ANALYTICAL CHEMISTRY
1272 sensitivity of the reaction, although the rate of formation of the precipitate was decreased. CONCLUSIONS
A solution of 2,5-dih~-droxy-1,4-benzoquinone in acidified nlcoho1 produces a pinkish precipitate 4 ith scandium n hich is obtained, although slowly, in the presence of large concentrations of sodium and ammonium salts. The interference of ammonium tartrate may be eliminated by the addition of exceSs concentrated hydrochloric acid. Oxalate and phosphate interfere; fluoride appears to increase the sensitivity of the test. The test is more specific when run in this manner than 1% hen an aqueous solution of the ammonium salt of the quinone i q used, as described in the literature (3). The scandium ion. in fact, does not readily form a precipitate with the ammonium salt. JVhile many ions interfere, a t least half of these can readily lie removed by precipitation as sulfides in acid solution. It is of interest to note that in the strongly acid solutions employed only highly charged ions or hydrated ions which are acidic in the Lewis sense, give a positive test. Of the lanthanide rare earths, the cerium subgroup ions apparently do not interfere, while the yttrium subgroup ions are indistinguishabl? from scandium. Thorium cannot readily be distinguished from si~andium.
ACKNOWLEDGRIENT
The authors wish to thank the Edwal Laboratories for supplying a sample of the quinone emplo>-ed. The work described rvas supported in part by a Frederic Gardner Cottrell grant from the Research Corporation, whose finnncial assistance is gratefully acknowledged. LITERATURE CITED
(1) Feiel. F.. ”Snecific and Snecial Reaneiits.” Sew I-ork. Elsevier , j
,
PEilishiiigACo.,1940.
~
( 2 ) Flagg, J. F., “Organic Reagents Used in Gravinietric arid T-olu-
metric hiialysis,” Sew Tork, Interscience Publishers, Inc.,
1948. (3) Frank, R. L., Clark, G. K., and Coker, J. S., .J. Ant. Chent. Soc., 72, 1827 (1950). (4) Fresenius, R., and Jander, G., “Handbucli der analytischen
Chemie,” Part 111, p. 734, Berlin, Springe],T-erlag, 1940-42. (5) Kusnetsov, V. I., J . Geii. Cliern. (CT,S.5’.R.),14, 89T (1944). (6) Pokras, L., and Bernnys, P. AI., J . Am. C h e m . Soc., 73, 7 (1951). (7) Sandell, E. B., “Colorimetric Determination of Traces of Xetals,” pp, 392-3, Xew York, Interscience Publishers, Inc., 1944. (8) Wenger, P. E., and Duckert, R., ecl., “Reagents for Qualitative Inorganic Analysis.” pp. 163-4, S e w T o r k , Elsevier Publishing Co., 1948. RECEIVED August 13. 1 9 2 .
.\ccepted May 7, 1953.
lodometric Determination of Hypophosphorous and Phosphorous Acids RICHARD T. JONES AYI) ERR-EST €1. SWIFT Gates and Crellin Laboratories of Cherriistry, California Institute of Technology, Pasadena, Calif. an investigation of the oxidation nit~c~h:tnism of liyD pophosphorous acid, quantitative procedures were needed for the determination of hypophosphorous and phosphorous URISG
acids, alone or in mixtures. Such methods would be valuable because phosphorous acid is almost a h a y s present in hypophosphorous acid preparations ( 1 ) and is frequently an intermediate product in the oxidation of hypophosphorouP acid.
hn cspcriniental study of thew methods has rwulted in niodification of the Wolf and Jung method viith promlures 11ascd upoil the folloiving assumptions: Phosphorous acid can be oxidized quantitatively hy excex:: iodine in an essentially neutral solution in 1 hour; oxidation of hypophosphorous acid under the same ronditions is negligible. Hypophosphorous acid can he oxidized quantitatively by excess iodine in 1 to 2 S hydrochloric acid in 3 hour;?.
PREVIOUS METHODS
Acid-Base Titrations. (K.4 = 1.0 X lo-*) can base, especially when a presence of phosphorous 7 X lo-’) such titrations
Pure hypophosphorous acid solutioris be titrat’ed accurately viith standard pH meter is used. However, in t h r acid ( K A = 1.6 X lo-? and K,i = are much less reliable.
Oxidation Methods. Procedures involving the oxidation of hypophosphorous acid to phosphoric acid with excess st,andard permanganate in hot solutions have been proposed. However. t>he instability of permanganate ucder such conditions led Kolthoff (3) to propose oxidation in a cold sulfuric acid solution for 24 hours, with blank experiments being used to compensatcx for permanganate decomposition. Kosaegi ( 4 ) has indicat?tl that errors exceeding 1 t,o 4% are possible in this method. Rupp and Fink ( 5 ) appear to have been the first to propose the use of iodine as a n oxidizing agent. However, Wolf and Jung ( 6 ) have pointed out inconsistencies in their results and have outlined methods for determining hypophosphorous and phosphorous acids separately b u t not in the presence of each other. Phosphorous acid was oxidized with excess iodine in a bicarbonate-carbonic acid buffer; after about 1 hour the excess iodine was back-titrat,ed with standard arsenous acid. Hypophosphorous acid was oxidized ( t o phosphorous acid) with excess iodine in an approximately 0.5 Y sulfuric arid solution for. 10 hours; the solution was then neutralized with excess sodium bicarbonate and treated by the procedure for phosphorous acid. Results for phosphorous acid were excellent. However, those for hypophosphorous acid were 0.4yo loner than the corresponding gravimetric results. Kamecki ( 2 ) proposed the oxidation of hypophosphorous a d to phosphorous acid by iodine in 1 to 4 AT sulfuric acid a r i d claimed t h a t the oxidation of phosphorous acid to phosphoric is slow under such conditions. However, experiments described below indicate an appreciable oxidation of phosphorous acid under these conditions,
EXPERIIIEN1.A L
Apparatus and Reagents. Volume nicasurement,s were used exclusively, t,he apparatus was especially calibrated. arid appropriate temperature corrections ivere made. Hypophosphorous acid and hypophosphitc solutions for preliminary experiments r e r e prepared from either t,echnical grade 5OTC hypophosphorous acid ~ o l u t i o nor from reagent grade sodium hypophosphite. Purified hypophosphorous arid \vas prepared by a special procedure ( 1 ) and was used for the final itonfirmatory analyses. All other rhemicals werp reagrnt or analytical grade. Standard thiosulfate solution3 and iodine solutions were prepared and standardized by conventional methods. Phosphate buffers were prepared from monosodium phosphate monohydrate and sodium hydroxide. A l l solutions were tested for oxidizable impurities. Oxidation of Iodide by Oxygen. The possibility of oxygen error was checked by allon-ing 1 to 2 S hydrochloric acid solutions containing known volumes of standard 0.1 S iodine solution to stand in the dark a t room temperature in glass stoppered flasks for 3 t o 5 hours. Some of these had been flushed with carbon dioxide. S e r t , the acidity was either left the same or adjusted to a p H of 4 to 5 and the iodine titrated with standard thiosulfate solution. The solutions which had not been treated with rarbon dioxide Irere generally high more; the carbon diosidr-treated solution i’rror, Thereafter, carbon dioxitic vias u chloric acid solutions. Similar test!: on iodine in a neutral 1)ho::ph:tte buffered solution, pH 7 . 3 , showed no error when the polutions were adjusted t o a pH of 4 to 5 Jvith acetic acid before the thiosulfate titration. Oxidation of Phosphorous Acid by Iodine. ACID SOLETIOKS. T h ? extent t o which iodine oxidizes phosphorous acid (in a 1 to 2 .V hvdrochloric acid solution) was investigated in order to establish the feasibility of a determination based upon the oxidation of hypophosphorous to phosphorous acid- under these conditions. First 1.5 -T hydrochloric acid solutions, containing known
1273
V O L U M E 2 5 , NO. 8, A U G U S T 1 9 5 3 Table I.
Oxidation of Phosphorous dcid hy lodine ill 1.5 !V Hydrochloric .icid Solutions
118P02 T a k e n Xes.
(for 2e Change) 2.18 2 18 2 18 2 18 2 00 2 . 00
Period of Standing.
ti 0 1 1 I'i 3 3.3 4 18 3 4
Iodine Reduced. 1leq. 2 322 2.338 2.348 2.658
lI3POa Oxidizd 1Ieq 0.14 0.16 0 17
2.223 2,243
0.23 0.33
0.48
sholv that there was no appreciable iodine oxidation of hypophosphorous acid in solutions of pH above 5 even after 1 hour. Solutions of 1.,5 S hydrochlorir acid containing hypophorphorous acid and known quantities of iodine were prepared and allowed to stand in the dark for various periods of time. The solutions were t'reated with carbon dioxide. After standing, the hydrochloric acid was neutralized with 6 S sodium hydroxide, and the solution was buffered at a p H of about 7.3. An hour was allowed for complete oxidation of all phosphorous acid, and the excess iodine was titrated with thiosulfate after the p H had been adjusted t o about 5 with acetic acid. The total equivalents of oxidized material were the same for a11 solutions standing from 3 to 5 hours during the hypophosphorous acid oxidation; those standing for about 2 hours were lower by a few per cent. Therefore, a t,ime interval of 3.5 hours appears adequate for this oxidation. These experiments resulted in two procedures based. on the foilo\ving reactions:
+ HgO + ICH3PO4 + 2H- + 31- (neutral solution) HIP02 + Ha0 + 13- = HaPo? + 2H+ + 31- (1.5 F HCI)
HIP02
(1)
(2)
Tahle 111. Confirmatory Analyses of HypophosphorousPhosphorous Acid Solutions Acid Sormality _____ Calcd. f r o m formalities found by
Formality Found by Procedures I and I 1 H3P03 HaPo2
Found by Titration pH
0.00006 0.00006 0.00006
0.02361 0.02360
0.02373 0 02373 0.02373
7
0.02370 0.02369
0.00008 0 00008 0 00008 0 00008
0.05060 0.05058 0,05060 0,05064
0 05021b
7
0.05072 0.05070 0.05072 0,03076
0 . 00009 0 00009 0 . 00009 0 0000!1
0.1731 0.1730 0.1730 0.1728
0.1730 0.1730
7
0.1732 0.1731 0 1731 0.1729
0 2134 0.3133 0 2135 0.2133 0 2134 0 2134 0.2135
0.0002 0,0003 0.0001
...
... ...
, . .
...
... ...
kI 81'0 2 Taken, 111
Jf
1.10
1.10 1.10 1.10
Oxidation of Hypophosphorous Acid by Iodine i n Neutral Solutions H3P03
Taken, IIlM
0.080 0.080 0.08 0,08:
Period of Standing, Hours
Hap03 Found.
1
0.078 0.078 0.078 0 . 08.5
.
1 3
2 18
.I4
111
Cnder the conditions present there iy no appreciable iodine oxidation of hypophosphorous acid during a 1 to 2 hour standing period, and during 1 7 to 18 hours only 0.35% of the hypophosphorous acid is oxidized t o phosphoric acid. Therefore, less than 0.1% of a n y hypophosphorous acid could be oxidized to phosphoric acid during a 4-hour standing period. A large relative error would result only if a small quantity of phosphorous acid were determined in the presence of a large quantity of hypophosphorous acid. .ACID SOLUTIONS.Experiments were made to establish the (>oriditionsunder which hypophosphorous acid could be oxidized quantitatively in 1 t o 2 Y hydrochloric acid solutions and t o
...
...
'
o,iiio 0 2133c
... ... ...
4
,..
... ... ... ... ...
,..
... ... ...
0.00016 0.1303 0 1363 7 0.1365 0 . 136.5 0.00016 0.1363 0.1363 F HaPo*. .4t this pH. total acid calculated as 1.5 X F Value from this single titration uncertain, because of erratic behavior of P H meter. e Titration of phosphorous acid solution (no added hypophosphorous acid),
+
'
Table 11.
Procedures I and 11"
Procedure I, based on Reaction 1, will determine the phosphorous acid present (even in the presence of hypophosphorous acid) if the pH is adjusted to approximately 7.3 and 1 hour is provided. Procedure I1 first utilizes Reaction 2 (which requires a solution about 1.5 F in hydrochloric acid and a standing time of 3 hours), fallowed b y Reaction 1 under the conditions of Procedure I. I n Procedure 11, 2 equivalents of iodine are reduced per mole of phosphorous acid present', plus 4 equivalents per mole of hypophorphorous acid. The moles of hypophosphorous acid present are c d x d a t e d by subtracting the equivalents of iodine reduced in Procedure I from those reduced in Procedure I1 and dividing the rc,wlt l)y 4. PROCEDURE I
Take 25 ml. of phosphate buffer solution (prepared as described above) in a 250-ml. iodine flask. Pipet into the flask 25 ml. of the hypophosphorous-phosphorous acid solution t o be analyzed. Then add exactly 25 or 50 ml. of 0.1 N standard iodine solution (depending upon the quantity of phosphorous acid present). Stopper the flask, gently sivir], and place in the dark for 1 hour. Finally, add 5 ml. of 6 S acetic acid and immediately titrate
ANALYTICAL CHEMISTRY
1274 the excess iodine to a starch-iodine end point with 0.1 Ar standard thiosulfatc solution. PROCEDURE I1
Take 30 ml. of G S hydrochloric acid in a 250-ml. iodine flask. Pipet into the flask 25 ml. of the hypophosphorous-phosphorous acid mixture to be analyzed. Flush the solution and flask thoroughly with carbon dioxide. Pipet in 50 ml. of 0.1 A' standard iodine, wash down the sides of the flask, stopper, and gently swirl. Then, place the flask in the dark for 3.5 hours. Add G -V sodium hydroxide equivalent to the hydrochloric acid present, swirl the flask gently, and add 25 ml. of the phosphate buffer. Again waih the sides with water, stopper the flask, gently swirl, and place in the dark for 1 hour. Then, add 5 ml. of 6 S acetic acid and titrate the ewess iodine immediately with 0.1 S standard thiosulfate solution t o a starch-iodine end point. Table I11 presents data obtained from confirmatory analyses made by these procedures and from neutralization titrations in which the end points were obtained by the use of a glass electrode p H meter. For comparison purposes, there is shown the sum of 1.5 times the formality of the phosphorous acid plus the formality of the hypophosphorous acid for those titrations carried to p H i
when less than 1 mole % of phosphorous acid was present. The (lata indicate that mixtures of thwe two acids can be analyzed with a precision within 0.1 to 0.2% for mncrntrations of 0.02 F or greater. Close agreement with the :widinictiic*determinationx indicates a comparable accuracy. ACK VOW LEDGRI EYT
The authors are intlt+tcd to Wilmer Jenkin. for his advice and of the s:mple* of purr assistance, especiallv i n the pr~p:~ratiori hypophosphorous acid. LITERATURE CITED
(1) Jenkins, W. 9., and Jones, R. T., J . Am. Chem. Soc., 74, 1353
(1952). (2) Kamecki, J., Roczniki Chem., 16, 199-206 (1936). (3) Kolthoff, I. M , , 2 . anal. Chem., 69, 36 (1926). (4) Koszegi, D., Ibid., 70, 347-9 (1927). (5) Rupp, E., and Fink, A,, Arch. Phnrm., 240, 665 (1902). (6) Wolf, L., and Jung, W,, 2. onorg. cillgem. Chem., 201, 337-46 (1931). RECEIVED for review January 27, 1953. Accepted May 13, 1953.
Simultaneous Colorimetric Determination of Copper, Cobalt, and Nickel as Diethyldithiocarbamates J . M. CHILTON, Oak Ridge !\-ational Laboratory, Oak Ridge, Tenn. colorimetric determination of small amounts of copper T with sodium diethyldithiocarbamate has become a classical procedure, being first proposed by Callan and Henderson in 1929
ANALYTICAL PROCEDURE
HE
( 2 ) . The fact that cobalt and nickel also give colored complexes with this reagent was noted by Holland and Ritchie in 1939 ( 4 ) , but its use for the determination of nickel was not proposed until 1946 (1). The fact that it is possible to determine copper, cobalt, and nickel simultaneously with this reagent was demonstrated by Nasanen and Tamminen (5),but they gave no details for the formation and separation of the metal dic,t,h?.ldithioc:~rl):imates. The method involves the formation of the carlmmates at, a pH of 8.5 to 9.0 in an aqueous solution to which pyrophosphate and citrate have been added and subsequent extraction of the colored complexes with carbon tetrachloride.
An aliquot of the sample, containing 5 to 25y of copper, 4 to 207 of cobalt, and/or 2 to 1Oy of nickel in a volume of 5 to 100 ml., is transferred to a 125-ml. separatory funnel. A similar amount of water is placed in another funnel to be used for n blank determination of the reagents employed. To the sample and blank are added 2 drops of phenol red indicator and 4.0 ml. of 4 % sodium pyrophosphate solution. If solution is not basic, ammonium hydroside is added to change the indicator color. citric acid is added, and the sample One milliliter of again made basic with ammonium hydroxide, If a precipitate forms, addition of more citric acid, followed hy ammonium hydroxide, will usually bring it into solution. To the c l e ~ solution are added 2.0nil. of 0.1% sodium diet,h!-ldithioi,art,nmate. Two milliliters of varkion tetr;tc~hlorideare added and the flask
I
APPARATUS AND REAGENTS
Transmittancy measurements were made with a Beckman D U spectrophotometer, using 1.0-cm. Corex cells and tungsten filament lamp. Extractions with carbon tetrachloride were made in 125-ml. pear-shaped separatory funnels. Stock solutions of the ions of copper, cobalt, and nickel were made by dissolving 0.01 mole of the metal in nitric acid and diluting to 1 liter with redistilled water. Spectrographic analysis of the metals showed no detectable cobalt or nickel in the copper, < O . l % copper and 0.4% nickel in the cobalt, and
