shunt (ratios 1 to 0.0001). The resistance of the indicator circuit should be as lo~vas possible to increase the sensitivity. Care must be taken that all the connections are sound. REAGENTS
The ferrous ion was gent:rated from it 0.01 zV solution of ferric ammonium sulfate. A 0.01 JI solution of manganous sulfate was used to decrease the potential of the permanganate system. The potassium permanganate solution was standardized against primary standard ethylenediamine ferrous sulfate and was found to be 0.1231 *V. PROCEDURE
About 5 ml. of 0.01 S ferric ammonium sulfate and 0.2 nil. of 0.01 iV manganous sulfate were added to the titration cell. The potential of this solution is usually greater than the reference voltage, 0.87 volt. The ferric ion solution was then brought to the reference point by generating the required amount of ferrous ion. Aft.er the potential was adjusted, the unknown solution of permanganate ion was added. The timer was reset and ferrous ion generated in appropriate intervals until the galvanometer returned to zero. The diffusion current reached equilibrium in about a minute in the vicinity of the end point. Rather than adjustment of the current exactly to zero, some point near zero current may be used as the reference point. The curves were not plotted, but the return to the original reference potential \vas taken as the end point. The titration curve for a sample of 0.003 microgram is shown in Figure 2. If the samples were titrated just to the end point, additional samples could be added to the same solutions. RESULTS
The results of this work are shown in Table 11. ACKNOWLEDG.MENT
The authors gratefully acknowledge the aid, in the form of a postdoctoral fellowship to one of them (W.D.C.), furnished by the Eugene Higgins Trust Fund. LITERATURE CITED (1) Clark, W..M., “The Determination of Hydrogen Ions,” 3rd ed., p. 303, Baltimore, Williams & Wilkins Co., 1928. (2) Cooke, U’.D., Reilley, C. S . ,and Furman, N. H., ANAL.CHEhf., 23, 1662 (1951).
(3) Elmore, W.C., and Sands, Mathew, “Electronics-Experimental Techniques,” p. 181, Kew York, McGraw-Hill Book Co., 1949. (4) Furman, S . H.. Reilley, C. K.,and Cooke, W. D., ANAL.CHEM., 23, 1665 (1961). (5) Muller, Erich, “Electrometrische hfassanalyse,” 6 Aufl., p. 90, Dresden, T. Steinkopff, 1942. (6) Sease, J. W., Niemann, C., and Swift, E. H., ANAL.CHEW,19, 197 (1947). RECEIVED May 10, 1951.
