Control led- Potential Coulometric Determination of Copper and Uranium W. D. SHULTS und P. F. THOMASON Analyfical Chemistry Division, O a k Ridge Nafional Laborafory, Oal: Ridge, Tenn. $Solutions of copper and uranium sulfate can be analyzed with relative standard deviations of about 0.1% by controlled-potential coulometric titration if both ions are reduced at -0.3 volt and copper only is reoxidized at +0.175 volt. The method is not affected by the copperuranium ratio.
*-SILVER WIRE
SOLID KCI; AgCI
T
controlled-potential coulometric titration of uranyl ion in a citrate electrolyte was reported by Booman, Holbrook, and Rein to give results reliable to 0.1% (2). This method is particularly suited to the analysis of radioactive materials, which led Rush (8) to study the method for the determination of uranium in a sulfate electrolyte. Merritt, Martin, and Bedi used the controlled-potential coulometric deposition of copper on platinum as one test of the performance of their apparatus (7). The reduction of copper at a mercury cathode was used as one test of the performance of the ORNL electronic controlled-potential coulometric titrator ( 5 ) . By careful choice of potentials, Farrar, Thomason, and Kelley titrated copper and uranium successively in homogeneous reactor solutions where the quantities of copper and uranium titrated varied little from 0.3 and 5 mg., respectively (3). Their method is not applicable to samples in which the copper-uranium ratio varies considerably. The method described here was developed to eliminate this difficulty. It is based on the fact that the copper reduction a t a mercury electrode is reversible, whereas the uranium reduction is irreversible in sulfuric acid electrolyte. By making two titrations, the first a reduction in which the copper and the uranium are both reduced, and the second a reoxidation of the then reduced copper, both ions can be determined with excellent precision and accuracy even when the amounts of copper and uranium vary considerably. HE
APPARATUS AND REAGENTS
Electrical Apparatus. The ORNL electronic controlled-potential coulometric titrator was used throughout this study. This instrument is a
492
ANALYTICAL CHEMISTRY
Figure 1 .
Coulometric titration cell
simplification of the one described by Booman ( I ) and its ciwuitry and performance have been rcported (5). It is designed to supply elcctrolysis current to the electrode system while maintaining the potential of the working electrode a t some chosen value with respect to a refertmce electrode. Polarity of the workiilg electrode and direction of current flow can be chosen to run either reducti'm or oxidation titrations. A portion of the current that flows is integrats?d electronically by an integrating operrLtional amplifier such that the charge axumulated on a capacitor is a measure of the coulombs of electricity that h,tve been used; the relationship betweeii readout voltage (charge on condenser) and coulombs is known from calibration procedures. I n practice a titration is made, after proper adjustment of the control potential, by allowing the electrolysis to proceed to completion--until the current has decreased to background, usually 50 ha. The rettdout voltage a t this time is measured by a Rubicon (Rubicon Co.) poteniiometer and is multiplied by the appropriate coulometric factor to obtain the weight of the substance titrated. Mechanical Apparai;tls. The titration vessel used in t'his study was designed by Rush (8) 2nd is diagramed in Figure 1. It is a g iiss vessel about the size of a 30-ml. beaker. Fresh mercury can be addec to and sample wastes removed from the cell through the two-way stopcocl; fitted to the
bottom. A Teflon cap supports a silver-silver chloride-potassium chloride reference electrode, a separated electrode compartment containing I N sulfuric acid, and a nitrogen inlet tube. A port for sample entry is provided. Stirring is provided by a glass disk, 1 cm. in diameter and positioned horizontally a t the mercury-electrolyte interface, with a 4-mm. glass rod sealed a t its center and protruding through the cap so that it can be driven by an 1800 r.p.m. Bodine motor (Bodine Electric Co.). Connection to the mercury is made by a platinum wire sealed into the vessel near the bottom and a coiled platinum wire serves as the separated electrode. Reagents. A standard copper sulfate solution was prepared from spectrographically pure copper metal. The stock uranyl sulfate solution was standardized gravimetrically b y the ammonium diuranate method (9) and volumetrically by ferric sulfate titration (4). Appropriate dilutions were made from these solutions in Normax glassware (Owens-Illinois) to obtain solutions of the desired concentrations. All other reagents were C.P. grade. The nitrogen gas was purified before use by passage through two chromous sulfate scrubbing towers. PROCEDURE
Approximately 6 ml. of clean mercury, 5 ml. of 1N sulfuric acid, and the sample aliquot are placed in the cell. The solution is deaerated 5 minutes with a stream of nitrogen gas. Adjustments are made so that a reduction a t -0.3 volt will be made and the electrolysis proceeds until the current falls to 50 pa., at which time the titration is stopped and the readout voltage (ER) is measured, The integrating capacitor is shorted (zeroed), adjustments are made so that an oxidation a t +0.175 volt will be made, and the second electrolysis proceeds until the current falls to 50 pa. This titration is stopped and the readout voltage (Eo) is again measured. The amount of copper titrated is calculated by multiplying the copper coulometric factor by EO. The amount of uranium titrated is calculated by multiplying the uranium coulometric factor by ER - EO. DISCUSSION AND RESULTS
Figure 2 graphically indicates the relationship between the ratio of per
cent copper oxidized to per cent copper reduced and electrode potential when copper is electrolyzed a t a mercury electrode. Essentially complete (99.9%) reduction occurs a t -0.075 volt or more negative, while essentially complete oxidation occurs at -t-0.150 volt or more positive, silver chloride being the reference electrode in each instance. A. potential of -0.1 volt proved satisfactory when the reduction of copper ions was used as a test of the performance of the ORNL electronic controlled-potential coulometric titrator (6). For the oxidation of copper amalgam, a potential of +0.175 volt was chosen. Aliquots of standard copper sulfate solutions were titrated by coulometric reduction a t a mercury cathode whose potential was controlled a t -0.1 volt (silver-silver chloride). When each reduction was complete, the resulting copper amalgam was oxidized coulometrically, the mercury being an anode whose potential mas controlled a t $0.175 volt (silver-silver chloride). An indication of the precision and accuracy obtainable when copper is determined by coulometric reduction of cupric ions and by coulometric oxidation of copper amalgam is given in Table I. The standard deviation values were calculated using the known values of copper content rather than the mean of the observations., The controlled potential coulometric reduction of UOz++ to uranium(1V) at a mercury cathode has been reported (2, 3, 8 ) . This reduction is generally believed to involve formation of an unstable UOzf ion which disproportionates to UOz++ and uranium(IV), the net result being conversion of UOL++
I
Comparison of Results of Copper Determinations by Coulometric Reduction and Reoxidation (A 1-ml. aliquot of standard copper sulfate solution was titrated in 5 ml. of 1N H8Oc by reduction at -0.1 volt and reoxidation a t +0.175 volt. Reference is saturated AgCl.
Coulometric factor: 13.10 mg. Cu/volt. Underlined digits are doubtful.) Copper, blg. Present Found Std. Dev., Mg. Oxidation Reduction Oxidation Reduction 0.0966 0.0963 0.0971 0.0965 0.0971 0.0963 0.0966 0.0971 0.00054 0.0971 0.00047 0.0966 0.0958 0.0971 0.0966 0.097i 0.0971 0.0965 0.0968 0.0971 0.0966 0.0966 0.0971
I
I
I
0.963 0.963 0.964 0.963 0.965 0.968
4.830 4.830 4,830 4.830 4.830 4.830 4.830
4.832 4.835 4.833
II.
0.963 0.963 0.963 0.963 0.965 0.963
0.963
0. (1028
0.0030
0.3070
0.0073
0.962
4.832 4.835 4.837 4.839 4.835 4.836 4.84O_
4.840
4.836 4.831 4.833
Coulometric Titration of Cu*+ and UOz++ by Total Reduction and Selective Oxidation
(A 1-ml, aliquot of standard solution was titrated in 5 ml. ol' 1N HtSOi by reduction of Cu++ and UOz++ a t -0.3 volt and reoxidation of Cu(Hg) ,it $0.175 volt. Reference is saturated AgCl. Coulometric factors: 49.16 mg. U/volt, 13.10 mg. Cu/volt. Underlined digits are doubtful.) Readout Voltages Reduction, Uranium, Mg. Copper, Mg. UOz++plus Oxidation, Difference, Present Found Present Found UOZ + + cu++ Cu(W 0 0 0.0735 0.0735 0.0000 0.966 0.963 0 0 0.0736 0.0735 0.0001 0.96: 0.963
I
-0.05
0.0 t0.5 r0.1 +0.!5 CONTROL POTENTIAL, E v s . A g AgCl ~
Figure 2. Relationship between logarithm (per cent copper oxidized/per cent copper reduced) and electrode potentia I
0.960 0.966 0.96tj 0.966 0.968 0.968 0.96:
Table
1' I
perimentally by couloinetrically reducing a uranyl sulfate solution at a mercury cathode whose potential was controlled a t -0.3 volt and then attempting to reoxidize electrically the uranous sdfate which had been formed. Only background currents were obtained a t potentials less positive
Table 1.
I
-4.01
-0,4
to uranium(1V) when the electrolysis is carried to completion (6). Extremely positive potentials are required to electrically oxidize uranium(1V) to U02++ and this is usually attributed to the difficulty of forming UOz+ from uranium(1V). The irreversible behavior of uranium was verified ex-
0.1170 0.1170
0.0000 0.0000
0.1170 0.1170
0 0
0 0
5.750 5.750
5.752 5.752
0.1904 0.1904 0.1904
0.0735 0,0734 0.0735
0.1169 0.1170 0.1169
0.965 0.96s 0.963
0.963
0.962 0.962
5.750 5.750 5.75Q
5.741 5.752 5.741
0,1540 0.1542 0.1543
0.0368 0.0370 0.0370
0.1172 0.1170
0.483 0.483 0.483
0.482 0.485 0.485
5.750 5 * 750 5 * 75Q
5.762 5.752 5,766
0,1244 0.1245 0.1244
0.0073 0.0073 0.0071
0.1171 0.1172 0.1173
0.092
0.096 0.09s 0.093
5.750 5.750
5.76Q
5.752 5.762 5.765
0.3080 0.3076 0.3078
0.0735 0.0735 0.0736
0,2345 0.2341 0.2342
0.965 0.966 0.966
0.96i]:
0.962i
11.500 11.5012 11.500
11.528 11.508 11.513
0.1030 0.1030 0.1031
0.0737 0.0735 0.0736
0.0293 0.0295 0.0295
0.96s 0.966
0.965
0.96jj 0.96!! 0.96.L
1.438 1.438 1.438
1.440
0.1173
0.092.
0.092
0.96;;
1,450 1.450
VOL. 31, NO. 4, APRIL 1959
493
than +0.200 volt (silversilver chloride); more positive potentials resulted in dissolution of mercury. Standard solutions containing varying amounts of copper and uranium mere analyzed. Two electrolyses were required for each aliquot: a coulometric reduction with the mercury cathode potential controlled a t -0.3 volt and a coulometric oxidation with the mercury anode potential controlled a t +0.175 volt. The number of coulombs of electricity consumed during the oxidation titration is a measure of the quantity of copper titrated, whereas the number of coulombs consumed during the reduction is a measure of the total amount, copper and uranium, titrated. The uranium was then calculated by difference. Table I1 gives the results of these titrations. The data of Tables I and I1 indicate
that good prc oision and accuracy are obtained by t:?is method and the data in Table I1 indicate that the method is unaffected by varying quantities of copper and uranium. There was no investigation of the effect of varims fission products or ionic impurities because this was nude by Farrar, Thomason, and Icelley (5) in a study of the controlled-potential uranium titraiion. The method described has be’tn in use for some time without incider t,
LlTI I!ATURE CITED
(1) Booman, G. L., ANAL.CHEII. 29, 213
(1957). (2) Booman, G. L., Holbrook, W. B., Rem, J. E., Zbfd., 29, 219 (1957). (3) Farrar, L. G., Thomason, P. F., Kelley, M. T., Ibid., 30, 1511 (1958).
(4) Ginocchio, B. J., “Uranium, Auto-
matic Potentiometric Ferric Sulfate Method,” Method No. 1 219224 (2-2458), ORNL Master Analytical Manual, TID-7015(Section 1). (5) Kelley, M. T., Jones, H. C., Fisher, D. J., ANAL.CHEM.31, 488 (1959). (6) Kolthoff, I. M., Lingane, J. J., irPolarography,J’ 2nd ed., p. 462, Interscience, New York, 1952. (7) Merritt, L. L., Jr., Martin, a. L., Jr., Bedi, R. D., ANAL. CHEII. 30, 457 (1958). (8) Rush, R. M., “Determination of Uranium by Controlled Potential Coulometry,” ORNL preliminary report, RMR-2 (March 5, 1957). (9) Shults, W. D. 11, “Uranium in HRE Solutions, Ammonium Diuranate Gravimetric Method,” Method No. 9 082205 (1-22-53), ORNL Master Analytical Manual, TID-7015(Section 9). RECEIVED for review September 10, 1958. Accepted November 28, 1958. Southeastern Regional Meeting ACS, Gainesville, Fla., December 1958.
Determination of Ammonium and Potassium Ions in Mixtures of Alkali Metalis Use o f M erc ury (I I) -( Eth y I e ne d i ni ‘1riIo )tetra a c et i c Ac i d FAWZY S. SADEK and CHARLES N. REILLEY Deparfmenf of Chemistry, Universify o f Norfh Carolina, Chu,t>elHill, N. C. Two rapid volumetric methods for the determination of ammonium and potassium ions in mixtures of alkali metals are proposed. For ammonium ion, the acid liberated by the reaction between the mercury(l1) complex of (ethylenedinitri1o)tetraacetic acid and ammonium ion is titrated; the method is selective for ammonium ion in the presence of free ammonia and other alkali metal ions. It is virtually free of anion interference, except phosphate, and gives accurate results in the semimicro and micro ranges. A second procedure, applicable to ammonium and potassium ions, is based on precipitation of their tetraphenylborate salts, dissolution in N,N-dimethylformamide, treatment with excess mercury(11)-EDTA, and back-titration of the liberated EDTA with standard magnesium solution. This procedure can also be applied to the analysis of other compounds which form insoluble tetraphenylbora tes.
W
HEN sodium, potassium, and ammonium ions are present together in a sample, laborious gravimetric separation and determination are often employed for no rapid and accurate
494
ANALYTICAL CHEMISTRY
volumetric methods are known. The present paper shows that, by a combination of the r-rethods developed, this problem can 1:e largely overcome. The methods arc satisfactory on microand semimicro iscale even when the alkali metal COI tent of the sample is very small. Free ammonia in solution in the absence of other basic constituents is usually best determined by direct titration with standard acid using an appropriate indicator Ammonium ion is usually determin od gravimetrically by precipitation wihh chloroplatinic acid or volumetricall~ by Kjeldahl or Conway procedures. The volumetric estimation of amnionium ion using the formaldehyde re5 ction discovered by Plochl (16) is ra:tid but side reactions also occur, as shown in the extensive studies made since its discovery. Baur et al. and Williamt; et al. (3, 25) showed that an equivalerit of acid is usually freed in the intera:tion of an ammonium salt with formalc chyde to form hexamethylenetetramine according to the following over-all reaction: 6 HCHO
+
4 NH.: + (CHzhNl -I- 6 HzO
+ 4 H+
(1)
The acid liberated forms the basis for the analytical estimation of the ammonium ion. Boyd and Winkler (?‘) conducted kinetic studies and concluded that the reaction was actually more complex than indicated by Equation I, and that products other than hexamethylenetetramine are formed. Among these products, various workers (1, 7, 11, 16) have identified mono-, di-, and trimethylamines, and condensation products of formaldehyde with monoand dimethylamines. Kolthoff (16) stated that the formation of these byproducts was caused by the presence of excess alkali or formaldehyde and that the titration error increased with contact time of the reactants. The proposed estimation, based on the reaction of ammonium ion with mercury(11)-EDTA, is free of these side reactions and offers certain advantages.
DETERMINATION OF AMMONIUM ION USING MERCURY(I1) CHELONATES
A number of mercury(I1) chelonates behave as Lewis acids towards ammonia HgY --2
+ NHa e HgYNH, *
(2)
