Cathodic Action of the Uranyl-Lactate Complex at the Dropping Mercury Electrode TSAI-TEH LA1 and 81-CHENG W A N G Department o f Chemical Engineering, Cheng Kung University, Tainan, Taiwan, China
b The complex of uranium with lactic acid has been stuclied polarographically. The limits 0.f reversibility and complexaiion, which are the function of pH and the concentration of lactic acid, are shown. Ella E7I8 are first used in deriving an equaiion for studying a polynuclear complex. Partial polymerization to form a dimer takes place in the reduced forni of the complex. The complex forms and electrode reactions in various givsn ranges of pH value and concentralion of complexing agent are proposed,
T
HE STOICHIOMETRY of the uranyllactate system in acid solution was determined by absoiytion spectra and Job's continuous variation method by Feldman and Havill ( 4 ~ who , found that only a combining ratio of unity ia cvident a t p H 3.5. The uranyl lactste complex was proved titrimetrically and spectrolhotometrically, in I n equimolar mixture to form a trimer before dimer formation is complete ( 5 ) . Breda and coworktm (1) have studied the polarography of oi-er 25 metal ions in organic h y d r o s j ~ acid supporting electrolytes. This work way untlcrtaken to present comidete informatio 1 about the iiolarogral)hic behavior of uranyl-lactate complex.
EXPERIMENTAL
Apparatus. -411 polarograms were taken with a Fisher Elecdropode in ronjunction with art external potentionie tcr
.
H-type polarographic cells containing a saturated calomel reference electrode separated from t h e solution compartment by a fine sintered-glass disk and a potassium chloride-agar plug were employed. The dropping mercury electrode had a flow rate of 1.445 mg per vcond, and a drop time of 4.866 seconds at -0.200 volt us. S.C.E. The polarographic cell was thermostatted at 30' & 0.01' C. Dissolved o\;ygen was removed by bubbling purified nitrogen through the solution 20 minutes before meawrement. pH measurements were made with a Beckman Model H-2 pH meter. Chemicals. Standard stock solution of uranyl perchlorate was prepared and standardized b y t h e method described in our previous paper (8). A 2 M stock solution of lactic acid mas prepared and the exact concentration was determined potentiometrically with carbonate-free sodium hydroxide solution which was standardized with potassium biphthalate of the Xational Bureau of Standards. Fisher certified reagent grade gelatin was used as a suppressor. RESULTS AND DISCUSSION
As can be observed from Figure 1, the half-wax potential is shifted to the negative side by the complexation of uranyl ion with lactate, and by the variation in pH value, showing that hydrogen ion is involved in the electrode reduction of the complex. In the experiments reported in this paper, the hydrogen ion concentration of the solutions was so low that the uranium(V) complex may be regarded as stable with respect to disproportionation, or i t may disproportionate at such a rate t h a t the reduction would not affect the half-wave potential and shape of the wave (3, 7 ) . The unsymmetry to half-ware of polarograms 111 and IV in Figure 1 is responsible for the inequality of a to b in the equation of Edt. = C1 CZ log +/(id
...-._.-. -*-
__
~. 0
-01
R
7 .
-02
Ed e
Figure 1 .
/
'
016
1 5
-1.1 5
-
--0.3 5
(
-04
,
-05
E , volt.
Typical polarogiams
X lO-3M U02:ClO&, 0.2M NoCIO~, 0.00570 gelatin 1.017
- i)b.
value of 0.049 =t0.002, 0.056 i 0.03 and greater than 0.06 volt exist. The temperature coefficient of halfwave potential corresponding to each category of polarograms is -0.5 mv./ ' C., -0.2 mv./O C., and -1.7 mv./' C. respectively, indicating that only the group of polarograms having E1'4 Ea14 value greater than 0.06 volt is irreversible. The group of polarograms with E1/4- Ea,&value of 0.049 volt, which is maintained constant for a wide range variation of p H value and concentration of chelating agent of the polarographic solution, is considered to be a reversible process, and the deviation from the theoretical value of 0.057 volt for 1electron reduction is attributed to the effect of polymerization as described later. The limitation of region of reversibility and chelation in relation to the concentration of lactic acid and p H ia shown in Figure 2. The electrode reaction is still rpversible even in low p H value, and becomes irreversible when the pH value is higher than a certain value, which is a function of the concentrations of lactic acid. Nature of Diffusion Current. The effect of temperature and the height of mercury column on t h e diffusion current was demonstrated b y recording the polarograms of a 1.017m.11 uranyl perchlorate solution in 0.2111 lactic acid and sodium perchlorate and 0.005% gelatin a t pH 3.5. The temperature coefficient of diCfusion current is 1.8% per degree, and
+
Reversibility. Reversibility of t h e electrode reaction was judged from the value of El 4 - Ea,, and the temperature coefficient of half-wave potential. I n an inspection of the value of El,4 - E ~ Mi t, was found that three categories of polarograms with the
PH
Figure 2. chelation
Limit of reversibility and
1.017 X 10-3M U02(CIOal2, 0.2M NaCIOI, 0.005% gelatin
VOL. 35, NO. 7, JUNE 1963
905
the value of id/h1I2is 0.438 A 0.006 for the height of mercury column from 59 cm. to 87 em. Therefore, i t was confirmed that the electrode reaction was entirely diffusioncontrolled. Effect of Polymerization on the Shape of Polarogram. If M f m and X-1 denote metallic ion and complexing agent, respectively, the half-cell reaction of the complex, J I a X p ( a m - p " , may be written in the form bMaXpcam-pl)
+ abne
X-'
06
08
I.0
(1)
The potential of the electrode must then be that given by the Nernst equation: 1 RT E d , c .= E" + -- - 1nK ab nF (bp
30
30
+
(bp -
ab n F
2.0
+
aMoXP(b(m-n) -ai)
1
b IO
- an) In
(Ella-
1 . RT - In 30+b E114 - ESl4= -
[X-l]
nF
1 RTln3
=
I + ;
b
b'nF
243
+ 602 (%) (3)
and
Table 1.
4.60
5.00
5.60 CHA= 0.376M 2.70 3.30 5.80 CHA= 0.756M 3.50 4.50 4.95
906
EI/~)
polarogram should produce a straight line of reciprocal slope 0.059,'bn which can be used for the criterion of reversibility of the electrode reduction. Table I gives the values obtained by applying some polarograms to Figure 3. The value of a is unity, indicating that the uranyl-lactate complex exists in monomer; however, the lalue of b increases with increasing the concentration of lactic acid, proving that polymerization occurs in the reduced form of the complex. Since the electrode reaction is reversible and potentials mere determined to the nearest 1 mv., the value of b estimated from Figure 3 must be reliable. The deviation of b irom integral number may be attributed to the existence of a mixture of monomer and dimer in the experimental condition. Half-wave Potential Is Function of Diffusion Current. When polymerization occurs. and if a # b , then from Equation 2 we get
Values of a and b in Various Polarographic Solutions
pH Value
CHA= 01047M 1.75 3.65 4.97 CHA= 0.095M 2.05 3.95 5.30 C H =~ 0.189M 2.70
(4)
where E7,8and are the potentials a t which the current is equal to seven eighths and one eighth of the diffusion current, respectively. I n an attempt to obviate the calculation in estimating the values of a and b, the relation among the values of (El,z E71s)/(Ein - Elid, b/a, (Em - E3141 and b is shown in Figure 3, which was prepared from Equations 3 and 4. For example, if we have (E112 - E718)/ (Ells - Ell*)= 1.33 and (E114- E3d = 0.043/n then, from Figure 3, we get b/a = 2 and b = 2 , therefore a = 1. With the value of a and 6 , a 1)lot of log i / ( i d - i ) * ' a 2's. Ed.c.for a given
(Writing CHICHOHCOOH a8 HA and CHA= [HA]
CHA= 0.024M 2.70 3.95
E~/@),/(EI,~-
+ [Av])
El12= E a
a
b
0.8 1.o
1.oo 0.91 0.94
0.059 0.060 0.053
1.o 1.o 1.1
1.01 1.02 1.01
0.059 0.051 0.057
1.o
1 .o 1.1 1 .o
0.90
0.059 0.057 0.057
1.1 1 .o 1 .o
0.9
1.01 1.06 1.13
1.oo
0.050 0.052 0.056
1 .o 1 .o 1 .o
1.2 1.3
1.14 1 .I5 1.02
0.051 0.049 0.055
1 .o i .O 1.o
1.3 1.4
1.I5
0.049 0.049 0.050
1 .o 1 .o 1 .o
1.4 1.4 1.3
1.oo
ANALYTICAL CHEMISTRY
1.17 1.12
14
Figure 3. Estimating b / a and b from (€11~ - €%)/(Ex- €11~) and Ex - €%
ab
I n the case, a # b, the polarogram is unsymmetrical to half-wave. The value of a and b can be estimated by the following equations which were derived from Equation 2 .
12
IO
1.1 1 .o
1 RT +ab n F --
In K -
GnF R T (bp - aq) In [ X - l ]
1 .o
1 .o 1 .o
The half-ware potential is no longer only a function of pH value and the concentration of chelating agent but also a function of diffusion current. a - b RT The term Eli* 7 In should be used instead of Elh alone in finding the ligand number. For a given a and b,
+
(*)
1.o
1.1
The id value was found to change around 15% for fivefold changing of CHA. Introducing the A h i d x-alue and a = 1, b = 1.4 into the equation,
- 0.40
C HA
PH
3C24
0
0 0
4.5 4.0
L L I
LA 6
-I 0
PH
Figure 4. \'ariation of
Eli2
with pH value
Figure 5.
1.01 7 X 1 O-3M UOz(CIO4)2,0.2M NaCIOn, 0.005% gelatin
-b
r ($)I!
1ZT Aln we get 1 mv., nF . . \I-hich is IF'ithin esperimental error. Thereby el,^ was used to find the number of hydrogen .ons involved and the ligand number of the complex. u
ab
Half-wave Potential and Chelate Form. T o determini? t h e composition of uranium lactate complex and its electrode reaction, the variat.ion of half-wave potential with pH value and with t h e concentration of lactic acid was measured (Figures 4 and 5). It is obvious from Figure 4 that the curves steepen rapidly when the reaction approaches irreversibility, showing that more hydroxyl ions are involved in the electrode reaction and the excessive hydroxyl ions cause i,he reaction to be irreversible. ilccording to the value of the slopes in Figures 4 and 5, five kinds of reactions are presumed to be predominant in the following range (See Figure 2). (A) pH < 3.8, < 0.4-11. The value of A E ~ / ~ / Ais ~ -0.052 H and AEl~zjAlogCHa is --0.057; therefore one hydrogen ion is involved in the electrode reaction and U(V1) has one more ligand than does U(V). It is known t h a t uranyl ion forms a 1:l complex with lactate ion in this condition ( 5 ) and we c o n c l u d ~ that ~ the U(V) is uncomplexed. I n most cases uranyl ion shows a coordiration number of four; therefore thc. uranyl lactate chelate is probably PH
HO, \
+
+
where K is the apparent equilibrium constant (including activity coefficients) for the reaction UO,(OH)2;\-
+ UOz+ + H + UOz(OH)2- + UOz+' + H.4 =
Using (El,z),as -0.18 volt ( 6 ) , pKa = 3.74 (Z), and data on Figure 4, we calculated K 9. Apldying Equation 7 for one unit of pH changing from 2.5 to 3.5, t'he theore& ea1 value of the change of half-wave potential is calculated to be (Ei,z)PH 3.5
- (Ei,s)pH 2.5
=
-0.05
Thus the deviation of the slope -0.052 from -0.06 on E,,z us. pH curve in this range was verified. (E)pH > 3.8, CEa < 0.1Jf. The slope of Eliz US. pH is -0 01 volt, presuming that no hydrogen ion is involved in the electrode reaction. The slope of El,*us. log CHAis -0.057 volt, hence the electrode reaction can be expressed as U02(0H)2Ae = UOQ(OH)~- .Li-
+
+
From this reaction we get
+ +
0.06 log K' 0.06 log [ A - ] ( E i / ~ ) s 0.06 log K' 0.06 log CEA
+ 0'06log (10-uH
+
)
~ O - P K ~ 10-pKa
H
-P
UO*(OH)z-
(8)
where K = K.KH.A From Equation 8 we obtain
and the electrode reaction is
+ H+ + e
(7)
10-pKa)
=
' 0
I
TJOZ(OH)ZA-
+
(Ei/t)s 0.06log K 0.06 log CHA 0.06log (10-p1
3.8, 0.1 0.4Jf. The value of L I E ~ ~ ~ . ' A ic ~ -0.056, H and AE,:Z/'Alog Ca.k is zwo, showing t,hnt one hydrogen ion is consumed in bhe reduction and lJ(V1) has the same number of ligand as C(V). The value of a and b in Table I shows that the polymerzation of G(V) takes place in t h k region. The electrode reactions are U02(0H)2A-
+ H+ + e = UOz(OH)A-
+ I120
and
- (Eii2)pH3.8 -0.01 volt
2U02(OH)A-
=:
(U02)2(0H)zAz-'
VOL. 35, NO. 7, JUNE 1963
907
(E) p H > 3.8, Cas> 0.73f. I n this region A E 1 , 2 1 ' 3 p H = -0.03 volt, hEl,z/'Al~g CHAE O > and b E 1 . 3 > ~ =1 Therefore,
ACKNOWLEDGMENT
(3) Davis, D. G., ANAL CHsnr. 33, 492
The authors thank thc Satiollal Council on Science Development for financial support of this project'.
( 4 ) Feldman, I., Havill, J. IT., J . Am.
-+ e = U02(OH)2L4-2
U02(0H)2A-
LITERATURE CITED
aiid
(1) Breda, E. J., Meites, L., Reddy, T. B.,
+
2u02(oH)2-l--2 21If
=
+
( uoz)z(o11)2.~2-221f20
!Test, P. W.,Anal. C h i m Acta 14, 300 (1956). ( 2 ) Cannan, It. li., Iiihrirk, .I., J . z4?ti. ChenL. SOC.60, 2314 ( I 938).
(1961).
(5y$$i:lk,;6: iil?l/,ly&
W.F., Ibzd., 76, 4726 (1954). (6) Harris, W.E., Kolthoff, I. M., Ibid., 69, 446 (1947). ( 7 ) Kolthoff, I. RI., Harris, W.E., Ibid., 68, 1175 (1946).
(8) Lai, T. T., Chang, T. L., AXAL.CHERI. 33, 1193 (1961).
RECEIVED for review Soveniber 13, 1'362. Accepted March 20, 1063.
Simultaneous Determination of Hypobromite, Bromite, and Bromate Using Ammonium Sulfate M. H. HASHMI and AYYAZ AHMAD AYAZ West Regional laboratories, C.S.I.R.,
Lahore, Pakistan
b The reaction between ammonium sulfate and hypobromite has been developed into a simple and accurate analytical procedure for the simultaneous determination of hypobromite, bromite, and bromate. The influence o f various salts has been discounted. Bromite and bromate added to hypobromite have been determined accurately. Chlorides and bromides, i f present, do not interfere in the determination. Phenol reacts with bromite and hypobromite, and cannot be used for determination o f the latter in the presence of bromite.
of ammonium sulfate in place of phenol has many advantages as it avoids the phenol-iodine reaction and renders the procedure simple and accurate. Animonium reacts selectively u-ith hypobromite, but not with bromite, and can be used for the estimation of hypobromite in the presence of bromite. The influence of various salts ha$ been eliminated; and hromite arid bromate, added to hypolirornite, h a w been determined accurate!).. C'hloridrk and bromides, if present, do not intcrfere with the determinatiol1,s. EXPERIMENTAL
D
of the kinetics and stability of hypohalites (L), there arose the need of a n analytical procedure for the simultaneous determination of hypobromite, bromite, and bromate ions. Many attempts have been made for the determination of hypobromite in the presence of bromite and bromate ions (1-3). Most of the methods (1, 3 ) are based on the use of phenol for the destruction of hypobromite and involve many complications. The slow action of iodine on phenol leads to 101%-results, and the end point with starch, though sharp, is not permanent (1). The method based on the oxidation of hypobromite to bromate also involves difficulties (6, 7 ) because the bromide present is also oxidized and leads to incorrect results. -4 method has been described ( 5 ) for the estimation of nitrogen uqing hypobromite. In the preqent paper, the same reaction between animoniuin sulfate and liypohroniite has been developed into a siiiiple and accurate analytical procedure for the simultaneous determination of hypobromitc, bromite, and bromate ions. The use
908
URISG THE STUDY
ANALYTICAL CHEMISTRY
Materials. 1111 reagents were analytical grade. Sodium hypobromite was prepared by dissolving 3.5 mi. of bromine in 250 ml. of 10% sodium hydroxide solution a t 0" C. and then making u p to 1 liter. It was stored in a n amber bottle a t room temperature (20" to 35" (3.). Sodium bromite was prepared by the procedure of Chapin ( 1 ) and standardized with arsenious oxide each time ~
Table 1. Recovery o f Added Hypobromite, Bromite, and Bromate from Hypobromite Solution hlillimoles .- __
Comp~nent Added Hypobromite 41 .30
Found
nevi ation
i6.12 3 21
40.9,j 60.25 68.37
76.00 3.24
-0.35 -0.07 +0.15
2.80 2.44 (5.70
2.i6
60.32 68.22
Bromite I3rornate
10.16 13.41 16.70
2.41 6 O!)
10.25 I:< 40
16.78
-0.12
+0.0:3 -0 04 -0.08 -0.01
+00!)
-0.01 $0.08
before use. Secessary correction for ?odium bromate present in sodium bromite solution was always made by potassium iodide-thiosulfate titration. Arsenious oxide (0.12~1: solution) mas prepared by dissolving approximately weighed reagent in 10% sodium hydroxide solution. The solution was made acid to phenolphthalein. Ten grams of solid sodium bicarbonate was added, and the solution was diluted to 1 liter and standardized with iodine. Potassium iodate (0.1~17 solution) n-as prepared from the rcagent previously dried a t 120" C. Procedure. Potassium iodate was used as the primary standard, and the determination was carried out by t h e following methods: (a) To 5 ml. of a niisture containing hypobromite, bromite, and bromate, mas added 3 to 4 grams of solid potassium iodide, followed by 10 ml. of 4 S sulfuric acid. T h e solution was diluted t o twice its volume, and the liberated iodine was titrated with thiosulfate using starch as indicator. 'The titer included h!-pobromite, bromite, and bromate. (b) Five milliliters of mixture was added to a flask containing an excess of ammonium sulfate and about 1 gram of sodium bicarbonate. rlfter 10 minutes, 3 to 4 grams of solid potassium iodide and 10 ml. of 41V sulfuric acid lvas added, and the solut'ion was let stand for 5 minutes before it was diluted to twice its volume and titrated with thiosulfate. The titer included bromite and bromate. (e) To 5 nil. of mixture, a known excess of st,andard arsenious oxide was added. After 5 minutes, 4 to 5 grams of sodium bicarbonate was added fol!owed by dilute acetic acid to neutralize sodium hydroxide until each drop of acetic acid gave free effervrscence. The solution was t i t r a t d with iodine using starch as indicator from IThich the combined normality of hypobromite and bromite was dctermined.
