Chemistry of plutonium (V)

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Title:
Chemistry of plutonium (V)
Series Title:
United States. Atomic Energy Commission. MDDC ;
Physical Description:
7 p. : ill. ; 27 cm.
Language:
English
Creator:
Kraus, K. A
Moore, G. E ( Gary E. ), 1942-
U.S. Atomic Energy Commission
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Technical Information Division, Oak Ridge Operations
Place of Publication:
Oak Ridge, Tenn
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Subjects / Keywords:
Plutonium -- Isotopes   ( lcsh )
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federal government publication   ( marcgt )
non-fiction   ( marcgt )

Notes

Bibliography:
Includes bibliography references : p. 7.
Statement of Responsibility:
by k. A. Kraus and G. E. Moore.

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University of Florida
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oclc - 277230266
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AA00009293:00001


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MDDC- 906




UNITED STATES ATOMIC ENERGY COMMISSION










CHEMISTRY OF PLUTONIUM (V)

1. Potential of the Plutonium (V)/(VI) Couple
I onic Species of Plutonium(V) in Acidic Solutions
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K. A. Kraus
G. E. Moore


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This document consists of 7 pages.
Date Declassified: April 16, 1947


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; CHEMISTRY OF PLUTONIUM(V)
S:1. Potential of the. Plutonium(V)/(VI) Couple
Ionic Species of Plutonium(V) in Acidic Solutions


By K. A. Kraus and G. E. Moore


AB T rrie

SThrough lotentlomerici trations, the potential of the Pu(V)/(VI) couple in chloride and perchlo-
ri~~e oltidnas was found io be ca. -0.93 v near pH 3 at 25C. Since this value is In reasonably good
agreement with a potential previously determined for an 0.5M HCI solution, and since the hydrolytic
behayor of Pu(VI) insufficiently we.lkaown,.it can be concluded that the Pu(V) species in acidic so-
Lt: upas,; PqQat with an undetermined number of water molecules of hydration. This conclusion was
Sqpn:S4ned through pH measurements.during reduction of Pq(VI) to Pu(V) near pH 3.

1.. TRO WCIOi "

SAs it became apparent that Pu(V) is reasonably stable in solutions of low acidity (even in the
pesebnc of excess of certain -reducizig agents) ad- that Pu(V) solutions can be prepared instantane-
odily by reduction with a number of reagents, iineludlng iodide ions, it appeared feasible to determine
tliepofentfil df the Putr)/(Vt) cbiplb tirdugh'dtiect emf measurements. The potential of this couple
: hiNd previously beeitetifated froni dispt-portknation data By Connick, McVey, and Sheline' for an
0.5M HCI solution, and thus a measurement at considerably lower acidity (e.g., pH 3) coupled with a
S..,d.owle .dgg. te I. plyti. behavior ofpPi(fl wpuld permit identification of the species of Pu(V)
Swich. ta ,s':. lt ons. : .


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).,..p. ,t.ial.,.e Py/(.) ,couple in chloride andperchlorate solutions was determined.
through potentiometric titrationsof Pu(VI) solutions using iodide ions as reducing agents. This meth-
od was chosen rather than the potential determination of a series of solutions for which the ratio of
Pu(V) to Pu(VI) is knowi, since (a) the concentration of Pu(V) cannot be estimated with precision spec-
trophotometrically because of the.low extinction coefficient of its most prominent peak, and since (b)
by the use of this method an independent check could be obtained on the "reversibilitv" of the couple.
The. titration oa Pu(VI) chloride solution was carried out in the standard manner near 25C using
a saturated KCl-calomel electrode as a reference, and using saturated KCI as the liquid junction. In
order to minimize possible Interference through completing of Pu(V) or Pu(VI) by chloride ions, a
very dilute chloride solution was chosen for the experiment.
Several titrations of perchlorate solutions were carried out. Since the results agreed reasonably
.well, onl.the last f these will.be,descrijed here. The titration was carried out in a special vessel,
t Yhic. h :for, temperature, control of .both the solution and the saturated KCl-calomel reference electrode
S was paLrtally immpersd a thermostat whose temperature was 25.0 0.1"C.

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The titration vessel was conical in shape (maximum capacity 10 ml; 3 to 5 ml of solution were
used in the experiment), contained a side-arm for the insertion of a microburette, a small platinum
wire (used as the redox electrode) sealed into the wall of the vessel, and a cap constructed of a stand-
ard tapered female ground glass j;int, through which a small air stirrer could be inserted. Contact
with a salt bridge was made using pa asbestos wick troughtl. t)ottom of the titration vessel. To a-
void precipitation of potassium perchlorate, the bridge was divided by two microstopcocks to permit
the use of three different electroltes in' series 1r.45M NaCIO,;'FG681M NaNO,, and saturated KC1,
the latter being in a ca. 25 cc conicalvessel Into which'a itandard;(Beckman) calomel electrode was
placed. It was found necessary to permit thermal equilibration of the calomel cell for several days to
obtain reproducible readings.
Both before and after the titration of the perchlorate solution, the calomel electrode was cali-
brated against a quinhydrone electrode. For this purpose quinhydrone was dissolved in 0.09921M HC1.
Assuming the standard quinhydrone half-cell potential to be 0.6990 volts"' and assuming t.p eaft-
ity coefficient of oxonium ions to be 0.841 for 0.1M HCI,4 the calomel half-cell potential was found to
be 0.25430.0003 volts, The difference of this value from that usuallyassumed (-0.4246 .1 s)' is
probably due to the inclusion of the arithmetical sum of the various,liquid junction poteatialp ,of
system .*. i,'.i .,
The potential measurements were carried out with a Beckman Model G pH meter (cMtdfkiou-
tion) and a Leeds and Northrup Type K potentiometer (perchlorate solution)'.' Some dffiufiW4i ~Ire
encountered in finding stable potential readings after the addition of each new aliquot O-'s 4 TAf lbdide
solution. The potentials drifted for many minutes, and to get the "final" reading it was necessary
either to wait ca. one-half hour to one hour to get constancy, i.e., drift less thawBA8 Siraiay*,e pr
to extrapolate a series of readings. Probably only a small part of this drift was due to poor stirring
of the solution. Instead, the drift suggests that either the Pu(V)/(VI) couple is reversible ani to a
limited extent, or, more probably, that thel measuring electrode did not behave perfectly., I 4w.'t.,
reasonably good agreement of the experimental. points with a line of.theoreticl.. slpe, (pee gai" .)
it can be concluded, however, that the drift in the readings did not seriously affect tih finple uts.
In this connection it ia interesting that a similar.difficulty was later observed In.the detpzwlaihaof
the potential of the Np(V)/(VI) couple ;.,:. p .
Analysis for the various oxidation states of plutonium was carried' out spectroplot6mditrtcilfwith
a Beckman quartz Model DU spectrophotometer. Using the extinction coeffictiets given Isiathei
paper,' the calculations were carried out according to the "3-point ratio" method previously de-
scribed.' The plutonium concentrations were determined through standard radiochemlIm alfpiAassay.
The Pu(VI) solutions were prepared from Pu(VI) nitrate solutions through two "barium (poly)plu-
tonate" cycles,' consisting of precipitation with excess barium hydroxide and dissoliititlinr a small
amount of the requisite acid. The composition of the solutions before titratiod was
a) Chloride solution: 6.25 x 10-'M Pu, pH 3.18, 0.002M C1, L* = 0.01 *.
Sb) Perchlorate solution" 1 65 '10'-M Pu, pH 3.14, 0.039M AC10, T = 0.! 1.
When the iodide titration was ca. 50% complete, the ionic strength of the Pu(VI) perchlorate solu-
tion was increased to 1.0 through addition of concentrated sodium percMlorate, to determine'tWe differ-
ence in the potentials at the two ionic strengths.
Solutions (a) and (b) contained 96% and 6%l Pu(VI) and ca. 3% and ca. 26% Pu(V), reqpetively,
as determined through a combination of the iodimetric and spectrophotometric data. (The end-point of
the potentiometric titration could be Identified with reasonably high precision since a large potential
change occurs there. This suggests that an iodimetric titration of Pu(VI) to Piu(V) near OiB i could be
used in the quintitative analysis of plutonium. The method appears to be limited only by the' purity of
the original Pu(VI) solution with respect to other oxidation states of plutonium, since as Ahown In

j represents ionic strength.




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MDDC .- 906 [3


separate experiments, the presence of excess iodide jons does not cause reduction to Pu(IV) or
Pu(Il) near that pH.) Solution (b) also contained several per cent polymeric Pu(IV).
Using a glass electrode and a Beckman Model G pH meter, the pH of the solutions was measured
before and after the titration.

3. RESULTS AND DISCUSSION

The results of the potential measurements have been summarized in Figures 1 and 2, where the
observed potenials have been plotted against
logt, (determined from the iodimetric data).
A line of slope -0.0591 was drawn through the experimental points, since they should fall along
such a line according to the equation
E = E' 0.0591 lo 0.0591 log + (1)
where E = obsetred cell potential
E = standard cell potential (without liquid junction)
a = activity coefficient subscriptt denotes species)
e = sum of liquid junction potentials
Parentheses indicate concentrations in moles/liter.
The agreement and thus the precision of the measurements appear to be satisfactory, especially for
thh perchlorate solution (Figre 2), where sufficient time was taken to obtain equilibrium readings,
indicating that the drifts in potential reading did not seriously interfere.
The cel potentials for unit ratios of Pu(VI) to Pu(V) where 0.687 volts (chloride solutions),
0.6808 volts (perchlorate solution, I = 0.11), and 0.6848 volts (perchlorate solution, 4 = 1.0). From
these the formal half-cell potential of the Pu(V)/(VI) couple for the chloride solution can readily
be evaluated by assuming zero liquid junction potential (at the saturated KCI bridge), and -0.246 volts
for the standard potential of the calomel half-cell.' The value thus obtained for the Pu(V),'(VI) po-
tential is -0.933 volts. The uncertainty of this value has been arbitrarily set at0.005 volts, since
the calomel eletbfode was not separately calibrated and differences of this order of magnitude can be
expected. The assumption that the liquid junction potential was zero is probably reasonably good.'0
Evaluation If the formal potential for the Pu(V)/(VI) couple in perchlorate solution is consider-
ably more difficult because of the uncertainty connected with the estimation of the various liquid
junction potentials of the system Of these, the liquid junction potentials at the junction 1.45M
NaCIO1 0.0992M litl (satiirated with quinhydrone) Is both the most difficult to evaluate and probably
also the largest. The assumption that the sum of the liquid junction potentials and the potential of the
calomel half-cell is -0.25430.0003 volts is applicable only for the standardization measurement,
since the liquid junction potential would be different during the measurements on the plutonium-
perchlorate solution. Since this change of the liquid junction potentials cannot be readily evaluated,
it will be assumed that iJ remains constant, although this will then introduce a large uncertainty (ar-
bitratily set at0.015 volts) for the value of the Pu(V)/(VI) potential in perchlorate solution. Using
-0.254 volts as the reference potential, the value -0.935 volts for the potential of the Pu(V)/(VI)
couple at ionic strength j = 0.11 can be calculated;


*Parentheses indicate concentration in moles/liter.
I: "Formal" potentials refer to unit ratios of the concentrations of the oxidized and reduced
species, rather than unit activities of all species-'ts is the case with "standard" potentials.
SFor the reaction Pu(V) -4 Pu(VI) + E-.

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0.690


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0 -0.60 -O40
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Figure 1. Observed cell potentials kor chlori t ~satioa.
o Observed potential
Solid line has theoretical slope q1 0.001 volts


The half-cell potential at ionic strength L = 1' appears to be practically the same since the small
difference in the obseaged cell potentials lies within the expected change of the liquid junctim poten-
tials, and thus no further significance can be attached to it. .


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MDDC 906


The agreement of the potentials in the chloride and perchlorate solutions is surprising in view of
the unicetainties connected with the measurements of each potential. The data show, however, that it
:'is entirely'unlIkely tht tinder the conditions of the experiment chloride ions complex Pu(V). This is
id agreement with later experiments.' (At the time these experimentA were carried out, chloride
complexes u f Pu(V) could not' entirely be excluded, even for the highly dilute chloride solutions used,


0.740


0.330



0.720


710 -


w0o700
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M(Pu(VI))
M (Pu (V))


0.00 +0.20 +0.40 40.60


Figure 2. Observed cell potentials for perchlorate solutions.
o Observed potentials M = 0.11
A Observed potentials p = 1.0.
Lines A and B have theoretical slope of 0.0591 volts


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MDDC 906


since very little.,as known regarding the chemical properties of Pu(VJ. irom qur pqiesent aowLpdge
of its chemistry, however, i. appears very unlikely that complegxing .,QLPuI ) wyuld tke.place: evenRin
concentrated chloride solutionm, More experiments are necessary to;prove this point co P.lrstiely,
Chloride completing of J u(VI) wou.d be expected to,occur oWly!at very much,higher concentaaio"s,
of chloride ions.')
Comparison of the potential data determined near pH 3 with those determined at higher acidity
(0.5M HCI) is of interest since from this comparison identification of the Pu(V) species in acidic so-
lutions is possible. The value reported by Connick et al.' for 0.5M HCI is -0.912 volts. Some com-
plexing of Pu(VI) by chloride ions is known to occur in 0.5M HCi" and therefore a somewhat more
negative potential would be predicted for the uncomplexed species. Since the stability content of the
chloride-Pu(VI) complex is unknown, the difference in the potentials can only be estimated, and prob-
ably amounts to ca. 10 to 20 my. This would make the Pu(V)/(VI) potential in 0.5M HC1, after cor-
rection for chloride complex ing, ca. -0.92 to -0.93 volts- in reasonably good agreement with the
value determined here at pH ca. 3-showing that the potential of the couple is practically pH inde-
pendent in this range of acidities.
It is generally assumed that Pu(VI) in acidic solutions is PuO2,4+, and it was shown that prima-
rily the same species is still present near pH 3." Since the constancy of the potential shows that
hydroxide ions are neither liberated nor used up in the reduction of Pu(VI) to Pu(V), one can conclude
that Pu(V) in these solutions is symbolically represented by PuO,+ with an undetermined number of
water molecules of hydration. This conclusion is in accord with that based on disproportionation
studies."
The conclusion [hat the ormula PuO2 represents the ionic species of Pu(V) in acidic solutions
is also confirmed by 'he results of the pH measurements before and after reduction, from which the
number of hydroxide ions (or oxonium ions) liberated in the reaction Pu(VI) + 1- Pu(V) f 1/2 I, can be
deduced. Since Pu(VI) near pH 3 is primarily PuO,+, and since hydroxide ions are neither liberated
nor used up in the oxidation of I- to 1,, all pH changes reflect the difference in the number of hydrox-
ide ions (or oxide iors) of PuO,*+ and the P.(V) species in the solutions.
The pH of the chloride solution changed from 3.18 to 3.25 during the titration. Since this pH
change is very small compared with that expected on liberation of one hydroxide ion per Pu(VI) ion
reduced, it can be concluded that the number of hydroxide ions on Pu(V) and Pu(VI) is the same, i.e.,
that for Pu(V) the formula is PuO,*. The slight increase in pH cannot be due to dilution of the solu-
tion, but probably is due to the fact that the Pu(VII solution previously had been at a high pH and that
because of the hysteresis effect of Pu(VI) a small amount of this species was still hydrlyzed." The
pH of the perchlorate solution changed from 3.14 to 3.31. The bulk of this changed can be accounted
for by dilution resulting from the addition of the concentrated sodium perchlorate solution during the
experiment. The rest of the change is probably primarily due to the change in ionic strength of the
solution aind the resulting change of the activity coefficients of the oxonium ions. From both of these
experiments, therefore, as w. 11 as from number of similar experiments, the cor( lusion ju I
reached, that the ionic species of Pu(V) in acidic solution is PuO,', is we! *'uh-ta:a'.:atcd.










'This potential was originally given as 0.92 volts by Connick, McVey and Sheline, and was
apparently re-evaluated by Connick," whose value i- used here.










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t iit" c K. kelmel:y: i" The *5 oxidat a ion, state of plutonium, CN-191.2,


.,. 1." e..Prin.. lpLsd of Eaperimental and Theoretical Electrochemistry, First Edition,

a; EditinsM-'rw 39:5 7 '194 3
4i. P. .i -iAO Ja y .1. (.. d,).
4. sc .hr .r G. i.. .Che. .s. 47:198 (18.).

15. Latinier, W. E., ~ kidation Petentials, Pre~ntieea1l Inc., New York, 1938.
6. MagEussEi, L. B., 1. C. Hindman, mad T. I. LaChapelle, Formal oxidation potentials for nep-
Stuniun couples nd stability of oxidation sttes in solution, CN-2767, March 1945.
;7. Gevaatman, L. II., K. A.; iaus, Chemistry of Pu(V), Mon- ce-80, May 27, 1946.

8. Kraus, K. A., Studios on ply oeric Pn(IV); depolymerilation of polymeric Pu(nI) solution in
nitric acid, CN-3309, June 30, 1945.

7. 9. Moore, G. S.,.K. A. Kraus, Preparation of Pu(VI) solutions containing various negative ions,
CL-P-432, July 20, 1945.

10. Daniels, 1., M. h. Matches, and I. W. Williams, Experimental Physical Chemistry, Third
.. .: Edition, McGraw-Hill, Inc., New York, 1941i-

11. Kasha, M. and G. ;. Sheline, Ionic equilibria and reaction kinetics of aqueous solutions of
plutonium, Ci-iSC January 17, 1943.

12. Connick, R. E., Oxidation states, potentials, equilibria, and oxidation-reduction reactions,
Plutonium Progress Record Volume 14a, Chapter 3.

13. Kraus, IL A., Oxidation-reduction potentials of plutonium couples as a function pH, CN-2830,
January. 23, 1947.

'14. Krau,. K. A., and J. R. Dam, Hydrolytic behavior of plutonium(VI) CN-2831, October 14, 1946.







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