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Stability constants of the chloride and nitrate complexes of neptunium (V) and neptunium (VI)
J. inorg, nucl.Chem., 1970,Vol. 32, pp. 977 to 986. PergamonPress. Printedin Great Britain
STABILITY NITRATE
CONSTANTS COMPLEXES
OF THE CHLORIDE OF NEPTUNIUM (V) NEPTUNIUM (VI)
AND AND
N. S. A I - N 1 A I M I * , A. G. W A I N and H. A. C. M c K A Y Applied C h e m i s t r y Division, A.E.R.E., Harwell, Berkshire, England
(Received 30 July 1969) A b s t r a c t - S t a b i l i t y c o n s t a n t s of the chloride and nitrate c o m p l e x e s of n e p t u n i u m (V) and n e p t u n i u m (VI) have been determined in HCIO4 of concentrations ranging from 0.3 to 0.8. This was done by measuring the potentials of chemical cells containing equimolar proportions of n e p t u n i u m (V) and n e p t u n i u m (VI) along with varying concentrations of the complexing ligands. T h e valency of the n e p t u n i u m was adjusted by controlled-potential coulometry. F r o m the results a preferred m e a n value of the stability c o n s t a n t of NpO2CI + at 25 _ 0 . 0 1 ° C in 0.3-0.5 M H(CIO4, C I ) o f 0 . 4 6 ± 0 . 0 3 mole -z I. was obtained. For NpO2NO.~ + there appears to be s o m e variation in the stability constant, from 0.105+_0.005 mole -I I. in 0.4 M H(CIO4, NO3) to 0 . 1 3 0 ± 0 . 0 6 mole 11. in 0.6 M H(CIO4, NO:0. T h e r e was s o m e indication of the formation of small a m o u n t s of the second complexes, N pOzCI2 and N pO._,(N O:~).,. INTRODUCTION
VERY little is known about the stability constants of the chloride and nitrate complexes in aqueous solutions of neptunium (V) and neptunium (VI). This paper describes measurements in a perchloric acid medium at 25°C. The method was to determine the potentials of cells of the type HcHXIo 4
l~t NpO2CIO4 NpO2(CIO4)2
C Cx. - - -
I HC104 Cx INpO2ClO4 ~Cvp~CvP NpO2(C104)2
1 '
½c),,
Liquid junction (X = CI or NO3), where CH had values in the range 0.3-0.8 M, Cx never exceeded about 0.2 M, and Cup was equal to 3.96 × 10-4 M. Since complexing was weak, the potentials, AE, of the cells were small (<1.5 mV), and very careful work was necessary. An important feature was the use of controlled potential coulometry to adjust the valency of neptunium. EXPERIMENTAL
Cells T h e cell for the potential m e a s u r e m e n t s consisted of two stoppered glass c o m p a r t m e n t s each of 35 ml capacity, separated by a N o . 4 glass sinter impregnated with precipitated silica. Diffusion between c o m p a r t m e n t s in such a cell a m o u n t e d to < I per cent of the solution in a c o m p a r t m e n t over *All c o r r e s p o n d e n c e to be a d d r e s s e d to N u c l e a r Research Centre, T u w a i t h a , Baghdad, iraq. 977
978
N . S . AI-NIAIMI, A. G. WAIN and H. A. C. McKAY
24 hr. Each compartment contained a bright platinum electrode and was provided with means for bubbling nitrogen for stirring and to give an inert atmosphere. The neptunium valency was adjusted in a coulometer cell with three compartments separated by No. 4 glass sinters impregnated with precipitated silica. The centre compartment contained neptunium in perchloric acid of the appropriate molarity, a bright platinum working electrode made from a 2.5 cm square piece of 36 gauge platinum gauze shaped into the form of a cylinder and welded to a 15 cm length of platinum wire, and was provided with means for bubbling nitrogen. The two outer compartments contained perchloric acid of the same molarity as the acid in the centre compartment; in one was a platinum counter electrode and in the other a calomel reference electrode. The cell was used in conjunction with a Harwell type coulometer 95/2220-1/6 Serial No. B.7 [1]. The platinum electrodes were cleaned before each experiment by washing in hot aqua regia, rinsing with distilled water and heating in a bunsen flame.
Preparation of solutions Neptunium (V) solution. Neptunium dioxide was dissolved in boiling 12 M nitric acid containing 0.1 M fluoride. Successive evaporations to near dryness followed by further additions of strong nitric acid were carried out to eliminate the fluoride. After a final evaporation the neptunium was dissolved in 1 M nitric acid and oxidised to the hexavalent state. Sodium nitrite was added to reduce neptunium to the pentavalent state and neptunium (V) hydroxide was precipitated by adding excess sodium hydroxide. The precipitate was washed ten times with dilute sodium hydroxide and dissolved in 0.5 M perchloric acid. The hydroxide was reprecipitated, washed as before and dissolved in perchloric acid to give a solution of neptunium (V) in 0.5 M perchloric acid.
Acids Nitric, hydrochloric and perchloric acid solutions were prepared by diluting the Analytical Reagent concentrated acids. The nitric acid solutions were freshly prepared before each experiment from solutions that had been purged with oxygen-free nitrogen.
Analysis Neptunium. This was analysed by a-counting using a Simpson proportional counter and by aspectrometry using either an ion-chamber or a solid-state detector in conjunction with a Laben 512-channel analyser. A further check on the neptunium concentration was made with the controlledpotential coulometer.
Acidity determinations The acidity of the solutions was determined by titration with standard sodium carbonate solution, using a mixed indicator of bromocresol green and dimethyl yellow.
Procedure Samples of neptunium (V) solution were oxidised completely to neptunium (VI) and then half reduced in the controlled potential coulometer, giving a solution containing a 1 : 1 mixture of neptunium (V) and neptunium (V1). This method of valency adjustment avoids the uncertainties of ordinary chemical methods. For the potential measurements equal volumes of this solution were transferred to the two halves of the cell, which contained equal volumes of perchloric acid of the appropriate molarity. The initial volume of solution in each half-cell before titration was 20 ml. Oxygen-free nitrogen was bubbled through each compartment and the solution in one half-cell was titrated from a polythene weighing bottle with either hydrochloric or nitric acid of the same molarity as that in the cell. In this way the acid concentrations of both half-cells were kept constant during the titration. Before each titration urea was added to the nitric acid solutions at a concentration of 1 per cent of that of the acid, at least 2 hr before use, to suppress the formation of nitrous acid, which reduces neptunium (VI) to neptunium (V). Sulphamic acid and hydrazine were rejected as alternatives because it was not possible to obtain stable potential readings when these reagents were used. The potential of the cell was measured to ---0.02 mV using a digital voltmeter (Solartron type LM 1. G. W. C. Milner and G. Phillips, Coulometry in Analytical Chemistry, p. 69. Pergamon Press, Oxford (1967).
Stability c o n s t a n t s
97 9
1420.2). T h e readings were stable before and during the titration and ranged between (+0.10_+ 0.01l and (--0.10 + 0.01) m V for the different e x p e r i m e n t s before the titration c o m m e n c e d . After addition of the titrant readings were taken at intervals of 15 min, which allowed time for equilibrium to be established. All e x p e r i m e n t s were carried out at 25 -+ 0-01°C. RESULTS
AND
DISCUSSION
F o r both chloride and nitrate complexing we have the reactions NpOf ++ nX- ~
NpO,~X,, (~-'')+
N p O 2 + + n X - ~- NpO,,Xn (1 ")+ where X is CI or NO3, n = !, 2, 3 . . . . . and 6fi,, and 5fi,, are the complexity constants of neptunium (VI) and neptunium (V) respectively. T h e potential AE at 25°C of the cells used is given by A E = 0 . 0 5 9 1 5 5 log
[(l+6~a+6fiza~+...)/(l+Sfl~a+Sfl2aZ+...)]+Ej.
where a is the anion concentration and E~ is the liquid junction potential. Tables 1 and 2 give the experimental results. Attempts to work at lower acid concentrations did not give reproducible results. Other authors[2-4] have reported reduction of neptunium (VI) in 1 M or 2 M HCI solutions in the presence Table 1. Potential of the chemical cell for Np(V)/Np(V1) chloride s y s t e m s in H CIO4 of different concentrations (Cn)
CH = 0"4
CH = 0-3
(mV)
[Cl-] (× 102)
0.24 0.46 0.66 0.85 1.03 1.19 1.30
2.22 4.14 5.80 7.28 8.58 9.74 10.75
AE
C~ = 0"5
(mV)
[CI-] (× 102)
AE (mV)
(× 10z)
0.12 0.30 0.30 0.36 0.51 0.55 0.60 0.68 0.80 0.99 1.16 1.34
1.29 2.50 2.95 3.54 4.72 5.52 5.72 6.68 7.75 9.70 11.41 12.95
0.19 0.38 0.60 0.76 0.91 1.03 I. 17 1-30 1.39 1.47
1.92 3.70 5.36 6.90 8.33 9.68 10.95 12.10 13.20 14.25
AE
[el-]
T h e limit of error on the A E readings is --+0.02. 2. L. B. M a g n u s s o n , J. C. H i n d m a n and T. J. La Chapelle, The Transuranium Elements (Edited by G. T. Seaborg, J. J. Katz and W. M. Manning), Nat. Nucl. Energy Ser. Division IV, Vol. 14B, paper 15.4. McGraw-Hill, N e w York (1949). 3. R . W . Stromatt, R. M. P e e k e m a and F. A. Scott, H W - 5 8 2 1 2 (1958). 4. D. C o h e n and B. Taylor, J. inorg, nucl. Chem. 22, 151 ( 1961 ).
980
N . S . AI-NIAIMI, A. G. WAIN and H. A. C. McKAY
Table 2. Potential of the chemical cell for Np(V)/Np(VI) nitrate systems in HCIO4 of different concentrations (Cn) Cn = 0.4
CH = 0.6
CH = 0.5
CH = 0'8
AE (mV)
[NO3]total (× 102)
AE (mY)
[SO3]tota, (× 102)
AE (mY)
[SOa]tota I (× 10z)
AE (mY)
[SOa]tota l (× 10z)
0.06 0.12 0.15 0.21 0.24
2.96 5-51 6.66 8-75 10.60
0.10 0.18 0.22 0-28 0.34 0.40 0.46
3.7 6.9 9.7 12-1 14.2 16.2 18.0
0" 17 0.24 0-39 0.43 0.48 0-48 0.58
6.42 8.29 13-1 14.5 15.9 17.1 19.5
0.07 0.13 0-23 0.31 0-37 0.41 0.48 0-51 0.54
3.07 5.94 8.59 11.10 13.35 15.45 17.50 19.4 21.2
The limit of error on the AE readings is _+0.02.
of metallic platinum or gold, making accurate potential measurements impossible, but we observed stable potentials for a period of at least four hours at chloride concentrations up to 0.15 M and hydrogen ion concentrations up to 0.5 M. Liquid junction potentials were calculated from Henderson's equation. We feel justified in using this equation in the light of Smyrl and Newman's [5] recent study. They showed that even for junctions of a more extreme character than ours, the potentials vary very little with the method of making the junction or with the ionic strength, and that the Henderson equation provides a good approximation [6]. In any case, the effect on the calculated B-values of allowing for the liquid junction potentials is scarcely outside our limits of error. Figures 1 and 2 show plots of Y against a, where 0.059155 log Y = A E - - E j Y = (1 + ~ l a + eB2a2 + . . . ) / (1 + ~ l a + 5B2a2 + . . . ). At each point the vertical line indicates the limits (__+0.02mV) set by the precision of reading the digital voltmeter. The values of a for nitrate have been corrected for a, the degree of dissociation of nitric acid, using values reported by Davis and De Bruin [7]. It will be seen that to a fair degree of approximation Y= l+ba, where b is a constant, and that ignoring E j has little effect. The obvious interpretation is then to equate b with eB1 and to equate all the other B-values to zero. However, this interpretation cannot be accepted without reservation. 5. W. H. Smyrl and J. Newman,J. phys. Chem. 72, 4660 (1968). 6. P. Henderson, Z. phys. Chem. (Leipzig) 59, 118 ( 1907); 63, 325 (1908). 7. W. Davis and H. J. De Bruin, J. inorg, nucl. Chem. 26, 1069 (1964).
Stability constants
981
2
-
1.06
I
1-04
1.06
~I.02
1.04
1.02
--
1.oo 0
I 2
I 4
I 6
I 8
l I0
I 12
I 14
a
1 . 0 0
I I I 16 18 2 0 x I0 2
Fig. 1. Variation of Y of the Np(V)/Np(Vl)-chloride system in HCIO4 with CIat C , = 0.5. ( 1) Ye×P.vs. [CI ] ; (2) Y...... tea (allowing for E~) vs. [CI-].
In the first place we cannot absolutely exclude the formation of neptunium (V) complexes, because (1 + 6/31a)/(1 + ~/3,a) ~ 1 + (6/3_ ~/3~)a when 5/31a <~ 1 ; this is the same linear law as before, but with b now identified as 6/31- 5/31. Nevertheless, we do not think it likely that ~/31 is significant in comparison with 6/31 because we have obtained no definite indication of neptunium (V) complexing even with the more highly complexing anions fluoride and sulphate [8]. 8. N. S. AI-Niaimi, A. G. Wain and H. A. C. McKay. To be published.
982
N . S . AI-NIAIMI, A. G. WAIN and H. A. C. McKAY
-
1.03
1-03
--
1.02
1.02
--
1.01
1.01
--
1.00
Y
1.00 0
2
4
6
8
I0
12
14 0
16 18 2 0 x
102
Fig. 2. Variation of Y of the Np(V)/Np(VI)-nitrate system in HCIO4 with [NO3]tot.i at CH = 0.8. (1) Yexp.vs. [NO3]total; (2) Y~..... ted (allowing for Ej) vs. [NO3]total.
It is also possible that a second neptunium (VI) chloride or nitrate complex is formed. This possibility is best explored by plotting ( Y - - 1 ) / a against a, as suggested by Leden[9]. If Y=
l + ba + ca 2,
then ( Y - - 1)/a = b + c a , 9. See F . J . C . Rossotti and H. Rossotti, The Determination of Stability Constants, p. 108. McGrawHill, New York (1961).
Stability c o n s t a n t s
983
so that the intercept is equal to b and the slope to c. In the absence of neptunium (V) complexing c can be identified with 6/32. T h e plots, shown in Figs. 3 and 4, show a fair amount of scatter. This is to be expected, since we are c o n c e r n e d with a numerically small ca 2 term, amounting only to a minor correction to the Y-values and having only a small effect on the calculated values of b. We cannot hope to obtain more than a rough estimate of the values of c f r o m our data, even though these values may be of appreciable magnitude. T h e results for C n = 0.4 and 0-5 in the chloride s y s t e m seem consistent with c-values between 0 and 0,4. I f c = 0 we should have constant values o f ( Y - l)/a, while if c = 0-4 we should have a line of the slope indicated. F o r C , = 0.3 the
0-58 0.56 0.54 0-52 0.50 0-48 046 0.44 0-4:
I
,
t
)-48
,
0.46 0.44 i
121 0-42
( Y-I)/=
D-40 0-38 0.36 0.34 ,.32 0-50 0.48 0-46 (3)
044 0,42 0-40 0.38
i
1 I l l 1 2 3 4
I S
I 6
l 7
l I t 8 9 I0 0 x I0 2
~ II
I 12
I 13
J 14
t IS
I 16
17
Fig. 3. Variation of ( Y - l)/a o f the Np(V)/Np(VI)-chloride s y s t e m in HCIO4 with C1-1 at (l) Co = 0.3; (2) Cn = 0.4 and (3) Cn = 0.5 (Y values corrected for liquid j u n c t i o n potentials).
984
N. S. AI-N1AIMI, A. G. W A I N and H. A. C. McKAY
(I)
t
0.14 0.12 :3.10 D.08
0.14 0.12
Y-I)/o
(2)
0.10 0.08
D.14
(Y-I)/o
~3.12
(3)
:3-I0 0-08
0.14 0-12 040 0.0~
2
4
6
8
IO
12 14 16 oxlO 2
18
20 22 24
Fig. 4. Variation of (Y-- 1)/a of the Np(V)/Np(VI)-nitrate system in HCIO4 with [NO3]total at (l) C a = 0.4; (2) CH----0-5; (3) CH = 0"6 and (4) CH -- 0.8 (Y values corrected for liquid junction potentials).
trend of the points suggests a steeper slope, with say c -- 0.8. In the nitrate system c may lie between 0 and 0.1. The effect on the values of 6/31 in both systems will never amount to a reduction of more than - l 0 per cent. The experimental values of b, for the various values of c just indicated, are given in Table 3. The "preferred values" of b correspond to what seem the most likely values ofc. For the chloride system any variation in b with CH lies within the experimental error, but for the nitrate system there seems to be a definite concentration dependence. The only previously reported investigation of the stability constants of the chloride and nitrate complexes of neptunium (V) and neptunium (VI) is that of Cohen et a/.[10], who studied the effect of nitrate and chloride ions on the rate of the isotopic exchange reaction of neptunium (V) and neptunium (VI). They found that nitrate ion appears to have little effect, but chloride ion exerts a marked catalytic effect. They interpreted their results in terms of formation of chloride complexes of neptunium (VI) but not of neptunium (V). Extrapolation of their data to 25°C gives 6/31 - 0.42 mole -11., which is in reasonable agreement with our results. No value has been reported for 6/31for the NpO2NO3 ÷ complex. However, Keder[11] has reported a study of the absorption spectra of neptunium (VI) in 4 M H(CIO4, NOz) solutions at wavelengths of 1230nm. The shifts in the apparent molar extinction coefficients are small and difficult to interpret, but are 10. D. Cohen, J. Sullivan and J. C. Hindman, J. Am. chem. Soc. 77, 4964 (1955). 11. W. E. Keder, J. inorg, nucl. Chem. 16, 138 (1960).
Stability constants
©
985
g....ggg r-,
4-1 +1 +1 +1
E
gggg
©
2
.=. Z6&&&
+~
© Z Z
~-I +1 +1 +E
+ © Z
8
i
+1 +1 lq
+~
~ Z Z 0
+1 +1 +P r,,,
I
+1 +1 +1
986
N . S . AI-NIAIMI, A. G. WAIN and H. A. C. McKAY
compatible with our results. Kirkbright and Yoe [ 12] report no variations of molar extinction coefficients in solutions from 0.16 to 2.7 M HNO~, but they found a decrease of 70 per cent in 6.9 M HNOa and of 6 per cent in a solution of 5 M NaNO3 and 1.9 M HNO3. Their results again indicate that the complexity constants must be small. Rykov and Yakovlev[13], in their paper on the kinetics of the reaction between neptunium (IV) and neptunium (VI) in nitrate solution, state, without giving experimental evidence, that from an analysis of the dependence of the molar extinction coefficient on [NO3] only one complex N p O ~ N O J is formed at [NO3-] < 2 M, with a stability constant of 0.4 mole -11. at 25°C. They then use this value to interpret their kinetic results. Although obviously rather uncertain, this value provides further general confirmation that 6/31is quite small. Acknowledgements-The authors thank Mr. R. A. P. Wiltshire for the a-pulse analysis and Dr. 1. L. Jenkins for useful discussions. 12. G. F. Kirkbright and J. H. Yoe, A nalyt. Chem. 35, 806 (1963). 13. A. G. Rykov and G. N. Yakovlev, Soviet Radiochem. 8, 26 (1966).
STABILITY NITRATE
CONSTANTS COMPLEXES
OF THE CHLORIDE OF NEPTUNIUM (V) NEPTUNIUM (VI)
AND AND
N. S. A I - N 1 A I M I * , A. G. W A I N and H. A. C. M c K A Y Applied C h e m i s t r y Division, A.E.R.E., Harwell, Berkshire, England
(Received 30 July 1969) A b s t r a c t - S t a b i l i t y c o n s t a n t s of the chloride and nitrate c o m p l e x e s of n e p t u n i u m (V) and n e p t u n i u m (VI) have been determined in HCIO4 of concentrations ranging from 0.3 to 0.8. This was done by measuring the potentials of chemical cells containing equimolar proportions of n e p t u n i u m (V) and n e p t u n i u m (VI) along with varying concentrations of the complexing ligands. T h e valency of the n e p t u n i u m was adjusted by controlled-potential coulometry. F r o m the results a preferred m e a n value of the stability c o n s t a n t of NpO2CI + at 25 _ 0 . 0 1 ° C in 0.3-0.5 M H(CIO4, C I ) o f 0 . 4 6 ± 0 . 0 3 mole -z I. was obtained. For NpO2NO.~ + there appears to be s o m e variation in the stability constant, from 0.105+_0.005 mole -I I. in 0.4 M H(CIO4, NO3) to 0 . 1 3 0 ± 0 . 0 6 mole 11. in 0.6 M H(CIO4, NO:0. T h e r e was s o m e indication of the formation of small a m o u n t s of the second complexes, N pOzCI2 and N pO._,(N O:~).,. INTRODUCTION
VERY little is known about the stability constants of the chloride and nitrate complexes in aqueous solutions of neptunium (V) and neptunium (VI). This paper describes measurements in a perchloric acid medium at 25°C. The method was to determine the potentials of cells of the type HcHXIo 4
l~t NpO2CIO4 NpO2(CIO4)2
C Cx. - - -
I HC104 Cx INpO2ClO4 ~Cvp~CvP NpO2(C104)2
1 '
½c),,
Liquid junction (X = CI or NO3), where CH had values in the range 0.3-0.8 M, Cx never exceeded about 0.2 M, and Cup was equal to 3.96 × 10-4 M. Since complexing was weak, the potentials, AE, of the cells were small (<1.5 mV), and very careful work was necessary. An important feature was the use of controlled potential coulometry to adjust the valency of neptunium. EXPERIMENTAL
Cells T h e cell for the potential m e a s u r e m e n t s consisted of two stoppered glass c o m p a r t m e n t s each of 35 ml capacity, separated by a N o . 4 glass sinter impregnated with precipitated silica. Diffusion between c o m p a r t m e n t s in such a cell a m o u n t e d to < I per cent of the solution in a c o m p a r t m e n t over *All c o r r e s p o n d e n c e to be a d d r e s s e d to N u c l e a r Research Centre, T u w a i t h a , Baghdad, iraq. 977
978
N . S . AI-NIAIMI, A. G. WAIN and H. A. C. McKAY
24 hr. Each compartment contained a bright platinum electrode and was provided with means for bubbling nitrogen for stirring and to give an inert atmosphere. The neptunium valency was adjusted in a coulometer cell with three compartments separated by No. 4 glass sinters impregnated with precipitated silica. The centre compartment contained neptunium in perchloric acid of the appropriate molarity, a bright platinum working electrode made from a 2.5 cm square piece of 36 gauge platinum gauze shaped into the form of a cylinder and welded to a 15 cm length of platinum wire, and was provided with means for bubbling nitrogen. The two outer compartments contained perchloric acid of the same molarity as the acid in the centre compartment; in one was a platinum counter electrode and in the other a calomel reference electrode. The cell was used in conjunction with a Harwell type coulometer 95/2220-1/6 Serial No. B.7 [1]. The platinum electrodes were cleaned before each experiment by washing in hot aqua regia, rinsing with distilled water and heating in a bunsen flame.
Preparation of solutions Neptunium (V) solution. Neptunium dioxide was dissolved in boiling 12 M nitric acid containing 0.1 M fluoride. Successive evaporations to near dryness followed by further additions of strong nitric acid were carried out to eliminate the fluoride. After a final evaporation the neptunium was dissolved in 1 M nitric acid and oxidised to the hexavalent state. Sodium nitrite was added to reduce neptunium to the pentavalent state and neptunium (V) hydroxide was precipitated by adding excess sodium hydroxide. The precipitate was washed ten times with dilute sodium hydroxide and dissolved in 0.5 M perchloric acid. The hydroxide was reprecipitated, washed as before and dissolved in perchloric acid to give a solution of neptunium (V) in 0.5 M perchloric acid.
Acids Nitric, hydrochloric and perchloric acid solutions were prepared by diluting the Analytical Reagent concentrated acids. The nitric acid solutions were freshly prepared before each experiment from solutions that had been purged with oxygen-free nitrogen.
Analysis Neptunium. This was analysed by a-counting using a Simpson proportional counter and by aspectrometry using either an ion-chamber or a solid-state detector in conjunction with a Laben 512-channel analyser. A further check on the neptunium concentration was made with the controlledpotential coulometer.
Acidity determinations The acidity of the solutions was determined by titration with standard sodium carbonate solution, using a mixed indicator of bromocresol green and dimethyl yellow.
Procedure Samples of neptunium (V) solution were oxidised completely to neptunium (VI) and then half reduced in the controlled potential coulometer, giving a solution containing a 1 : 1 mixture of neptunium (V) and neptunium (V1). This method of valency adjustment avoids the uncertainties of ordinary chemical methods. For the potential measurements equal volumes of this solution were transferred to the two halves of the cell, which contained equal volumes of perchloric acid of the appropriate molarity. The initial volume of solution in each half-cell before titration was 20 ml. Oxygen-free nitrogen was bubbled through each compartment and the solution in one half-cell was titrated from a polythene weighing bottle with either hydrochloric or nitric acid of the same molarity as that in the cell. In this way the acid concentrations of both half-cells were kept constant during the titration. Before each titration urea was added to the nitric acid solutions at a concentration of 1 per cent of that of the acid, at least 2 hr before use, to suppress the formation of nitrous acid, which reduces neptunium (VI) to neptunium (V). Sulphamic acid and hydrazine were rejected as alternatives because it was not possible to obtain stable potential readings when these reagents were used. The potential of the cell was measured to ---0.02 mV using a digital voltmeter (Solartron type LM 1. G. W. C. Milner and G. Phillips, Coulometry in Analytical Chemistry, p. 69. Pergamon Press, Oxford (1967).
Stability c o n s t a n t s
97 9
1420.2). T h e readings were stable before and during the titration and ranged between (+0.10_+ 0.01l and (--0.10 + 0.01) m V for the different e x p e r i m e n t s before the titration c o m m e n c e d . After addition of the titrant readings were taken at intervals of 15 min, which allowed time for equilibrium to be established. All e x p e r i m e n t s were carried out at 25 -+ 0-01°C. RESULTS
AND
DISCUSSION
F o r both chloride and nitrate complexing we have the reactions NpOf ++ nX- ~
NpO,~X,, (~-'')+
N p O 2 + + n X - ~- NpO,,Xn (1 ")+ where X is CI or NO3, n = !, 2, 3 . . . . . and 6fi,, and 5fi,, are the complexity constants of neptunium (VI) and neptunium (V) respectively. T h e potential AE at 25°C of the cells used is given by A E = 0 . 0 5 9 1 5 5 log
[(l+6~a+6fiza~+...)/(l+Sfl~a+Sfl2aZ+...)]+Ej.
where a is the anion concentration and E~ is the liquid junction potential. Tables 1 and 2 give the experimental results. Attempts to work at lower acid concentrations did not give reproducible results. Other authors[2-4] have reported reduction of neptunium (VI) in 1 M or 2 M HCI solutions in the presence Table 1. Potential of the chemical cell for Np(V)/Np(V1) chloride s y s t e m s in H CIO4 of different concentrations (Cn)
CH = 0"4
CH = 0-3
(mV)
[Cl-] (× 102)
0.24 0.46 0.66 0.85 1.03 1.19 1.30
2.22 4.14 5.80 7.28 8.58 9.74 10.75
AE
C~ = 0"5
(mV)
[CI-] (× 102)
AE (mV)
(× 10z)
0.12 0.30 0.30 0.36 0.51 0.55 0.60 0.68 0.80 0.99 1.16 1.34
1.29 2.50 2.95 3.54 4.72 5.52 5.72 6.68 7.75 9.70 11.41 12.95
0.19 0.38 0.60 0.76 0.91 1.03 I. 17 1-30 1.39 1.47
1.92 3.70 5.36 6.90 8.33 9.68 10.95 12.10 13.20 14.25
AE
[el-]
T h e limit of error on the A E readings is --+0.02. 2. L. B. M a g n u s s o n , J. C. H i n d m a n and T. J. La Chapelle, The Transuranium Elements (Edited by G. T. Seaborg, J. J. Katz and W. M. Manning), Nat. Nucl. Energy Ser. Division IV, Vol. 14B, paper 15.4. McGraw-Hill, N e w York (1949). 3. R . W . Stromatt, R. M. P e e k e m a and F. A. Scott, H W - 5 8 2 1 2 (1958). 4. D. C o h e n and B. Taylor, J. inorg, nucl. Chem. 22, 151 ( 1961 ).
980
N . S . AI-NIAIMI, A. G. WAIN and H. A. C. McKAY
Table 2. Potential of the chemical cell for Np(V)/Np(VI) nitrate systems in HCIO4 of different concentrations (Cn) Cn = 0.4
CH = 0.6
CH = 0.5
CH = 0'8
AE (mV)
[NO3]total (× 102)
AE (mY)
[SO3]tota, (× 102)
AE (mY)
[SOa]tota I (× 10z)
AE (mY)
[SOa]tota l (× 10z)
0.06 0.12 0.15 0.21 0.24
2.96 5-51 6.66 8-75 10.60
0.10 0.18 0.22 0-28 0.34 0.40 0.46
3.7 6.9 9.7 12-1 14.2 16.2 18.0
0" 17 0.24 0-39 0.43 0.48 0-48 0.58
6.42 8.29 13-1 14.5 15.9 17.1 19.5
0.07 0.13 0-23 0.31 0-37 0.41 0.48 0-51 0.54
3.07 5.94 8.59 11.10 13.35 15.45 17.50 19.4 21.2
The limit of error on the AE readings is _+0.02.
of metallic platinum or gold, making accurate potential measurements impossible, but we observed stable potentials for a period of at least four hours at chloride concentrations up to 0.15 M and hydrogen ion concentrations up to 0.5 M. Liquid junction potentials were calculated from Henderson's equation. We feel justified in using this equation in the light of Smyrl and Newman's [5] recent study. They showed that even for junctions of a more extreme character than ours, the potentials vary very little with the method of making the junction or with the ionic strength, and that the Henderson equation provides a good approximation [6]. In any case, the effect on the calculated B-values of allowing for the liquid junction potentials is scarcely outside our limits of error. Figures 1 and 2 show plots of Y against a, where 0.059155 log Y = A E - - E j Y = (1 + ~ l a + eB2a2 + . . . ) / (1 + ~ l a + 5B2a2 + . . . ). At each point the vertical line indicates the limits (__+0.02mV) set by the precision of reading the digital voltmeter. The values of a for nitrate have been corrected for a, the degree of dissociation of nitric acid, using values reported by Davis and De Bruin [7]. It will be seen that to a fair degree of approximation Y= l+ba, where b is a constant, and that ignoring E j has little effect. The obvious interpretation is then to equate b with eB1 and to equate all the other B-values to zero. However, this interpretation cannot be accepted without reservation. 5. W. H. Smyrl and J. Newman,J. phys. Chem. 72, 4660 (1968). 6. P. Henderson, Z. phys. Chem. (Leipzig) 59, 118 ( 1907); 63, 325 (1908). 7. W. Davis and H. J. De Bruin, J. inorg, nucl. Chem. 26, 1069 (1964).
Stability constants
981
2
-
1.06
I
1-04
1.06
~I.02
1.04
1.02
--
1.oo 0
I 2
I 4
I 6
I 8
l I0
I 12
I 14
a
1 . 0 0
I I I 16 18 2 0 x I0 2
Fig. 1. Variation of Y of the Np(V)/Np(Vl)-chloride system in HCIO4 with CIat C , = 0.5. ( 1) Ye×P.vs. [CI ] ; (2) Y...... tea (allowing for E~) vs. [CI-].
In the first place we cannot absolutely exclude the formation of neptunium (V) complexes, because (1 + 6/31a)/(1 + ~/3,a) ~ 1 + (6/3_ ~/3~)a when 5/31a <~ 1 ; this is the same linear law as before, but with b now identified as 6/31- 5/31. Nevertheless, we do not think it likely that ~/31 is significant in comparison with 6/31 because we have obtained no definite indication of neptunium (V) complexing even with the more highly complexing anions fluoride and sulphate [8]. 8. N. S. AI-Niaimi, A. G. Wain and H. A. C. McKay. To be published.
982
N . S . AI-NIAIMI, A. G. WAIN and H. A. C. McKAY
-
1.03
1-03
--
1.02
1.02
--
1.01
1.01
--
1.00
Y
1.00 0
2
4
6
8
I0
12
14 0
16 18 2 0 x
102
Fig. 2. Variation of Y of the Np(V)/Np(VI)-nitrate system in HCIO4 with [NO3]tot.i at CH = 0.8. (1) Yexp.vs. [NO3]total; (2) Y~..... ted (allowing for Ej) vs. [NO3]total.
It is also possible that a second neptunium (VI) chloride or nitrate complex is formed. This possibility is best explored by plotting ( Y - - 1 ) / a against a, as suggested by Leden[9]. If Y=
l + ba + ca 2,
then ( Y - - 1)/a = b + c a , 9. See F . J . C . Rossotti and H. Rossotti, The Determination of Stability Constants, p. 108. McGrawHill, New York (1961).
Stability c o n s t a n t s
983
so that the intercept is equal to b and the slope to c. In the absence of neptunium (V) complexing c can be identified with 6/32. T h e plots, shown in Figs. 3 and 4, show a fair amount of scatter. This is to be expected, since we are c o n c e r n e d with a numerically small ca 2 term, amounting only to a minor correction to the Y-values and having only a small effect on the calculated values of b. We cannot hope to obtain more than a rough estimate of the values of c f r o m our data, even though these values may be of appreciable magnitude. T h e results for C n = 0.4 and 0-5 in the chloride s y s t e m seem consistent with c-values between 0 and 0,4. I f c = 0 we should have constant values o f ( Y - l)/a, while if c = 0-4 we should have a line of the slope indicated. F o r C , = 0.3 the
0-58 0.56 0.54 0-52 0.50 0-48 046 0.44 0-4:
I
,
t
)-48
,
0.46 0.44 i
121 0-42
( Y-I)/=
D-40 0-38 0.36 0.34 ,.32 0-50 0.48 0-46 (3)
044 0,42 0-40 0.38
i
1 I l l 1 2 3 4
I S
I 6
l 7
l I t 8 9 I0 0 x I0 2
~ II
I 12
I 13
J 14
t IS
I 16
17
Fig. 3. Variation of ( Y - l)/a o f the Np(V)/Np(VI)-chloride s y s t e m in HCIO4 with C1-1 at (l) Co = 0.3; (2) Cn = 0.4 and (3) Cn = 0.5 (Y values corrected for liquid j u n c t i o n potentials).
984
N. S. AI-N1AIMI, A. G. W A I N and H. A. C. McKAY
(I)
t
0.14 0.12 :3.10 D.08
0.14 0.12
Y-I)/o
(2)
0.10 0.08
D.14
(Y-I)/o
~3.12
(3)
:3-I0 0-08
0.14 0-12 040 0.0~
2
4
6
8
IO
12 14 16 oxlO 2
18
20 22 24
Fig. 4. Variation of (Y-- 1)/a of the Np(V)/Np(VI)-nitrate system in HCIO4 with [NO3]total at (l) C a = 0.4; (2) CH----0-5; (3) CH = 0"6 and (4) CH -- 0.8 (Y values corrected for liquid junction potentials).
trend of the points suggests a steeper slope, with say c -- 0.8. In the nitrate system c may lie between 0 and 0.1. The effect on the values of 6/31 in both systems will never amount to a reduction of more than - l 0 per cent. The experimental values of b, for the various values of c just indicated, are given in Table 3. The "preferred values" of b correspond to what seem the most likely values ofc. For the chloride system any variation in b with CH lies within the experimental error, but for the nitrate system there seems to be a definite concentration dependence. The only previously reported investigation of the stability constants of the chloride and nitrate complexes of neptunium (V) and neptunium (VI) is that of Cohen et a/.[10], who studied the effect of nitrate and chloride ions on the rate of the isotopic exchange reaction of neptunium (V) and neptunium (VI). They found that nitrate ion appears to have little effect, but chloride ion exerts a marked catalytic effect. They interpreted their results in terms of formation of chloride complexes of neptunium (VI) but not of neptunium (V). Extrapolation of their data to 25°C gives 6/31 - 0.42 mole -11., which is in reasonable agreement with our results. No value has been reported for 6/31for the NpO2NO3 ÷ complex. However, Keder[11] has reported a study of the absorption spectra of neptunium (VI) in 4 M H(CIO4, NOz) solutions at wavelengths of 1230nm. The shifts in the apparent molar extinction coefficients are small and difficult to interpret, but are 10. D. Cohen, J. Sullivan and J. C. Hindman, J. Am. chem. Soc. 77, 4964 (1955). 11. W. E. Keder, J. inorg, nucl. Chem. 16, 138 (1960).
Stability constants
©
985
g....ggg r-,
4-1 +1 +1 +1
E
gggg
©
2
.=. Z6&&&
+~
© Z Z
~-I +1 +1 +E
+ © Z
8
i
+1 +1 lq
+~
~ Z Z 0
+1 +1 +P r,,,
I
+1 +1 +1
986
N . S . AI-NIAIMI, A. G. WAIN and H. A. C. McKAY
compatible with our results. Kirkbright and Yoe [ 12] report no variations of molar extinction coefficients in solutions from 0.16 to 2.7 M HNO~, but they found a decrease of 70 per cent in 6.9 M HNOa and of 6 per cent in a solution of 5 M NaNO3 and 1.9 M HNO3. Their results again indicate that the complexity constants must be small. Rykov and Yakovlev[13], in their paper on the kinetics of the reaction between neptunium (IV) and neptunium (VI) in nitrate solution, state, without giving experimental evidence, that from an analysis of the dependence of the molar extinction coefficient on [NO3] only one complex N p O ~ N O J is formed at [NO3-] < 2 M, with a stability constant of 0.4 mole -11. at 25°C. They then use this value to interpret their kinetic results. Although obviously rather uncertain, this value provides further general confirmation that 6/31is quite small. Acknowledgements-The authors thank Mr. R. A. P. Wiltshire for the a-pulse analysis and Dr. 1. L. Jenkins for useful discussions. 12. G. F. Kirkbright and J. H. Yoe, A nalyt. Chem. 35, 806 (1963). 13. A. G. Rykov and G. N. Yakovlev, Soviet Radiochem. 8, 26 (1966).