17!)0
WARRENL. REYNOLDS AND ROGERH. WEISS
tion of tlic sample gas with something in the sample iiilct system, ion source or analyzer system of the spcvtronieter. If adsorption of some species is the priiiiary process on which these ions depend and the adsorbing surface is only partially covered a t low imssurcs, a first-order pressure dependence is expected. At higher sample pressures the surface should be coiiipletely covered, with the resulting effect being the apparent zero-order dependence observed. Support of this thrsis was given by an expcriiiient in which the sample gas pressure in the source was greatly reduced. z\t pressures froin of the normal operating pressure, the ion to currents a t lower electron accelerating voltages varied directly as the pressure. I t is interesting to note that the initial break of the ionization efficiency curves for all the ion fragiiirnts exhibiting an apparent double break was 12.5 & 0.5 volt for fluoride fraginents and 11.5 f 0.5 \. olt for the chloride fragmclits. This sinal1 variation suggests that these fragment ions riiay be of t 1icrni:il origin and not priinarily due to electron I)oiiibardincnt of gascous niolecules. Fibwre 3 shows the Crf and CrOF+ curves as tlicy appeared :it various pressures. The upper nuiiibcr of each C r + curve is the relative ion current fur 45-volt electrons. The lower number is the value a t 25 electron volts. The number on the CrOF+ curve is the relative ion current for 45volt electrons. I t would be teiiiptiiig td interpret these results
Vol. 81
in terms of the bond strengths in molecular or ionic chromyl halides. However, the paucity of ionization potential data for CrO+ or CrX+, coupled with the uncertainties involved in interpreting such data as presented in Figs. 2 and 3, make any such attempts tenuous. It is apparent that no simple assumptions can be made which represent the processes occurring in the mass spectrometer ion source. While the results listed in Table I appear to have some consistency, their meaning in terms of thermodynamic properties in these molecules is clouded by the several likely processes which may be occurring. Attempts to measure ionization efficiencies of the chlorofluoride fragment ions were unsuccessful, except for the parent ion. The value of 14.0 & 0.2 volt was obtained for the appearance potential from a measurement on a mixture containing very little chloride. The aniount of chlorofluoride present was large enough to give reliable results, and the very small amount of CrOC12+ from the chloride did not interfere. I t was thought that the appearance potential curves for some of the ions of a mixture of all three compounds might exhibit two or more “breaks” because of the probable difference in appearance potential of the ions which are common fragments of two or all of the compounds. In the course of work reported here, no such breaks were observed on any mixture. AMES,IOWA
[CON.I.KII~III [ o s VKOM n i u SCIIOOL OP CIIEMISTKY, UNIVERSITYOI? MINNESOTA]
Iron(II1) Complexes in Non-aqueous Solvents. I. The Solvolysis and Chloride Complex Constants in N-Methylacetamide BY
\\‘ARREN
L. REYNOLDS A N D ROGERH. Wmss RECEIVED OCTOBER 31, 1958
~~~cctr~,~ili~~toiiictric I I I I ~ : I S U ~ C I I I ~ I>~ivl~lctl I~S v;ilucs of 1.4 x 10-3 mole IiLcr-1 for tlie first sulvolysis constant of Fc+Ja i d 7.5 x 1 1 t z ;tilt1 I:{ litcr IIII)IL~--’ for tlir stcl)n.isc f~~riiiatiori constants of 1;eCl + * atid PeCla+, rcspectivcly, iii S-iiictli)lucetariiide. ‘flie iiio1:ir e~tiiictione ~ r l l i c i c i ;it ~ t ;Hi:; ~ iiip o f I‘eS+’ (wliere S - is the solvent anion), PcCl +* and FcCl?+ were found to bc over tlic l . x x 1 1 1 3 , 1 .$I x 1 0 3 : i r i d $1 3 x 1 0 3 . rvspectivcly. T h c average latcnt heat of vaporization of S-iiictl~~lacelaiiiiiic t eii~p~r:itiirc interval f r ~ m i115 to 26’ wits 14.9 kc:il. mole-1.
Experiiiierits, preliiiiinary to an investigation of
+
the kirictics of the Fe(1I) Fe(II1) isotope exchange rcnction i n the solvent N-methyl‘acetamide ( X l l L l ) ,showecl that Fe+$formed complexes with the :inions o f various strong miueral acids in this solvelit. In 1;ig. 1 are showii spectra of anhydrous FeCl:, tlissolved iri NLI.1 in the presence of 0.1 ill
The solvent NMA was chosen because it could bc obtained in relatively anhydrous condition (water concentration less than 0.002 Jl) and because of its high dielectric constant.
Experimental Reagents.-Iron(II1)
chlt~ ide was prel)nred by p;issiiig
liytlr~cliloric.~)cn~Iiloric anti sulfuric acids and in dry Clr gas a t rooin teiril)er;iture through iroii powdcr in tlie tlic. absciicc (Jf aiiy adtlcti acid. Presence of watcr presence of a trace of moisture. The product w a s :iii:ilyzed iodonictrically and foiiiid t u coiitain 96. I I I ) t o a t ivast one iiiolar had no :tppreciahle effect on main iinpurit:~probably was watcr; tlic :ilnotlllt o f water t I i i w spectra. ilp1xirently coinplexes between iritrduccd with the 1’cCl3 in ~ i ~ ; i k i nag0.1 ruillirnolar soluw:itcr I;c +:$ aiid tile various ;inions arc formed which have t i o ~of ~l’c(II1) was tlius 1iiuc11less than the ~niriiriiuii~ which could be detected by a Karl Fisclicr titration. tliiTcrcii t nbsorl)tion spectra. ’Therefore, before content Hydrochloric acid solutions were prelxircd by Imssing d r y t lie rate coilstants for the isotoiiic cwhingc of iron I-ICI gas’ into NRIA. ‘l‘lic solutions were ~tu11t1:rrdizedl ) v I)c,twcc.ii l;c(l I ) : i i i t l tlic varioiis I;e(III) c o i i ~ p l c ~ sdilutiiig ail aliqiiot a t Ic.i.;t twifi,ld with witer mltl titr;itiiil: ;is iiidic:ttor. wit11 standxrd K:iOlI using ~~licnol~~lith;ilciii c.o111(1 I x tlc.tc.riiiiiivt1, it was iicsccss:try to cli.tcriiiiiw NaCl solutiuiis were prepired by dissolviug allthe associatioil constaiits oi soiiic’ o f lhcsc coiiiplcxcs _Standard _ ~ a i i d to clc.tcriiiiiie the solvolysis coiistmt of I’e + 3 ( I ) “lni,rganic Synthmes,” Vol. I , XIcCraw-Irxll I3ook C o . , I n c , i i i NA1.l. N e w York. N. Y . , p. 117.
IRON(III) COMPLEXES IN NON-AQUEOUS SOLVENTS
April 20, 1959
value of 205" for the normal boiling point3 the average heat of vaporization calculated from the integrated ClausiusClapeyron equation was 14.2 kcal. mole-'. The value calculated from the data in Heilbron3 for the temperature interval 140.5 to 205" was 13.1 kcal. mole-'. This value is significantly greater than the value for water and probably results from the larger dipole moment of NlMA as compared to water. Fe(III) Complexes Present.-Iron(II1) spectra a t three different HC1 concentrations and an Fe(II1) spectrum in the absence of added aFid are shown in Fig. 2. The presence of an isosbestic point a t
02'
260
1791
280
300
320
340
S O
380
400
WAVE LENGTH IN Y L L I U Q I W .
Fig. l.--[FelII] = 8 X 10-6 111: (a) 0.1 61 HtSO,; (b) 0.1 A I HCI; (c) 0.1 Af HCI04; ( d ) no acid added; cell length 1.00 cm. hydrous Mallinckrodt analytical reagent in a known volume of solution. The anhydrous solvent was prepared by a slight modification of a method recommended by Mr. Laurance Knecht' of this department. Gaseous methylamine was distilled into excess glacial acetic acid. T h e resulting solution of metliylammonium acetate in acetic acid was r d u x e d overnight, followed by elimination of the major portion of the water produced in the amide formation by distillation a t atmospheric pressure until the temperature rose to approximately 120'. Concentrated sulfuric acid then was added to conibine with any unreacted methylamine and distillation was continued under reduced pressure. The first portions of distillate containing large amounts of acetic acid were discarded. The acetic acid content in the remaining distillate was reduced to less than polarographically detectable amounts by shaking with calcium oxide or with potassium carbonate and then distillation a t reduced pressure. The resulting solveiit containing some water was polarographically and spectrophotometrically pure. Distillation from phosphorus pentoside under reduced pressure introduced a small amount of an impurity reducible a t a dropping mercury electrode a t - 2 . 2 v. relative to a saturated calomel electrode. This material apparently did not affect the slvxtrum of the pure solvent or the spectra of the Fe(II1)CI- complexes. The solvent was analyzed for water by a Karl Fisclier titration and found to contain less than 0.002 molar water. Apparatus.-A Cary model 11 recording spectrophotometer was employed for scanning of slxctra. A Beckman model DU spectroi)liotometer was employed for absorbance measurements a t constant w:ive Iciigth. Since neither instrument was tlierniostated a t the desired temperature and since the solveiit froze at approximately 29'. it was necessarv to warm a sample somcwhat before placing it in the cell compartment. Thus the temperature of a sample was never accurately known. This affected measurements in two ways: namely, through the effect of temperature on the equilibrium constants involved and through the effect of an appreciable volume change on concentrations and resulted in an error of about 6Vc in the proportionality constant between absorbance and Fe( 111) concentration. Procedure.---Stock FeCI3, SaC1 and €IC1 solutions were atldrd to solvtwt in such rtniuunts an to give the desired concentrations of F e ( I I I ) , CI- arid H + . The absorbarices of s~,lutioiisin 1-crn. glass-stoppered silica cells were measured at 3 5 1 ..5 and :UY3 nip. Absorbance measurements were made a t X i 1 5 n i p because tliis wave lerigth wis apparently an isosbestic point for the two species FeCI + 2 and FeCL+, and a cliaiige of absorlxince a t this poitit indicated the appearance of atlclitioiial forms o f I'e(II1) i t i solution. The absorbance iiie;isiirements a t 3lZ nip w c w iisvrl to calculate the molar cstiiictioii ciielticierlts a r i t l equilihriutli coilstalits desired.
Results and Discussion Latent Heat of Vaporization.-'l'lic liuiling poirit o f S1I.I a t 2-4 i l i ~ i i .pressure was l 1 5 O . Using the (2) Unpuhlislled thesis research.
280
300
320
340
360
380
4
0
WAVE LENGTH IN MILLIMICRONS.
Fig. 2.-(a) 8 X low6ill FeCI3, no acid addcd; (b) 7.4 X Jf FeCI3,0.025 10-8 111FeCI3. 0.015 U IICI; (c) 7.4 X Af IICI; ( d ) 7.1 X 10-5 111FeClr, 0.10 M HCI; cell length 1.00 CIII.
approximately 351 nip suggests that only two Fe(III)-Cl- species are predominaiit in the concentration range 0.015 to 0.10 d l HCl. These wcre assumed to be FeC1+2 and FeCI:!+. The Fe(II1) spectrum in the absence of added acid was the same whether the dissolved Fe(II1) salt was the chloride, perchlorate or sulfate. Hence this spectrum was attributed to either the solvated species Fe(HS),+3 of the solvent HS or to the solvolyzccl species FeSf2 or to the aquo-ferric ioii or to the FeOHf2 ion in N I I A . This spectruiii was unchanged by the addition of water to the anhydrous solvent ([Hz0]
