Inorganic Infrared Spectroscopy - ACS Publications

---

John R. Fermro

Argonne National Laboratory

Inorganic Infrared Spectroscopy

Argonne, Illinois

The infrared spectroscopic study of organic compounds has been known for a number of years. H~T-ever,the application to inorganic compounds had a rather late start. Most of the early work on the spectra of inorganic compounds concerned itself with the Rnman method. Some reflection infrared met,hods (1) \\.ere used as early as 1930. The lack of high resolution spectrophotometers, plus the fact that adequate specimen preparation was difficult, 'delayed progress in this field. Since large particle size gave Based on work performed under the auspices of the U. S. Atomic Energy Commission. Paper presented at the Infrared Spectroscopy Institute, Canisius College, Buffalo, N. Y., August, 1960.

considerahle scattering losses most early infrared spectra of inorganic compounds gave broad, indeterminate absorptions, as compared t,o the sharp spectra obtained with organic compounds. With the introduct,ion of commercial high resolution spectrophotometers and new methods of specimen preparation, such as Nujol mulls, pressed discs (2-4, and recently t,he rare gas matrix isolation method ( 5 ) , the spect,ra obtained were found to be excellent. Revival of interest in inorganic compounds took place. Lecompte must, be given credit for a considerable amount of this revival. He and his co-workers examined numerous inorganic compoynds; the studies appeared

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201

in a series of papers in 1951 (6, 7). A very important paper appeared in 1952 by Miller and Wilkins (S), which added further impetus to this revival. They examined the spectra of over 150 salts in the sodium chloride region, and their correlation for analytical and identification work proved excellent. The Isst eight years have seen a tremendous development in the progress in inorganic infrared spectroscopy. I t is the purpose of this paper to point out some of this development with a particular emphasis on the application of infrared spectroscopy in solving various problems in inorganic chemistry. The paper is in no way intended to be a complete bibliographical report on this subject; it attempts to serve only as an introduction. Qualitative and Quantitative Analysis

I t is beyond the scope of this paper to present elaborate details concerning the qualitative and quantitative analysis of inorganic compounds utilizing infrared methods. When used with other physical measurements, the infrared method aids inorganic analysis. The numerous reviews (9, 10) in the literature of quantitative organic analysis cover principles also appropriate for analysis of inorganic compounds. Miller and Wilkins (8) reported the spectra of numerous inorganic compounds. I t is apparent from their work that a number of discrete regions of absorptions, characteristic of the type of anion present, existed. From their correlations, qualitative identification of various inorganic anions can be made. Similar correlations have since been made for various other inorganic compounds, and these can aid in making qualitative identifications. Several examples of quantitative infrared analyses have been reported. Hunt, et al. (11, 12) analyzed a number of rocks and minerals. Corbridge (13) has done some quantit,ativework on condensed phosphates.

ailother absorption was found in the 1800 em.-' region, and this was attributed to the "hridwd" carbonvl group. The assignment of the hand in the region of 1800 cm-' as a bridged carbonyl group finds confirmation from chemical means (20). Two bridging groups can be replaced by one molecule of acetylene, and the 1800 cm-' band disappears. Iron enneacarhonyl, Fe~(C0)swas found to have C-0 stretching frequencies a t 2080 and 2034 cm-1 due to the terminal carhonyl groups, and an absorption a t 1828 ern-' due to the bridged group (16). The structure proposed for this compound was that shown in Figure 1. Diffraction data for iron tetracarhonyl Fea(CO)lzare not available. From infrared studies (15) the structure proposed is shown in Figure 2, since bands are found a t 2020, 2043, and 1833 cm-I. Cotton (15) has found an additional band a t 1997 cm-' using LiF opt,ics, which indicate that Sheline's structure is incorrect. The proof of this structure is thus still open to discussion. The recent work of Cotton and Monchamp (21) in regard to the carbonyls of cobalt [ C O ~ ( C Oand ) ~ Cod(C0)12]is of particular importance. Original work with Co2(C0j8 had shown bands a t 2034, 2054, 2077, and 1858 cm-I; from this it was inferred that both "end" and "bridging" carhonyls existed in the molecule, and that the structure proposed (Fig. 3) was that of two trigonal bipyramids joined a t an edge (17, 18). Using calcium fluoride optics, Cotton, et al. ( P I ) , found bands a t 2075, 2064, 2047, 2035, 2028, 1867, and 1859 cm-'.

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Structure of Complex Inorganic Molecules

Infrared spectroscopy has assisted in the determination of the structure of certain complex inorganic molecules, such as the metal carbonyls, inter-halogen compounds, boron hydrides, and nitrogen oxides. The practice used has been to Compare the observed spectrum with the spectrum expected for an assumed model, on the basis of group-theory-derived selection rules. The dependability of this procedure depends on the degree of reliability of the observed spectrum. It is not at all satisfactory to use this method if the observed spectrum is poorly resolved. In addition, failure to observe as many hands as are required by the selection rules is no basis to reject a certain model. The metallic carbonyls illustrate this point well. Metallic Carbonyls. Sheline (14-16) has examined the infrared spectra of the iron carbonyls, Fer(CO)s, Fea(C0)12, and Fe(C0)6. The carbonyls of cobalt. (17, IS), C O ~ ( C O )and ~ , CO~(CO),~, have also been studied. X-ray diffract,ion data (19) has indicated that in some of t,hese carbonyls there are two kinds of carbonyl groups; the "bridging" type and the "end" type. Infrared data of some of these compounds has shown absorptions in the 2000 cm-' region, and this has been attributed to the "end" carbonyls; while 202

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Journd of Chemicol Educotion

Figure 1. FeKOIe.

Proposed structure

Figure 2. FeJCOItb

Proposed structure for

Figure 3. Proposed structure for Coz(COls, D2hsymmetry, trigonal bipyramid (17).

The appearance of five terminal fundamental frequencies and two bridging bands ruled out all of the previously proposed structures. In a similar fashion the structure proposed for CO~(CO),~ was found to be in error (21). The new structure proposed has been substantiated by X-ray studies (Fig. 4) (22). The manganese and rhenium compounds [M2(CO)lo] do not possess bands in the 180Q-1900 cm-I region and are believed, as a result, not to contain a bridged structure (25). X-ray diffraction data has recently confirmed this (24). The interpretation of the structures of the simpler carbonyls was also aided by infrared spectroscopy. Chromium, molybdenum, and tungsten form hexacarbonyls of the type M(CO)s. Electron diffraction and X-ray studies have indicated a regular octahedral structure (25, 26). Infrared measurements have confirmed this. The infrared spectrum of Fe(CO)s has shown that the compound has a trigonal bipyramid structure (16). The assignment of the observed frequencies of the gas was successfully made on the basis of a trigonal bipyramid model (Fig. 5). Recently (27), the stereochemistry of five coordinate compounds of the type IJFe(C0)4,LFe(CO)r where L is PhaP, Me.NC, Et.NC,

type, then the selection rules predict eight infrared vi-

brations. Aluminum borohydride, Al(BH& was also investigated by infrared and results (50) were interpreted on the basis that the structure has hydrogen bridges as in diborane. Oxides of Nitrogen. The establishment of structures for the various oxides of nitrogen has also been helped by infrared measurements. Nitrous oxide was shown to be linear (5135).

Nitrogen tetroxide has been studied in the solid, liquid, and the vapor state and has led Sutherland (54) to predict a planar structure. Recent confirmation of this was made by Snyder and Hisatsine (35). Nitrogen pentoxide was shown to have the struct,ure in the liquid state (36).

Coordindion Compounds

Figure 5. FelCOls.

Proposed structure for

Ph.iYC, Bu.n'C, and (PhKC)&o+' has been examined. By means of the infrared spectra and comparison with the selection rules, the configurations were concluded to be very probably as in Fe(CO)s, trigonal bipyramid. Interhalogen Compounds. Several of these compounds have been investigated by the infrared method and structures elucidated. The compound IFs illustrates how Raman and infrared spectra can together aid in structure determination. The Raman spectrum of the liquid (28) and the infrared spectrum of the gas (28) were measured. Lord, et al., compared t.heir observations with the predicted number of frequencies for seven possible symmetries for the molecule (Table 1). Structures 5, 6, and 7 are immediately eliminated, as are structures 1 and 2. Thus, the structure of IFs is either 3 or 4, (4 being a tetragonal pyramid) providing the spectrum is highly resolved. In a similar way the structure of IF, was postulated to be a pentagonal bipyramid. Boron Hydrides. Infrared examination of diborane (B2HSshowed that it had eight bands which appear to be fundamentals (29). If the structure is of a bridge

Infrared spectroscopy has been a very useful tool in the field of coordination chemistry. I t has aided in the problems of structure determination, cis-trans isomerism, metal-ligand sit,e at,tachment, and in determining bond types in these complex molecules. Only several examples in which the infrared spect,ra of coordination compounds have been instrumental in solving the aforementioned problems will be presented. Numerous others are known, and may be found in the literature. Certainly a large effort in this field has been made by the group a t Notre Dame, including Quagliano and Mizushima. The structural isomerism arising in the complex ions, nitropent.amminecobalt(II1) [Co(NH&N02]++, and nitritopentamminecobalt (111) [CO(NH~)~ONO] ++ has been reported' (37). From analogy with the colors of compounds known to possess M-0 or M-N links, the stable yellow-brown form was considered to have the C e N link, while the unstable red form was considered to have a Co-0 link. The verification of these postulates has now been accomplished by infrared techniques (58). The infrared assignments Table 1.

Application of the Vibrational Selection Rules to the Possible Structure of IF6 (28)

Total Struc- Point ture group 1 2

Djh D8h

Cr, 4 CI, 5 Ca, 6 C*, 7 C. Observed

3

fie. quenctes

7 8 7 9

8

Raman Spectrn3 Total 3 6

7

I)

8

Infra,,d Polar- speci ~ e d trum

Coincidences

PolarTotal

ized

1

3

0

0

2

5

3

2 3

4

4

0 2

6 8 11 12

6 8 11 12 3

4

12

12

5

12 9

12 8

8 3

4

Volume 38, Number 4, April 1961

3 4

5 8 2

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203

made for these isomers are shown in Table 2.

Organometallic Compounds

Salts of Carboxylic Acids. The salts of the carboxylic acids have been extensively studied by Lecompte Toble 2. Differences in Infrared Absorption of the Ions (42-45). When dealing with salts of acids, the charac[CO(NH~)~NO~I ? and ICO(NH,)~ONO 1. ' teristic carbonyl absorption is lost and is replaced by Cm-1 Assignment two bands between 1550-1610 cm.-' and between 1300[Co(NHa)sNOnl 1430 NO, (st.) antisym. 1400 cm-' corresponding to the antisymmetrical and 1315 NO2 (st.) sym., and NH3 def. sym. symmetrical vibrations of the COO- group. 825 NOa (def.) I n the solid state, the two frequencies vary with the (st.) sntisym. lCo(NH~),ONOI+Z 1460 0-N-O 1065 0-N-O (st.) sym. nature of the metallic ion. For mono- and divalent elements there is a linear relationship hetweeu the electronegativity of the element and the asymmetric stretching frequency (46). Stimson (47) has also This was done by observing the growth of peaks a t noted similar dependencies in salts of substituted 1430 em-', 1315 em-', and 825 cm-', with a decrease benzoic acids. initially in intensities of the absorptions at 1460 cm-' and 1065 cm-' as freshly prepared red [CO(P*'R~)~O- Salts of Phosphorus Oq-Acids. The spectra of a number of inorganic salts of the phosphorus oxy-acids NO]C12aged. Eventually the peaks at 1460 cm-I and were measured in the NaCl region (IS, 48). The 1065 cm-' disappeared, and the spectrum became that absorptions of each type of aniou were characteristic, of t,he yellow brown [ C O ( N H ~ ) ~ N O ~ ] C ~ ~ . and correlations (48) were suggested for various strucInfrared spectroscopy has proved helpful in detertural groupings in these substances. The hydrated mining the site of attachment of the metal to the ligand salts all absorb in the OH region of water a t 3300 and ($9). I n the coordination compounds involving urea 1640 cm-'. I n addition, the spectra are rich iu the and thiourea as ligands, the problem involved is whether regions of 90C-1400 cm-', and these are attributed to one has a metal-nitrogen bond or a metal-oxygen P-0 associated vibrations (e.g., P=O stretching, bond, in the case of urea. The infrared spectra of ionic POO-, and P-0-P). these complexes showed that urea formed metalMetal Complexes of Acetylacetone. A number of nitrogen bonds with Pt(I1) and Pd(II), and metalmetal complexes of acetylacetone have been prepared oxygen bonds with Cr(III), Fe(III), Zn(II), and Cuand studied by infrared methods (49-56). The in(11). These conclusions were reached by a comparison frared results with acetylacetone are consistent with a of the spectra of the complexes with that of free urea. conjugated chelate structure resonating betweeu forms If the coordination involves M-0 bonds the specI and I1 on t,rum would be expected to differ only slightly from that of urea. The C=O absorption would shift toward lower frequency [the C=O absorption region a t 1700 cm-I is blank in the Cr(III), Fe(III), Zn(II), and Cu(I1) compounds, while in free urea it is a t 1683 cm-')I. If metal-nitrogen bonds are involved, then the spectrum is significantly different from that of the free urea molecule. The XH stretching, NH deformation, and CN vibrations would be expected to shift toward lower frequencies, as indeed they do in the I t would be expected that coordinatiou to a metal (111) Pt(I1) and Pd(I1) compounds. I n addition, the abwould shift the carbonyl frequency to lower values; sorption due to the carbouyl group is shifted to 1720 shift the C-H stretching absorption to higher freem-', because of the blocking of resonance between the bonded nitrogen and the C=O group, as one would M expect. Similar comparisons of the spectra of the metalthiourea complexes with that of free thiourea showed that all metal bonds were to sulfur (40). The spectra were all similar t,o the spectrum of thiourea, with the exception that the C=S frequency was shifted to lower values. Infrared spectroscopy has aided in distinguishing yuency (new environment more like benzene riug); between cis and trans isomerism in inorganic cothe antisymmetric and symmetric methyl stretching ordinated compounds. For example, cis- and transvibrations should remain about where they are in dinitrotetraammine cobalt(II1) chloride, [Co(NH& acetylacetone; the OH . . . . 0 absorption should dis(XOr)l]CI, can readily be distinguished by their specappear. All of these expectations have been realized tra (41). Fewer absorption peaks were present in the in the infrared spectra of these compounds. Effort,s to trans compound than in the cis. This can be a direct correlate the shift of the carbonyl frequency with the consequence of the selection rules, since the trans comstrength of the complex also have been made (56,66). plex has a cender of symmetry while the cis complex does not; therefore the trans compound might be exAn excellent review on the "Infrared Spectre of Transitional pected to give a less rich spectrum. However, only Metd Complexes" by F. A. COTTONhas recently appeared in one example was given. I t would be of interest if the book "Modern Coordination Chemistry" by J. LEWISAND other cis-trans isomers gave t,he same results.' R. G . W r ~ ~ r wInterscience s, Publishem Ino., New York, 1964. 204 / Journal of Chemical Education

Phosphorus Compounds

The industrial uses of phosphorus compounds in recent years have grown rapidly. Phosphorus compounds are used in the oil industry as additives, in detergents and insecticides, in food manufacture, in water treatment, and as solveut extractants in ore processing. I t is in the field of solvent extract,ion that many of the organophosphorus compounds have made valuable contributions recently. The infrared study of these compounds was started by Daasch and Smith (57). They puhlished spectra of 60 organophosphorus derivatives and proposed spectra-structure correlations. Bellamy and Beecher (58) extended these studies, as did others (48, 69-66). Most of this work was made on the neutral type of est,ers. The acidic type of esters has gained more attention in the past five years, with the groups at Argonne, Oak Ridge, and Harwell laboratories evaluating these acids as solvent extractants for various cations. In the course of this work many new acidic type organophosphorus compounds have been synthesized and subjected to infrared studies (67). The infrared studies of the acidic type of organophosphorus esters of the type (GO),POOH, G'OGPOOH (where G can he alkyl, aryl, or variant thereof) have been reported (67). These acids have been found to be diieric in most solvents, and intermolecular hydrogen bonding has been postulated. Despite the fact that these compounds are quite complex, the infrared spectra can be qualitatively analyzed in terms of vibrations involving the phosphorus atom and those involving the G group. The spectra are characterized by broad regions of absorption at about 2500-2700 cm-I and about 2300-2350 cm-' (67, 68); and these have been attributed to the bonded P-OH st,retching motion (57,668). In addition, a third, broad region appears at about 1680 cm-', which is attributed to the bonded P-OH deformation motion (69, 70). These regions all disappear on salt or metal complex formation, and are appropriately shifted to lower frequencies on denteration. The P 0 stretching vibration is very strong and is observed in the 1200 cm-' region. The identification of this absorption is aided by salt or complex formation. The strongest absorption appears at the 1000 em-' region and is prohably due to several vibrations P-O-(G) and (P)&G. Where G is aryl, the P-&(G) vibration (asymmetric) appears at ahout 9G6-1010 em-' (64). Where G is alkyl, it is found at about. 1000-1050 cm-'. I n aryl phosphorus acids the vibration a t 1180 em-' is attributed to the (P)-Gary1 stretching motion. The various CH vibrations are located in the expected positions of the spectra (CH stretching, 280W3100 cm-'; CH deformation, 1380-1480 em-'; C=C skeletals in the aryl ring 159Cb1610 cm-I, and 15001518 em-'; CH out-of-plane vibrations in aryl phosphorus acids, 700-835 cm-I).

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Differentiating Between Ionic and Covalent Bonds

The use of infrared spectroscopy in differentiating between an ionic bond and a covalent bond in a complex (in the ligand-metal bond), where the ligand is nitrate, carbonate, thiocyanat,e, and cyanide has been reported. The method has been particularly useful in the case of

nitrates aud carbonates, while not so successful in the case of thiocyanates and cyanides. Nitrates. The nitrate ion is considered to have the configuration of an equilateral triangle possessing the point group Dlh symmetry, In nit,rate compounds in which the nitrate is covalently hound the symmetry lowers, and is changed to that of a point group CZv symmetry (71). Thus in going from a Dlh symmetry to one of CZrthe p3 frequency in the symmetrical NOiion (point group Dlh) undergoes splitting into two new frequencies. One component is observed to shift toward higher frequency k4) a t about 1500 cm-', and the other toward lower frequency (Y1) a t about 1275 cm-l, both of point group CSusymmetry. I n addition, a t about 1000 em-' appears, as does the frequency a t about 800 cm-I. Gatehouse (7Z, 73) and Quagliano (74, 75) have recently examined the infrared spectra of several nitratocomplexes of the transition elements, and observed the changes in symmetry occurring when the nitrate is in and out of the coordination sphere of the cat,ion. They have been able to distinguish between the nitrate ion and a covalently bonded nitrate group in these compounds. Addision and Gatehouse (76, 77) have also examined several anhydrous nitrates in this way. Peppard and Ferraro (78) have applied this method to the study of the organophosphorus complexes of nitrates. They have examined the infrared spectra of the nitrate-bis-(2-ethylhexyl)phosphate complexes of M(V1) cations, and have noted the lowered symmetry in these complexes. Ferraro (79) studied the nature of nitrate in the TBP (tri-n-butyl phosphate) solvated nitrates [M(III), M(IV), and M02(II) type], and observed a greater lowering in nitrate symmetry in these complexes than is present in the hydrated nitrates themselves. The nature of the mineral acid (H1\T03) extracted into HDEHP (his-(2-ethylhexyl) phosphoric acid) and TBP was also studied by this method (80). The infrared spectra of several metallic nitrates were examined @I), and the observation made that in going from a monovalent metallic nitrate to a tetravalent metallic nitrate there is an increased lowering of the nitrate symmetry, and a t,ransition from a point group Dab symmetry to a point group Cy, symmetry. The observed changes in t,he nitrate symmetry in inorganic nitrate compounds has proved to he very useful in giving additional informat,ion as t,o the nature of the nitrate present. Carbonates. As with the nitrates, Gatehouse, et al. (823 have attribut,ed the decrease in symmetry arising in carbonates to covalent bonding. The carhonate ion has a Dl,. symmetry giving rise t,o four vibrations. When the carhonate ion is covalent,ly bouud bhrough one or two oxygens the symmetry is lowered to that of a point group C2,. The v , vibration in the carbonat,e ion is in a doubly degenerate state and is split into two component,^ when the carhonate ion is coordinat,ed, v4 aud ul. The following compounds illustrate the results obtained for a covalent carbonate. v1

[Co en Coal Cl

1577s

Dimethyl carbonate

1760s

YI

YZ

ye

or

vg

12818 1059w 1272s 1035w 830s 1280s 969s 793s

us

754m

The basic carbonates and hydrogen carbonates also Volume 38, Number 4, April 1961

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205

Table

3. Frequency Ranges of Carbonate Absorption Bonds (82)

Type of Carbonate Simple Basic Complex

Acid Organic

show this splitting. Table 3 illustrates the frequency range of the carhonate absorption bands.

(30&700 cm-'), and in the sodium chloride region (700-3600 em-'), in Table 4.

Far Infrared Spectra

The Infrared Speclm of Metal-Oxygen Double Bonds

A recent publication (85) reported on the infrared spectra of over 200 inorganic compounds in the cesium bromide region. Characteristic frequencies are listed for some inorganic ions in the cesium bromide region

The infrared spectra of several compounds containing metal-oxygen double bonds has been reported (84). The results show that in compounds where one metaloxygen double bonded link exists, an absorption occurs

Table 4.

(300-3600 cm-') (8, 83)

Characteristic Frequencies of Polyatornic Inorganic Ions

1100

1:

7 7;

co j HCO;

siog SILICATES NO; NO; NH; POI

S -

HPo;

DIFFUSE

H2PO;

ALL 6/11

so; so;

s,og seo;

sea; clo; 00; BrO;

5

vo;

croi Cf20>

Mood WOf MnO:

11W 1600

1800

C: HCOj OCN-

5

-,

** Literature value. S, strong; M, medium; W, weak; SP, harp. Journal of Chemicol Education

2400

1

ma

2600

1 1300

3400

w' I1 II

L *

* In most, but not all, examples.

/

2200

-S

CN'

SCNN: NH; P: HPO; WATER OF CRYST.

moo

BANDS

SHM7F

10;

206

I

s

HSO; S20f S20f

W

II II II II

M

M

1

a t 952-1087 cm-'. The compounds studied in this class were VOSOr, VO(acetylacetone), Ti(acety1acetone)%,K2MoOC15, K2CrOC16, and VOC1,. I n compounds where two oxygens are linked to the metal (e.g., Cr02C12and Cr02F2),the absorption occurs a t 9701016 cm-'. I n compounds where there are more than two metal-oxygen bonds (Vz06, CrOa, MoOa, KMnO4, KRe04, KZCr04,and KtCrzOr), the absorption occurs a t 825-1020 cm-'. The absorptions in this region are of diagnostic value if no atom of similar atomic weight is present which might give frequencies in the same region.

COTION,F. A.,

Summary

CH~DIN J.,, Compt. rend., 201, 552 (1935); Compt. ?end.,

Several examples have been discussed, which serve to illustrate the application of infrared spectroscopy to inorganic chemistry. With the availability of commercial high resolution infrared spectrophotometers, interest in inorganic infrared spectroscopy should continue to develop. The method is very useful and is certainly not rest,ricted to organic compounds.

MELLOR,J. W., "A Comprehensive Treatise on Inorganic and Theoretical Chemistry," Vol. 14, 1st ed., Longmans Green & Co.. New York. 1942. PENLAND, R. B., LANE,T. J., AND QUAGLIANO, J . V., J . Am. Chem. Soe., 78, 887 (1956). PENLAND, R. B., ET AL.,J . Am. Ckem. Sac., 79,1575 (1957). YAMAGUCHI, .4., ET AL., J . Am. Chem. Soc., 80, 527 (1958). F A U ~ TJ., PHILIP,AND QUAGLIANO, J. V., J. Bm. Ckem. Soc., 76,5346(1954). LECOMPTE, J., Rev. optipue, 28, 353 (1949). D u v a , C., LECOMPTE, J., AND DOUYILLB, F., Ann. Phys.,

AND

PARISH.R. V.. J . Chem. Soc.. 1440

(1960).

.

LORD.R. C.. ETAL..J . Am. Ckem. Soc.., 72.522 . (1950).. BELL,R. P., AND LONGUET-HIGGINS, H. C., PTOC. Roy. SOC. (London) 183A, 351 (1945). PRICE,W. C., J. Chem. Phys., 17,1044 (1949). PLYLER,E. K., A N D BARKER,E. F., P k y s Rev., 38, 1827 (1931): Phys. Rev.,41, 369 (1932). ~ICHARDSON, W. S., AND WILSON,E. B., J . Ckem. Phys., 18,694 (1950).

BEGUN,G. M.,

AND

FLETCHER, W. H., J . Ckem. P ~ K s 28, .,

414 (1958).

SUTHERLAND, G. B. B. M., PTOC.Roy. SOC. (London), 141A, 342 (1933).

SNYDER, R. G., A N D HISATSINE, I. C., J . Mo1. Speetro~eopy, 1,139 (1957).

Literature Cited

SCHAEFER, c . , AND MATOSSI, F., "Dm U ~ ~ I RSpektmm," TO~ Springer, Berlin, 1930. SCHEIDT, V., Z. Natu~forsek.,76, (5), 270 (1952). STIMSON, M. J., J . Am. Ckem. Soc., 74, 1805 (1952). STIMSON, M. J., Pkw. Radium, 15, 390 (1954). PIMENTEL, G. C., Spee. Ada, 12, 94 (1957). LECOMPTE, J., Anal. Ckim. Aeta, 2, 727 (1948). LECOMPTE, J., Cahien Phw., 17, 1 (1943). M~LLER, F., A N D WILKINS,C. H., Anal. Chem., 24, 1253 (1952).

MELLON,M. G., 'iAn~lytie~l Absorption Spectroscopy," 1st ed., John Wiley & Sons, Inc., New York, 1950. BARNES, R. B., GORE,R. L., LIDDELL, U., AND WILLIAMS, V. Z., "Infrared Spectroscopy, Industrial Applications and Bibliography," 1st ed., Reinhald Publishing Corp., New York, 1944. HITNT,J. M., WISHERD,M. P., A N D BONHAM, L. D., Anal. Chem.. 22.147811950). . . (12) HUNT,j.hi., AND TURNER, D. s., ~ n Chem., a ~ 25, 1169 (1953).

203,722(1936).

17.5 (1942).

DOU~ILL*, F., DUVAL,C., AND IIECOMPTE, J., Bull. 80c. ckim. Memoi~es,9,548 (1942). DUVAL,C., GERDING, H., A N D LECOMPTE, J., Ree. Trau. Chim., 69,391 (1950). KAGARISE, J., J . Pkj,~.Ckem., 59, 271 (1955). STIMSON, M. J., J . Chem. Phys., 22, 1942 (1954). CORBRIDQE, D. E. C., AND LOWE,E. J., J . Chem. Soe., 4555 (1954), 493 (1954).

BELFORD,R. I,., MARTELL,A. E., AND CALVIN,M. J., J . Inory. and Nuel. Chem., 2, 11 (1956). BELLAMY, L. J., AND BRANCH, R. F., J . Ckem. Soc., 4491 (1954).

MORGAN, H. W., US AEC Document No. 2659 (1949). LECOMFTE, J., D ~ C Z L S Fa~aday S ~ O ~ SSoe., 9, 125 (1950). BELLAMY, L. J., SPICER,G. S., AND STRICKLAND, J . D. H., J . Ckem. Soe., 4653 (1952). HOLTZCLAW, H. F., JR., AND COLLMAN, J. P., J . Am. Ckem. Soe., 79, 3318 (1957). D u v a ~ ,C., FREYMANN, R.,A N D LECOMPTE, J., Bull. me. Ckim. Fr., 19,106(1952). WEST, R., A N D RILEY,R., J. Inorg. and Nucl. Chem., 5, 205 - ..- (19.5RI - -- ,. DAASCH, I,. W., AND SMITH,D. C., '4nal. Ckem., 23, 863 \

COTTON,F. A.,

AND

WILKINSON, G., J . Am. C k m . Soe..

79,753 (1957).

PITZER,K. S., J . Am. C k m . Soc., 72. 1107 (1950): FATELEY. W. G.. AND LIPPINCOTT. L. R.; Spee. A&, i 0 , 8 (1957): (17) CABLE,J . W., NYHOLM, R. S., AND SHELINE, R. K., J . Am. Ckem. Soe., 76,3373 (1954); CAELE,J . W., AND SHELINE, R. K., J . Am. Ckem. Soe., 76, 3373 (1954); CABLE,J. R. K., C h m . Revs., 56, 1 (1956). W., AND SHELINE, R. A., ET AL.,J . Am. Ckem. SOC.,77,3951 (1955). (18) FRIEDEL, (19) POWELL, H. M., A N D EWENS.R. V. G.,J . Ckem. Sac., 286 SHELINE,R. K.,

AND

(1939). (20) STERNRERG, H. W., ET AL.,J . Am. Chem. Soe., 76, 1457 (1954). . . (21) COTTON, F. A., AND MONCHAMT, R. R., J . Ckem. Soe., 1882 (1960); BOR, G., AND MARKO,L., Spec. Ada., 9, 747 (19591. \-.--,(22) CORRADINI, P., A N D ERCOLI,R., Abstracts, 17th Inter(23) (24) (25) (261

national Congress, Pure and Applied Chem., Munich, 1959. Vol. I, p. 25. BRIMM,E. O., LYNCH,M. A., AND LESNY,W. J., J . Am. Chem. Soc., 76,3831 (1954). DAHL,L. F., ISHISHI, E., AND RUNDLE,R. D., J . Chem. Pkus.. " , 26.1750 (1957). , BROCKWAY, L. O., EWENS,R. V. G., AND LISTER,M. W., Trans.Faraday Soe., 34, 1350 (1938). RUDORF, W., AND HOFMANN, V. Z., Z.Pkysik Chem., 28B,

.

351 (1935).

.

(1951).

BELLAMY, L. J.,

AND

BEECHER,L., J. Chem. Soe., 475

(1952); 1701 (1952); 728(1953).

GORE,R. C., Discussions Faraday Soe., 9, 138 (1950). MEYR~K C., , AND THOMPSON, H. W., J . Chem. Soc., 225 (1950).

HOLMSTEDT, B.,

AND

LARRSON, L., .4ela C k m . Scand., 5,

1179 (1951).

BERGMANN, E. D., HAUER,V. Z. L., A N D PINCHAR, S., J . Ckem. Soe., 847(1952). MORTIMER, F. S., Spee. Ada, 9, 270 (1957). NYQUIST. , R. A.. J . Anal. Chem.. No. 4. 161 (1957). . . PUSTINGER, J. V., JR., CAVE,W. T., AND NIELSEN,M. L., J . Appl. Ckem., N o . 11, 909 (1959). M c I v o ~R. , A., ET AL., Can. J . Chem., 36,820 (1958). PEPPARD,D. F., FERRARO, J . R., AND MASON,G. W., J . Inoro. and Nuel. Ckem.. 4. 371 11957): ., J . Inorn. and Nuei Chem., 7, 231 (1958); J . Inorg. and Nuel. ?hem.,

..

.

12,60(1959).

BRAUNHOLTZ, J. T., ET AL.,J . Ckem. Soe., 868 (1959). THOMAS, L. C., Ckem. andlndustry, 197 (1957). MAARSEN, J . A,, S ~ TM. , E., AND MATZE,J., Rev. tmu. Ckim., 76,713 (1957). HERZBERG, G., "Molecular Spectra. and Molecular Stnrcture," Vol. 11: "Infrared and Rrtman Spectra of Polyatomic Molecules," 1st ed., D. van Nostrand Ca., Inc., New York, 1945, p. 178.. Volume 38, Number 4, April

1961

/

207

(72) GATEHOUSE, B. M., LIVINGSTONE, S. E., AND NYHOLM, R. S., J . Chem. Soe., 4222(1957). (73) GATEH0U8Es ,...B. M . t AND C O M y ~ s , J . ~ J . Chenz. See., 3965 %. .

(LVJS,.

(74) M ~ z u s a r S., ~ ~AND , QUAGLIANO, J. V., J . Am. Chem. Sac., 75,4870(1953). E. P., ET AL.,J. Am. Chem. Soe., 81,3818 (1959). (75) BERTIN, (76) ADDIBON, C. C., AXD GATEHOUSE, B. M., Chem. and Ind., 464 (1958). , (77) ADDISON, C. C., AND GATEHOUSE, B. M., J. Chem. Soc., 613 (1960).

.

208

/

Journol of Chemical Education

(78) PEPPARD, D. F., AND FERRARO, J. R., J. rnwg. and N u d . Chem., 10, 275 (1959). (79) FERRARO, J. R., J. Ino~g.and Nucl. Chem., 10,319 (1959). (80) FERRARO, J. R., J . Inwg. and Nucl. Chem., 15, 365 (1960). FERRARO, J, R,, J , MO1, 2709copy, 4, 99 (19 (82) GATEHOUSE, B. M., LIVINGSTONE, S. E., AND NYHOLII, R. S., J . Chem. Soc., 3137 (1958). (83) MILLER, .4eta,. 16,135 (1960). . F.,. ET AL.,. Spec. . . . . (84) BARRACLOUGH, C. G., LEWIS,J., AND NYHOLM, R,S., J. Chem. Soe., 3552 (1959).