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Subject Category: CBBGSTHf^aBBBRAL

4 /0 - . 35340

* . ';: #Ms< gBBfcOBl IliLllliN COMPLEXES • 'SI, John T. Barr and Charles A. Horton

Work Supervised by R. H. Lafferty, Jr. Laboratory Diriilon F. V. Hurd, Superintendent

CUSSIFIMTON c w a m

Sr ATOMIC *■•?«Snrt.i-.wt QW «.oH

—*Q " \uiur. nwOMSirioATION HRAMCS F *-.

CARBIDE ABB CARBOH CHSKXCALB CtMPAI* K-25 P lan t Oak Bldge, Tennessee

RESTRICTS) This document contains rpsgSCS^^atB ae defined in the Atonic Energy Act traneaittal or the disclosure of ita dSSteSta in any manner to an unauthorised person is prohibited.

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***• v* * ? iiiii'n'i'iiii'iV i 'r Date of leant: leant: of Date K-705 Sodwr: Report Bight chelated complexes and 7 uranium am ines, prepared fraft the the fraft prepared ines, am uranium 7 and complexes uranium chelated Bight Salicylic and la c tic acids, acetylacetone, and sim ilar ccmplaxing agents agents ccmplaxing ilar sim and acetylacetone, acids, and tic c agents la complexlng and with Salicylic ealta uranium of aolutlone organic of reaction pleting agent, the formation of a Werner-type Caspian resu lted . For For . lted com­resu Caspian compound the and Werner-type uraniua a of the with of formation reaction the solution e l the agent, t t i to l shoved pleting base group*, external an functional acidic only characterised. and contain which prepared been have respectively, bases, organic example, uranyl n itra te hexahydrate, sa lic y lic a d d , and pyridine in in pyridine and , d d a lic y lic sa hexahydrate, te of itra n addition uranyl example, the upon However, solutions. organic compounds In uraniua n itra te . Uranium(IV) chloride, la c tic a d d and pyridine formed dichloro- dichloro- formed pyridine and d d a aallcylatoaquopyridlnouranyl tic c • of la chloride, Uranium(IV) precipitate a dllactatodipyridinouraaiua(IV) gave . te solution itra n acetate butyl -hnllde n qiadnc wih oti bt alo and anlno both contain which , d d a quinaldlnlc and I-Phenylglydne lunllaolqornl UainI) hoie ece ih B- with reacted chloride Uraniun(IV) dlqulnaldlnatodlaquouranyl. Pyridine precipitated aquopyridinoumayl n itrate from organic solutions solutions organic from itrate n aquopyridinoumayl precipitated Pyridine pbenylglyclne in butyl acetate solution to precipitate dlchloroblt - - dlchloroblt precipitate precipitated to hexahydrate ts solution itra n acetate uranyl butyl organic in compounds In acid, pbenylglyclne uraniua with quinaldlnlc With readily reacted solution. groups, carboxylic from aqueous solution gave bls(R*phenylglyclnato)tetraaquouranlwi(n) bls(R*phenylglyclnato)tetraaquouranlwi(n) gave solution aqueous from The infrared spectra of 10 of the complexes were determined in the 2 to to 2 the in determined were complexes the of 10 of spectra infrared The f rnl tat eayrt, n tlyliolrnu I) hoie from chloride (If) trlpyrldinodluranium and hexahydrate, te itra n uranyl of chloride. on drcl t h uaimao ya codnt-oaet linkage. coordinate-covalent a by atom uranium the to directly bound bopin ad at .6 50, *0 5*6 ad *31wr msig le ... llev missing, were 5»*t6> 5*93/1 and 5*20, 5*02, 2.76, t a bands absorption RpeygylaoHpeygylornu(7, hc uo rcysaltton n stallitatlo recry upon which (R"phenylglyclnato)H-pbenylglyclnouraniua(I7), 15 solutions of uranium(IY) chloride. Aniline formed diaaillnoureniun(IY) diaaillnoureniun(IY) formed Aniline chloride. uraniua(IV) with chloride. uranium(IY) chloride of solutions absorption bands in the 11.2-12.0 and 15.0-13.2 15.0-13.2 and 11.2-12.0 the in bands absorption u t h R0- lnae ' RC00-U linkage. the to due In a ll complexes which contained pyridine, the characteristic pyridine pyridine characteristic the pyridine, contained which complexes ll a In before chelation could occur. Step-wise replacement of inorganic anions anions inorganic a by of attack replacement necessary was Step-wise preliminary a molecule occur. could un-lonlttd solutions, an of organic chelation stem in uranium before the t, a upon group th basic postulated was t I by complexlng agents could then proceed u n til the maximal coordination coordination maximal the attained. was til atom n u uraniua proceed the then number of could agents complexlng by mco rne fo hc i ws on ta te mn ntoe aos were atoms nitrogen amino the that found was it which from range, micron , 1951 7, y m ABSTRACT te: io-O URAHUM Uimon-TOB COttUDOS : itle T Subject Category: Category: Subject uhr J. . r ad A Horton A. . c and irr B T. . J Author:

l j

regions appeared to be to appeared regions cmCBm-CBBORAL

i * *

mmsi-vm mmm mm*m

Darlas the court* of a recent — arch for specific chelating astute for u m vith solution* of urtniun caapousds In organic solvent* (5), It ate found that the — type conplejdng agsnte which have been reported to be useful for the detection of uranium in afueoua solution* reacted with uraniun compounds in organic solvent** It wte Observed, however, that non/ of the product* which were for— i In noo-**— ous eolation differed greatly fron thoee foe— d in a*— ou* — die* Opon farther investigation, It we* found that the— product* were Vernar-typt cunple— s, rather than the in— r-ccnple* type of ccnpound which la fox— d to readily by ureal— .

theoretical Disc— loa

The concept of secondary, or coordinate, valence w— introduced by War— r to account for the fox— tloa of — tal and— compounds la which — dn or or— ale — 1— • — re thought to be bound to a — talllc ion by a coordinate- covalent sharing of — electron pair furnished by the nitrogen etc*. This theory was developed fron a study of the anlnocobaltlchlorld*s, who— structures are eh— 11 below*

♦ ♦ h 3 % ✓ "h3 ♦ MJN\ ✓ " Hs hjN*♦ CO •*-MKs MjK ”♦ CO ♦'Cl ter . „/• V.. V...

Eenanlnocobalt(XIZ) Chloride Chloropsnto— inocobalt(nz) Chloride ""1 + T 5 * /Cl Co MJ*^ f ^ Cl ^ n h 3 . id ■

Dlchlorotet— dnocobalt(xn) Chloride

It woe later found that other groups canid fans similar compounds. 8c— of the acre familiar of the— ore shown below.

•*“■**•# t v. g| .-mm. ***«**-"■' W**»*MIW Wr -4Mt« WWW 7

«■» Ml S ” -u Ms*\ /c« Cu Pt ccr CM''* Ns* Cl'* ^Cl

Potassium Tetracyaneto- Trlrhlnm— laoyUtlnai(I?) ccpperata(II) Chloride

T 1 “ i :

H C 8 n ^ ^SCM Cl^ ^C« 2No HO HC8^ ^MCM

m Ml

Sodiiaa Tetraisothlocyanato- J)ichlorodiaminoplatim»( 1Z) mercuryate(Il)

These Verner-type, or penetration, complexes contain inorganic anions, and art characterised by the fact that one or both of the Ions are complex, being formed by coordinate-coTalant bonds* Many solvated aolecules may be Included in this class. With the exception of a few neutral compounds similar to dlchlorodlsimni noplat lnum( 11), these complexes are more soluble in vater than in non-polar solvents, have high melting points, and behave as typical electrolytes in solution.

This idea vas later developed by Bilge!ck and Pauling to explain the formation of chelate compounds, "oniua" ions and hydrogen bonds, and is a fundamental part of the Levis acid-base theory.

According to these later-workers, such bonds occur by the interaction of an electron-donor (electrophobic groups such aa trlvalent nitrogen, carbonyl and hydroxyl oxygen, and sulfide sulfur) vith an acceptor group. The acceptance of a pair of electrons then results in the completion or expansion of an electron shell in the acceptor, or in the formation of a nev electron grouping*

The number of donor groups which msy associate themselves vith an acceptor le a function of the spatial requirements of the donor groups, and of the amount of expansion possible in the electronic configuration of the acceptor. For moat metallic lone this number le a characteristic constant knovn as the coordination number (C. I.), and varies from 2 to 8 as the' ionic radius of the metal increases. * a The coordinating groups are arranged in a definite epatial arrangeaent in the voluae around the central ton. This voluat la known as the coordination »ph*r«.

Polyfunctlocal organic ccaplexlng agents any fora 5- or 6-Mehbered ring! with Metallic Iona If the functional groups art in the proper positions. These complexes are known at chelate coapounds. Xxaaplta with ethylene* dlaalns and oxalic add are given below. These reagent* are bldentate (two points of attachaent) coeplexing agent*.

♦ CHt - W*“ CM. / o*c*o ter »r Pt

av„_ , " - CHf 0»C-0"* \>-C»0

lie (ethyleaedt ail nn) plat lnua( II) Potassiua Dioxalatoplatinate(n) Chloride

If one coordinate valance and one ionic valance are used to font the ring structure, the resulting compound is terand an Inner coaplex. If the coordination maker of the Metallic Ion it twice the ionic charge, and a bldentate reagent is used, the resulting compound le an inner ccaplex of the first order. These compounds are non-electrolytes and possess ■any covalent properties, such as low Melting points, Halted volatility, high in organic solvents, and low solubility in water. If the coordination maber of the ion le greater than twice the ionic charge, additional coordinating groups any attach theMaelvea to the central loo to coMplete the coordination maber. Bxaaplee of these are given below.

OeC-0^ yO-CeO „ o r \ ♦ Pt H S » W \

0 U Dlglyclnatoplatlnua(Xl) DiacetylacetonoAlpyridinoiron( XI) C.I. « twice ton charge C.H. greater than twice ion charge Hletorlcal Bedew

A few uraolua compounds have been described in the literature which nay have been Werner-type coaplexes; however, this structure has not beun definitely assigned to any previously reported coepound. Faecal (25) reported the exists nee of several cyanouranatee and Hauser (lk) and Xaghilleri and Qorl (16) prepared a mabar of oxalato coapounde which

r».« i 9 may have bid complex-loo structure. P i m i d i (6,9,10,11) reported a m i tel of uranyl compounds vith ortho-di phenols which be termed beteropoly aclde, but vhich also may be considered u Vernsr-type complexes. Andrew (9) peppered a complex naLatouranyl salt, but the ionic charge of the complex anion uni due to aa uncoordinated carboxyl group rather than to a reiidual charge of a coordinated function. This it probably true of may of the complex ureayl anion* found in aqueous solutions of the polyfunctional aliphatic aclde, tuch as citric, tartaric, and ascorbic aclde (3b), at ordinary hydrogen ion concentrations.

Much non work has been done on the uranium amines. Rasc&nu reviewed the literature up to 1930 (&7)» and described the preparation and properties of mny more conpoundi (98-31), Including such pyrldlno compounds as •quopyridino- and dlpyrldlnouranyl chloride and nitrate, and the antipyrlM, pbenacltln, and ’oyrlmidon analogs • Vllkle-Dorfurt and Schllepbake (b2) prepared the pentaaatlpyrlnouranyl perchlorate, and Inghlllerl and Oori (16,17) reported a serlee of anllino end quiallino salts, although their Interpretation of the structure• were obviously wrong in several instances. Lloyd and Cities (21) described an addition compound between hexmethylenetetrmlne and . Koatlgnle (23) prepared a series of amines frcn heterocyclic anines.

Spacu (36) described a series of amonlatee of uranyl and uranlua(IV) chlorides, nost of which vere stable only at low texperstures. Bchlesiager (37) reported conpounds of uranium (IY) chloride vith two sole exiles each of pyridine or trlethylanlne. Allison and Mann (1) prepared coordination conpounds of uranlun(77) chloride and bromide vith trlalkylphosphine■ and arelnsa. Rena reported (32), mistakenly, that uranium(IV) chloride dissolved in acetone to fora a yellow solution, and gave a voluminous precipitate upon the addition of pyridine or quinoline. Lafforty, Schuman, Radiner, and Salley (20) reduced uranl\sa(Tl) fluoride vith ethaaolaaine to obtain a tetrsms(ethanolemlno)uraniua(IV) fluoride vhich it probably beat explained by & Werner-type structure.

It Is beyond the scope of this report to review the inner complex compounds of uranium. Croxton has prepared a complete literature survey for uranium (8). Ware (39) has prepared a literature survey of organic analytical reagents for uranltaa reported before 19b5. Seville (2b) has given an excellent aunmary of the spatial requirements and functional groups involved in uranium chelate formation. Several hundred Inner complex compounds vere prepared by Project workers (12,13,33,58,39,41), any many more ccmplexlng agents vere screened for reaction vith uranium compounds in aqueous solution (18,22,2b,bo,43).

Rasc&nu prepared many of the uranium amines in eayl alcohol solution (27,28). Almost without exception, all others Investigators vorked vith aqueous media. The observations made during a recent study of the reactions of uranium compounds in organic solvents (3) led to a more complete investigation of the reactions of uranium compounds vith cam­ ple xlng agents in non-aqueous solutions. i| • |M|||MM«*• ■

10 nPtROOBTAL

Materials and Apparatus

The source of the uranium compounds has been given in a previous report (5) • tbs organic reagents used vers tho ccmmmrc tally available, reagent-grade materials. Solvents were of reagent grads, and were dried mad redistilled before nee.

Absorption spectre in the visible region were determined on a Beckman Model B Spectrophotometer, using a 1.000 ca. Corex cell and a blue- sens it ire photocell. Molecular weights were determined in a Menxles- Vright apparatus, using absolute ethanol ns the solvent. i Butyl acetate mas chosen for the solvent because of its good solubility characteristics, convenient volatility, and because it did not mask the reactions of the chelating agents. Ethyl acetate, dlaxane, msthylethyl ketone, and dibutyl carbltol were also found to be suitable solvents for this work, but strongly coordinating solvents such as dimsthylforn— ids and tributyl phosphate exerted strong masking action.

Pyridine was chosen as the base to he used in this study because hetero­ cyclic amines have been shown to have stronger coordinating tendencies with uranium compounds than do aliphatic amines (27), and because the pro­ ducts were more highly crystalline then those formed with other type baste.

Carbon and hydrogen mlcroaaalyses were obtained by the conventional combustion method. The residue from this combustion was weighed as urano-uranic oxide (U,0g) in order to obtain the uranium analyses. Titration with Burl Fleeter reagent was used to determine the water content of the ccmplemis.

Ths infrared spectra of 10 of the complexes were determined In tbs 2-15 fx region) using a Ferkln-Blmer double-beam recording instrument equipped with sodium chloride cells. The solid sample• were prepared by mulling in Bujol.

These spectra are presented with ths preparation of their respective com­ plexes. Wherever possible, the best assignment of the type of vibration v responsible for the bands is given In parentheses following the value of \ the absorption maxims of the bands which were present in the ligands or the uranlua compounds and dlsappemed in ths complex. The new bands which appeared in the complexes are also listed.

For purposes of comparison, ths spectra of uranium(IY) chloride, uranyl fluoride, hexahydrate, lactic acid, qulnaldlnlc acid, and dlnethylethanolmmlne were obtained. These spectra are presented in figures 1 and 2. The spectra of pyridine, salicylic acid, and H- phenyXglyclne were taken from the data given by Randall (25). Figures 1, 2, 4, 5, 7p 8, 9, and 14 show breaks In the 5«5 and 7*0 fx regions of the spectra due to lots of seneltlvlty by the instrument in regions of high absorption, which resulted in overcocpensatlon for the lfujol bands. %9T *» t t 4 6 9 r 6 9 4 3 NRRD BOPIN PCR SPECTRA ABSORPTION INFRARED URANYL NITRATE HEXAHYDRATE F RNU COMPOUNDS URANIUM OF i i I r I i i I i i FIGURE I 9

0 I t I 4 15 14 II It II 10 * + wV r

O W N A LO UV C A G IO 1 1 i 1 10 II It IS 14 IB

INFRARED ABSORPTION SPECTRA OF COMPLEXlKlG AGENTS

FI4U4I t

§

**4% --

■* Preparative Procedure

The Method used for the preparation of cCMplexas requiring the addition of an external bate mu toilii vara, filtered aolutloa of the ureal ue compound In butyl acetate and a eolation containing a two-fold exeeee of the chelating agent. Slightly aore than the theoretical amount of baee vae then added dropvise with rapid stirring. The .eolation waa than allowed to stand for a abort tine and vat filtered. Precipitation vae usually instantaneous and quantitative.

The anal nee were prepared by a tlallar addition of an organic baee to the uranlun salt solution. The sans Method was used to prepare complexes frca aaino acids and uranlun salts.

In nost cases, the ccmplexsi could not be purified by conventional recrystallitation Methods, either because of very low solubility, or because of s tendency toward oxidation or hydrolysis. After filtration of the Initial precipitates, these conpounds were subjected to several successive digestions with ethyl ether, nethyl or ethyl alcohol, and, usually, acetone, and were ♦ben dried by extended evacuation at roon temperature,

Dlnethylforaeaide showed strong solvent power for aany of the canplexes, but the complex could not be recovered from the solution, Addition of ether gave oils, and evaporation Induced ds composition.

R1ACTIQW WITH OROAHIC BASKS

Reactions with Pyridine

Uranyl Hltrate Hexahydrate. Tho addition of pyridine to organic solutions of many! nitrate hexahydvate resulted in iMMsdlste and quantitative precipitation of the uranium In the form of aquopyrldlnouranyl nitrate. In very dilute solutions the precipitation became slower, and did not form at all In polar or strongly coordinating solvents such as ethanol and trlbutylphosphate. The Halting concentration for lansdlate pre­ cipitation In several typical solvents has been reported (5),

It was not found necessary to use the dehydration technique described by Rasc&nu. This technique Involved the distillation of a portion of the solvent fron solutions of hydrated urrayl compounds. Warning or boiling the solution after the addition of the pyridine often hastened complete precipitation of tbs uranlun coupler, but in nost cases the products were the sane, regardless of the water cogent of the uranium compound used. For example, both uranyl nitrate hexahydrate and uranyl nitrate, dlhydrate formed aquopyrldlnouranyl nitrate upon addition of pyridine to their solutions. Rscrystalllcatlon of this monoasMln? frca pyridine gave the dlpyridlnouraayl nitrate. These compounds here been vell- characterited Ly Rascanu (27;20) • However, none of the explosive characteristics described by kin for the nitrates were noticed in this work. The infrared ape etna of dipyridinouranyl nitrate la given in figure 3« The pyridine bands which vanished in the ccopier were: 2.76, 3*02, 5.46, 5.93 (OH stretching), 6.33, 8.21, 9.36, and 10.09 (OC ring stretching) /t vibrations. Ho bands observed in uranyl nitrate hen- hydrate were absent. Hew bands appeared at 7*72, 7*90, 8.05, 8.10, 8.4, 8.6, 10.8, 11.2, 11.4, 11.9, and 12.3/1.

2 3 4 5 6 7 8 9 10 II 12 13 14 15

Figure 3 Infrared Absorption 8peetnas of U02(I0j )2*2C^H

Uranlus(lV) Chloride. A solution of 5 g. of uranium(lV) chloride in 50 nl. of butyl acetate was heated to 90 *C., and 2.0 ml. of pyridine was added to effect complete precipitation of tripyridinodluraaiuB(IV) chloride. The greenish-white precipitate was filtered and extracted several tinea with boiling ether. Tfce product weighed 6.6 g. after vacuum drying, and was Insoluble in methanol, ethanol, ether, dlonne, ethyl acetate, and benzene, slightly soluble in pyridine, and soluble in dimethylformamide, from which it was re precipitated as an oil by the addition of ether. Water decomposed the solid, with the formation of uranium(IV) oxide. The melting point was 165*C., accompanied by de­ composition.

AnalyBls: (C^J) jfUCl^g

Calculated C, 18.06* H, 1.52* 0, 47.75* Found C, 18.08* H, 1.89* 0, 48.40*

The composition of this compound was Independent of the concentration of the reactants, the order of mixing, or the solvent. This quanti­ tative precipitation was made with equal ease from a variety of solvents with a wide range of polarity sal coordination power, such as acetone, alcohol, tributyl phosphate, and dibutyl carbitol. .

The infrared spectrum of the compound prepared above is given in figure 4. The following bands which normally appear in the pyridine spectrua vanished in the complex: 2.76, 3*30, 6.33, 8.21, 9*36, 10.09 (OC ring stretching) Bind perhaps the 5*02, 3*20, 5*^6, and 5*93 (OH stret vibrations. The latter group, which are weak m M M M n i N

absorbers in pyridine, were doubtful since the complex hid a broad weak absorption in the *.25 to 6.0511 region. Rev bands which appeared were at: 2.85, 3.15, *.15, 7.29, 7.6>, 8.05, 8.10, 8.*0, 9.55, 10.85, 10.95, aad 12.5/1*

100

c t > \ S 8

2 3 4 5 6 7 6 9 10 II 12 IS' 14 IS

Figure * Infrared Absorption S pectra of 2I0C1^*3C^H^B

Reactions with Dlaethylethanolsml ne

Amino alcohols reacted similarly to other amines, with the exception that the presence of the second coordination group within the molecule alloved the replacement of the molecule of water which was present in the normal uranyl monoamines.

Uranyl Rltrate HexahydLrate. The addition of 5 ml. of dimethylsthanolamlne to a solution of l6 g. of uranyl nitrate hexahydrate in 50 ml. of butyl acetate gave an ineediate precipitate. The crude product vms digested with 50 ml. of ethyl ether, then 10 ml. of methanol, and finally, with two more 50 ml. portions of ethyl ether, and dried under racum. The yield was 7.5 g. of a red-orange, microcrystalline solid. This compound was slightly soluble in methyl and ethyl alcohol, and pyridine, and vas decomposed In water solution. It was Insoluble in all other solves teeted. '

AB.ly.iB! UOg(BOj)g* (CHjJgECHgCHgOH

Calculated C, H, 2.29)1 U, 1*9.27$ round C, 10.17$ a, 2.5$$ 0, >*9.75$ The structure of thie compound is probably:

h2c Vi*, I u i t # # *

Q i * * ***'■■»*.! ■ k i »— ------— —p ■ M M ..ii1iiiWKHLN,MMIN!R» •—... .

The Infrared absorption spectral is given in figure 5 <

Beads of dlrathylethanol —ins missing in the complex rare: 3.10 (HH stretching), 3.65 (CH* stretching), 5.85 (OO stretching), 7*90, 8.07, 8.20 (tertiary salr & ), 0.50 (asma), 8.70 (saae), 9*10 (seas), 9.25 (earns), 9.6$ (CH, rocking), 9.8, 11.2, ll.k5 (C-0 stretching), 12.9 (HHg defomationror CH rocking) jx vibrations. The 3.60, 9*75, and 11.77* uranyl nitrate hexahydrate bands rare also absent. A new band at f*08m ray be attributed to a quaternary nitrogen, and a nsv band a t 10M J4 to a CH^-041 linkage. Other nsv bands rare located at 7.25, 9**5, 9*92, 10.157 11.05, 12.07, 12.25, and lt.15 j*.

a

Figure 5 Infrared Absorption Spectrum of UOgfKOjJg^CHjJgHCHgCHgOH

Oraniua(rv) Chloride. With uraniua(IV) chloride an intractable black tar v&s obtained, from vhich no pure caspound could be isolated.

Reactions with Other Bases

The dark-green precipitate obtained from uranlum(X7) chloride and quinoline had a structure analogous to that of the compound obtained • t. from pyridine and uranlua(lV) chloride, but the greenish-tan solid obtained by the addition of anillra to solutions of uranlua(IV) chloride had the composition of oj^linouranium(TF) chloride.

Analysis: UCl^^^HHg

CalculatedCalcvi C, 25 H, 2.b9* B, k.95* Found C, 25.8# H, 2.72* », 5*2ty 17

RBACTI0H3 WITH ACIDIC CdffUBCHO AfflBfTB

Salicylic Acid

As vas the case vlth moat chelating agents with purely acidic functional salicylic acid alone shoved no reaction vlth either uranium(IV) chloride or uranyl nitrate hexahydrate in organic solvents. The addition of an external organic base censed Immediate reaction.

Uranyl Mltrate Hexahydrate. To a solution of 10 g. of uranyl nitrate hexahydrate and l6 g. of salicylic add In 100 al. of van butyl acetate, 4 ml. of pyridine vas added dropvise. After standing at rooa tempera- ture for a short tins all the uranlua had pradpltatedi leaving a nearly colorless supernote. The yield vas 11.0 g. of a dark-rod, crystalline salt, salicylatoaquopyridinouranyl nitrate. Recrystallisatlon froa ethanol gate aonodlnlc crystals vlth a density of 1.967* This compound sintered slightly et 10Q*C., and darfcensd at 150 *C., hut did not w it below 3?0°c. The product vas soluble In voter, pyridine, acetone, and methyl and ethyl alcohol, and was Insoluble In bentene, hexane, chlorofona, dloxane, ethyl ether, and ethyl acetate. The ab- sorptlon spectra In vater and in alcohol are given in figure 6.

The structure proposed for this compound is:

Calculated Mol. v t., 566 C, 25M f H, 2.1# u, b2.0# Found Mol. V t., 287 C, 2 5 .5 # B, 1 .7 # U, * 2 .3 # The presence of one molecule of vater vas shovn by titration vith Karl Fischer reagent.

The infrared absorption spectrum le presented In figure 7. Bands of pyridine mleslng in the complex were: 2.76, *.36, 5*02, 5.20, 5 .*6, 5*93 (C*H stretching), 6.33# 6.21, 8.72, 9 .# , and 10.09 (c«C ring stretchinghi. The 3*90, 6.02 (C«0 stretching), and 8.66jx salicylic acid bands'also vanished. Ho bands observed In uranyl nitrate hexa­ hydrate were absent. Rev bands appeared at 11.25, 11.5, 13.20, 13*53, and 14.9 ju *wm*n

100

£ /Assn / !• 0.otp / 50 f / Wofr / 0.00/ ¥

300 350 400 460 500 680 600 650 700 *•#**

Figure 6 8pectr* of BalicyUtoaquopyridioouranyl titrate in Water and Bthanol

T /I lir'*YV K 7*V k Jn f ¥ \

Figure 1 - 4 Infrared absorption Spectra of [HgO'CjH^W^C^OjJJ (60^)

This ccapouad foxmed a stable, acidic solution in eater, and was stable toward at Idly —smiirsl solutions, but was decomposed by strong acids and bases. Its eater solution was a conductor of electricity, and tests for n itra te loa were obtained. So change was noted upon the addition of a solution of 8-hydroxyquinollna. These factors led to the aselgnaent of a Werner-type structure for the compound, its use as a sensitive te s t fo r urscyl ccapounds in organic solutions has been reported (5).

Analogous compounds were prepared using as the external base asaonia (bright orange powder, non-fusible), n-butyl n in e (red-orange powder, no aelting point), and quinoline (brick-red crystals, Belting point above 300*CY}— — __ 19 In aqueous solution, aniline and pyridine have been reported to \for» anillno- and pyridinouranyl salicylate vith uranyl nitrate and salicylic acid (58). These are inner complexes which hare an extra coordinating group attaehedf and not Werner-type complexes.

Vo difference could be detected In the reactions of uranyl nitrate hexahydrate and uranyl nitrate dihydrate vith salicylic acid and the above amines.

Uranyl chloride Trlhardrate, Atteapts to prepare the corresponding complex of uranyl chloride were less successful. The addition of 3 ml. of pyridine to a hot solution of k g. of uranyl chloride trihydrate and 5 g. of eallcyllc acid in 50 ml. of butyl acetate precipitated a red tar which solidified upon rubbing under ether* Fractional crystallisation of this product In nethanol yielded only 0*5 g* of a relatively inpure reddleh-brovn salicyUtoaquopyrldlnouranyl chloride, n.p. above >0O*C., and 2.0 g. of aguopyridinouranyl chloride! a yellow powder with no Belting point below 360*C.

Analyses: (^0^ 585* (Cl)" Calculated C, 26.70* H, 2*22* U, 4t.ll* Found C, 26.21* H, 2.73* U, 45.13* TOgClg-VgO-C^

Calculated• C, 15*62* H, 1.59* U> 55*98* Found C, It .20* H, 1.76* U, 5t.t7* Uranium(IV) Chloride. The analogous uranium(IV) compound could not be isolated and purified because of suceptlblllty to hydrolysis and oxidation. When 5 ml. of pyridine was added to a solution of 10.0 g. of uranium(IV) chloride and 20 g. of salicylic acid in 125 ml. of butyl acetate! a greenish-white precipitate weighing 22.0 g. was obtained. This weight corresponded to a composition of approximately two moles each of pyridine and salicylic acid for each uranium atom. Analysis of the crude material confirmed this ratio.

Analysis: UClg(0^0^*20^11

Calculated C, 32.U0* H, 2.75* U, 32*U* Found, crude material C, 33*41* H, 5*02* U, 34*56

The nost probable structure for this compound Is: » i# W l

80

This infrared absorption spectral of this compound is given in figure 8. The following ligand bands sere lacking in the spectrum t 2.76, 5.02, 5.20, 5.46, 5*95 (C*i stretching), 6.33, 6*9** and 9*38 it for pyridine; 5.90 and 5.02 m for salicylic acid. Also the 10.5-10.7 w band of uranium(rv) ctflorlde vanished* Rev bands were observed at 2.95, b .10, b.85, 5.60, 7.10, 9.55> 10.85, IO.96, 11.26, 11.53* 11.72, U .96, 12.80, 13.05, 15*20, and lb .85 j i.

figure 8 Infrared Bpectraa of

When free of solvent th is product was stable, but It oxidised rapidly in acetone or alcohol solution, and hydrolysed rapidly in aqueous solution, with the formation of a precipitate of uranltm(Xf) oxide. Conditions for re crystallisation could not be found which did not yield Mixtures of uranyl products, for example, recrystallisation fro® pyridine gave predominately dipyrldlnouranyl chloride, a yellow, in­ soluble powder. Belting point 115*C«, with decomposition. Analysis: UOgClg*20^11 t Calculated C, 24.05* H, 2.02)1 found C, 83.85* H, 2.185 The molecular weight in boiling alcohol corresponded to one-third the formula weight. lactic Acid Lactic acid, acetylacetone, and other acidic completing agents reacted aa did salicylic acid with solutions of uranium compounds, with the exception that the products were aixtures of inner complexes and Werner- type complexes, fractional crystallisation from ethanol served to separate the more soluble inner ccmplexes from the less-soluble ionic complexes.

4 1*—«►*>•-

21 Prenyl KltrateHoxahydrete, The Infrared absorption spectrum of the bright yeliov iactat^iiopyridlnouranyl nitrate la given In figure 9*

Infrared Absorption Spec^uTS ^HgO • (C^H •POgtCHjCBOHCOO)]4! ^

Bands lacking in the complex were: 2.05 and 3*00j x of uranyl nitrate bexahydrate; 2.76, 4.36, 5.02, 5*20, 5*46, 6.94, 8.21, and 10.09 (C=*C ring stretching) /i of pyridine; and 5*02, 7*30, and 9*13 M of lactic acid. Rev bands rare observed at 6.90, 8.07, 11.13, 12.25, and 15.15 fk. Pranium(IY) Chloride. The infrared absorption spectrum of the light* green complex, dlchlorodilactatodipyridinouranium(lV) is given in figure 10. The probable structure; off this compound 1st

It vas prepared by a procedure similar to that used for the reaction of salicylic acid with uranium(IY) chloride. Figure 10 infrared Absorption Bpectrtm of (C^H^IJgOClgfCHjCHOHCOO)^

The following bands wars deficient in the complex's spectrumj 2.76, 3.28 (C-H stretching), 3*30, 3-32, 5*02, 5*20, 5*1*6, 5*93 (C-R stretching), 6.9t, 7*29, 9*36, 9*70 (C-C-C bend perpendicular to ring), and 10.09 (OC ring stretching) u for pyridine; 5.02, 5*85 (C«0 stretching), and 8.90/x for lactic acid; and i.On for uranlum(IV) chloride. Rev bands were observed in the complex at 7*3, 7*70, 11.90, and 13.08^1.

RSACTBSB WHO AMBK) ACIDS In most cases, amino acids did not require the presence of an external base for reaction with uranium compounds, and the products, especially with uranyl compounds, were much more similar to true chelates than vere those obtained from acidic complexers.

Qulnaldinic Acid > ttrany]r l titrate Hcxahydrate. The addition of 3 g* of qulnaldinic acid in 50 ml. of hot butyl acetate to m solution of 2 g. of uranyl nitrate bexahydrate in 50 ml* of butyl acetate gave an immediate precipitate. Digestion of the solid vith ether, followed by recrystallixation from ethanol, gave about 2 g. of light yellow, alcrocrystalline diquanaldlnato- dlaquouranyl, which melted at 243-2^5*0., vith decomposition. The pro* duct was slightly soluble in water, but was decomposed by strong acids • or bases. It was soluble in methyl and ethyl alcohol, acetone, and methyl ethyl ketone, and insoluble in ethyl ether, diaxans, ethyl acetate and bentene.

Analysis: uo^c^g^-ay) c Calculated C, 36.80)1 a, 2.78)1 0, 36.48)1 Pound C, 36.00)1 H, 2.43* U, 35.81)1; 23

The proposed structure it:

Figure XI Infrared Absorption Spectrua of BOgCC^HglfOg) *2HgO

Figure 11 is the Infrared absorption spectrua of this compound. The 6.35 (tvitter ion), 7.63 (tvitter ion), 8.73, 10.13 and 11.3/1 vibrations of quinaldinic acid and the 3*0, 6.65, and 12.0 jx bands of uranyl nitrate h/exahydrate disappeared, lev bands were noted at 6.00, 6.80, 6.83, 8.60, and 8.82/1.

Uraniua(IY) Chloride. A solution of 5 g. quinaldinic add in 75 nl. of warm butyl acetate vas added slowly to a solution of 2 g. uranium(IV) chloride In 30 ml. of butyl acetate. A curdy, cream-colored precipitate formed, leaving a slightly greenish supernate. This precipitate was digested twice with ether and recrystainted with difficulty from ethanol to obtain 2.3 g. of dichlorodlquipaldinatoqulnaldinouraniua(lV), a light yellow-green amorphous solid which sintered at l60*C., and melted with decomposition at 223*C. This compound vas slightly soluble in methanol, ethanol, and acetone, and vas Insoluble in other organic solvents. It dissolved In water with tbs formation of a brown pre­ cipitate and an acidic solution. Addition of sulfuric acid to this solution and distillation into silver nitrate solution gave a precipiate of silver chloride. mmm

' :

k u l j B U : CgYg-treyCjKgOglg

Calculated C, 43.75} H, 1.9# «, 28.91} Found C, 43.08jf H, 2.12} 0, 29.86}

This compound has the probable structure:

The infrared absorption spectra is given In figure 12. The following bands disappeared In this couples: 2,7$, 3.20 (C-H stretching), 3»30, 3.32, 5.02, 5.20, $M, 5.93, 6.9*. 7.29, 9-36, 9-70 (C-C-C bend perpendicular to ring), and 10.09 (C«C ring stretching) n for pyrl&lue; 3.90 w for salicylic acid; and 10.5-10.7, 13.37, and lb'0O 11 for uranium (IV) 'chloride. lev bands absorbed at: 3.25, 3*70, 5.10, 6.05, 6.15, 6.80, 6.90, 7.^2, 8.20, 0.60, 8.80, 8.90, 11.65, and 11.95/*•

Figure 12 Infrared Absorption Spectra of UClg ( M

■-Phenylglyclne

Pranyl Sltrste Basahydrate* Uranyl nitrate solutions developed a deep red color with B-phenyiglycine, but showed little tendency toward precipitation. Prolonged beating or the addition of non-polar solvent a did not Induce precipitation. This failure to fora an insoluble product nay Lave been due to the lover molecular weight of the complexing agent. The addition of a large excess of pyridine did.cause the formation of a small amount of tarry precipitate, but this could not be purified by crystalllratIon. Evaporation of the solvent also resulted In the formation of an intractable ^ t .

Uranlum(IV) 'Chloride. To a solution of 10.0 g. of uraniun chloride 150 rnJ» of butyl acetate ras added slowly a aclution of 15 g. of H- phenylglyclne in 75 ml* of butyl acetate. The resulting oil was stirred under ether until it solidified, and the solid vas digested several times.with ether. After drying, there remained 23*0 g. of a light green, amorphous solid. Because of its poor solubility characteristics and acceptability toward oxidation, this crude dlchlcrobls(V-phenylglyclnato) H-phenylglycino complex could not be purified further. It was very slightly soluble in alcohols, and oxidized upon standing, insoluble in ketones, ethers , esters, and chloroform, soluble In dimethyl*ormamide, and soluble in water with hydrolysis.

To prepare the bie(H-phenyig\ycinato)t*traaquouraniua(IV) chloride, 5 g. of the solid obtained suove was dissolve by heating in 25 ml. of water containing 1 ml. of 6 {[ , and tlmn 10 ml. of a saturated sodium chloride solution was added. After thorough chilling, the solid was removed by filtration and washed with ethanol and ether, and re- crystallited from a water-dioxane mixture. The light green solid weighed 2.0 g. after drying. This compound lost rater at 115*0., and darkened and decomposed slowly above 220*C. It was soluble in water, slightly soluble in methanol, and Insoluble in other solvents. All of the chloride in the molecule was precipitated by silver nitrate solution.

taalyils: jjn^O -U( CgHgHOg) 2j ^ (Cl'i^

Calculated C, 28.20)1 H, 2.29* Cl, 10,41* ,0, 10.58* Found C, 27.72* H, 3-12* Cl, 10.49* 0, 9.6*

The probable structure of this complex is:

H,0 0« C —— Ov 4 •c«o \ ’ /'

o 2h H / . V H h 2 c — N' f ^N— '2 I h20 I / V s csHs

The absorption in the visible region of an 0,002 M solution is given in figure 13. Figure Ik presents the infrared absorption spectna. fa* ****** The following bands of the components were absent in the complex: 4.26, 4.62, 4.99, 5.79 (acid C»0 group), possibly of phenylglycine, and the 3.00>4.0, and chloride. New bands were observed 6.25,at: 9-75, 10.0, 11.97# and and Randall (25). Itis Important to note that these assignments and Infrared Spectra correlations given by Barnes (4), Colthup (7), Kline and Turkevich ( Interpretations were made with the help of thespectra-structure %.T Spectrum of Bie(H-Phenylglyclnato)tetraaquouranium(lF) Infrared Absorption Bpectrum of rutCgHcHHCHgCOOjg^HgOl (Cl )g 13.35 u. Chloride la 0.002 M Aqueous Solution DISCUS8I0H Figure13 iue 1 n + Figure 14 OF RESULTS 6 . 69 ,7.68, 10 . 5 - 10.7 7 . 78 u bands of uraniua(lV) ,7.95, and 8.18 u 4 . . 26 19 ), 27 interpretations are based on more or less empirical relationships, and thus the conclusions presented are merely the speculations of the authors, nevertheless, it is interesting to observe some of the correlations and trends existing between complexes containing similar groups.

It will be noted that In each of the complexes in which pyridine is a component that sene of the bands normally present in pyridine are no longer present in the complex. This is also true for the other co­ ordinating groups studied, although the effect is less pronounced.

In the six complexes which contained pyridine the following pyridine hands vanished in all cases: 2*76, 5*02, 5.20, 5.46, 5.93 (OX stretching), and 10.09 (C-C ring stretching) u. The 6.33 and 9*3o M bands disappeared in all save l&ctatoaquopyridinouranyl nitrate; the 6.$4 n band in all save dipyrldinouranyl nitrate. This helps to hear out the theory that the pyridine molecules are Joined to tbs uranium atom by the two unshared electrons present on the nitrogen atom, since only the attach­ ment of a heavy atom would account for such a profound change in the spectrum of pyridine.

A new band at 11.2 u appeared in the sallcylato and lactato uranyl and the sallcylato uraniua(IY) complexes. Belated bands were noted at 11.63 M for the quln&ldlnlc acid derivative, and at 11*97/i for the pbenylglyclne derivative of uraniua(IV) chloride. A shift of this sort is expected with the attachment of heavier groups on the postulated U-0-C(0)-R linkages involved in these complexes. In addition, a band at 13*0-13*2 u appeared in all pyridine complexes except the dipyrldinouranyl nitrate and trlpyrldlnodluranl\a(X?) chloride. This night again be due to a U-O-C(O)-R linkage, which is absent in the latter two compounds.

In both the uraniun(IV) and uranyl sallcylatopyridino complexes the unexplained salicylic acid band at 3*99 /» disappeared. In the uranyl derivative the 6.02 jx carbonyl group stretching vibration and the 8.6611 band also vanished.. In both comparable lactic acid derivatives • the £.02 and 7.30 fx bands were absent. The missing 5M band is probably a carboxylic acid type carbonyl vibration. In both sallcylatopyridino complexes new bands were observed at 11.2, 11.5, and 13*2 n; in the lacktatopyrldino complexes at 13.1 jx* In both the uranlua(XV) and uranyl derivatives of qulnal&lnlc acid the Uguid bands normally at 5*23# 6.65# 7.63, 10.15# end U.5 M vanished In/ ; the respective complexes. In both complexes new bands appeared at about 6.0, 6.0, 6*9, 8.6, and 8.6 ja.

In dlmethylethanolaminouranyl nitrate many bands of dimethylethanolaalne attributed to the tertiary nitrogen disappeared. One band, which appeared In this complex at 7*08 jx, la probably due to quateralxatloa of the nitrogen. '

The Hechanltm of Complex formation

A typical method of preparation of inner complexes consists of the cautious addition of base to a solution of the ccmplexlng agents and a uranium salt in water, thus obtaining a precipitate of the coaplea (38). When, *n this work, the same method was extended to non-aqueous systems, with the substitution of an amine for the base, and n-butyl acetate or a similar oxygenated solvent for the water, a product waa obtained which was ionic In character. These compounds were more soluble in water than in organic solvent a, many did not malt below 560*C., and a portion of the original Inorganic anion was retained, held by ionic bonds to a complex cation. Also, the amine was found combined in this cation• Such properties are much more suggestive of Werner-type complexes than of true chelate compounds.

One possible explanation for the formation of such widely different products In different solvents may be found in a comparison of the degree of ionisation of uranium salts In aqueous and non-aqueous solvents. In aqueous solutions uranyl and uraniua(IV) compounds exist primarily as the completely ionised anion and a highly hydrated cation, while in relatively non-polar solvents such as n-butyl acetate the salts undoubtedly exist as molecular species, more or less eolvated, and at best only poorly ionised. In aqueous solution tbs cation would be susceptible to successive stepwise attack by complexing agents ic form a totally coordinated complex. On the other hand, in non-ionising solvents the separation of anion and cation proceeds only after the attack by an external base upon the molecule. The preferred attachment of this additional basic group weakens the attraction of the uranium atom for the electrons of the partly-covalent, partly-ionic shared pair of the inorganic anion. In the presence of an acidic function which could form a more covalent bond with the uranium than could the orginal anion (e.g., a carboxylic acid), stepwise attack on the cation could then proceed. The process would, however, be interrupted at the point where the total number of coordinating groups surrounding the uranium atom becomes equal to the maximal coordination somber of the uranium atom. At this point the bonds between the uranium atom and any remaining inorganic anions would become strongly ionic in nature, and the resulting Werner-type complex, being Insoluble In the organic solvent, would be removed from the sphere of reaction. This point apparently la reached after the substitution of one bldeutate organic acid type complexar in uranyl salts, and after the substitution of the second molecule of the camplexer In uranlum(lV) salts.

Some confirmation of the part which the base plays in this reaction may be found in tbs fact that complexing agents which contained only acidic functional groups, such as salicylic acid, had little effect on uranium compounds In organic solution and required the addition of an external base for reaction. On the other hand, a complexing agent such as quinaldinlc acid, which Itself contained a basic group, reacted readily with organic solutions of uranium salts to give chelate compounds.

This relationship may be seen In figures 15-17* These figures contain the spectra of solutions of uranium compounds to which various complexing agents were added, lone of the carboxylic acids added caused any significant change in tbs spectra of the uraniua compounds. Acetylacetone produced a decided bathochroalc shift in the spectra. Several of the hydroxyanthraquinones and flavones, which have a structure similar to that of enollzed acetylacetone, have shown similar results (5)* Home of these compounds gave precipitates with uranium compounds, however. 100

SPECTROPHOTOMETRIC EVIDENCE FOR THE REACTION OF COMPLEXING AGENTS WITH URANIUM (33T) CHLORIDE IN BUTYL ACETATE

m u n i is M W VT

i

SPECTROPHOTOMETRIC EVIDENCE FOR THE REACTION OF COMPLEX1NG AGENTS WITH URANYL NITRATE HEXAHYDRATE IN BUTYL ACETATE

FI0UKK IS

1 %.T Lactic acid added

SPECTROPHOTOMETRIC EVIDENCE FOR THE REACTION OF COMPLEXING AGENTS WITH URANYL CHLORIDE IN BUTYL ACETATE

FIGURE IT 52 The extensive alteration of the epectra produced by 5,7-dichloro-8- hydroxyquinoline is typical of the results obtained with amino adds, and is Indicative of actual chemical combination of the complexing agent vith the uranium compound.

The lack of reaction of hydrated uranyl compounds vaa not due to the protective action of the vater of hydration, as is demonstrated by a comparison of figures l£ and 17, which show th a t anhydrous uranyl chloride underwent no more reaction than did uranyl n itrate hexahydrate.

Structure Considerations

In both Inner complex compounds and Werner-type complexes, the uraalum(XV) ion has a coordination number of eight. The uranyl Ion has a coordination number of four, thus giving the uraniua(WI) Ion a coordination number of s ix . The difference in the covalent rad ii of uranium(IV) and uranyl ions would not be great enough to explain this difference, and the uranyl ion has been shown to have a coordination number of high as eight In basic aqueous solutions (15)« Instead, this apparently low coordination number of the uranyl ion may well be the result of the sharing by each oxide ion of two pairs of electrons vith the uranium ion, thus in effect satis­ fying two positions in the coordination sphere for each oxygen atom. A uranyl ion so constituted would contain the stable outer structure of an eight-electron shell, with a rrobable 7*f configuration. This la in agreement vith the concept of a linear uranyl ion (15)* The great stability of the uranyl Ion, and the profound tendency of uraniua(WI) compounds to acquire the required number of oxygen atoms to form tb s uranyl ion (3,5)# lends additional support to this Idea.

The changes observed in the Infrared spectra of the uranium complexes as compared to the spectra of their components help to confirm the hypothesis that the organic ligands are bound to the uranium(XV) or uranyl ion rather than being present as purely ionic components of the complexes. Pyridine appears to be bound by linkage between the tertiary nitrogen atom and the uranium. The amino acids are attached through both the carboxylic group and the amino nitrogen atom.

a The susceptibility toward hydrolysis and oxidation of the dichloro- uranium(IV) complexes Indicates that the chlorine atoms were s t i l l quite ionic in character, and thus were easily displaced from the « actual coordination sphere of the uranium atom by the action of strongly coordinating groups such as water and the oxide ion*

8UMKAKX

A study has been made of the action of organic bases, acidic complexing agents, and amino acids upon organic solutions of uranyl nitrate hen- hydrate and uraalm(IW) chloride.

•She type of product which was obtained from the reaction of complexing agents with uranium s a lts in organic solutions was a function of both the oxidation state of the uranium and of the type of complexing agent. Acidic complexing agents did not form isolable products with e ith er

v. . • .'.‘/ ' . i - 7 l ‘ .•’.'rrtriii r T . r * - r ^ f a i J "< itofsMwwst' iw»f»> 33 uranyl or uranium (IV) salts, although spectrophotcoetric evidence Indicated that enollzed 1,3-diketonee complexed to some extent. The addition of an external base to solutions of a uranium compound and coorplexlng agent precipitated Werner-type complexes.

Cduplexing agents which themselves contained an amino group reacted readily ulth uranyl sa lts to form inner complexes, and with uranium(IV) sa lts to form Werner-type complexes.

I t was proposed that the mechanism of the reaction depended upon pre­ liminary attack upon the uranium atom by a basic group, followed by step­ wise substitution of the inorganic anion by ccoplexing groups until the marl mm coordination number of the uranium atom was attained.

Interpretation of the Infrared spectra of 10 of the complexes and aamlnes confirmed the formation of a strong uraniua-to-nitrogen co­ ordinate-covalent bond, and shoved th a t, in the complexes, the coordinating groups were bound directly to the central uranium atom. Preliminary assignments have been made for absorption bands due to the RC00-U linkage.

ACKHOWI&XSEMEirr

The authors wish to acknowledge the assistance of Miss Francis Ball In the determination of the analyses, and of Miss Patricia Volta in the measurement of the Infrared absorption spectra.

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43. Woodward, R. W., Leh, D. A., and Tilson, Pi V., Organic Cocplexlng §!& Agents for Uranium, Carbide and Carbon Chemicals Corporation, T-12 t e ? * i? • ^ Plant, September 21, 1948, (t-106). 111 m If 1; ■ ■ 11 ; NOTEBOOK REFERENCES — m ’i - -t J . T . B arr, K-25 Hotebook Huabers 1498 and 1597* - ; I > C. A. Horton, K-25 Hotebook Jhabers 1455# p . 87-8, and 1620, p. 26-7# 89-73* U4'• ■

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