1.

COPPER (II) AND NICKEL( II) COMPLEXES OF 2-ALKOXYPHSNQLS

A Thesis presented to the University of* Surrey for the degree of Doctor of Philosophy in the Faculty of Science,

by

Stanley Leonard Jones

Tine Inorganic Research Laboratories, June, 1969*Department of Chemistry,University of Surrey,London, S 0¥„ll

ProQuest Number: 10800192

All rights reserved

INFORMATION TO ALL USERS The qua lity of this reproduction is d e p e n d e n t upon the qua lity of the copy subm itted.

In the unlikely e ve n t that the au tho r did not send a co m p le te m anuscrip t and there are missing pages, these will be no ted . Also, if m ateria l had to be rem oved,

a no te will ind ica te the de le tion .

uestProQuest 10800192

Published by ProQuest LLC(2018). C opyrigh t of the Dissertation is held by the Author.

All rights reserved.This work is protected aga inst unauthorized copying under Title 17, United States C o de

M icroform Edition © ProQuest LLC.

ProQuest LLC.789 East Eisenhower Parkway

P.O. Box 1346 Ann Arbor, Ml 4 81 06 - 1346

Summary

New copper(II) and nickel(II) complexes with 2-alkoxy- phenols have been prepared.

The visible spectra of the copper(II) complexes indicated a quasi-tetrahedral stereochemistry for two of them, and a distorted octahedral stereochemistry for the remainder. Magnetic susceptibilities for most of the copper(II) complexes have been determined over the range of temperature 80~3Q0°K, and the anhydrous complexes showed a variable degree of antiferromagnetic interaction. No unambiguous theoretical interpretation of the magnetic data as a function of the degree of aggregation was achieved, but in the case of bis(2~methoxyphenolato)copper(II), which has a quasi-tetrahedral stereochemistry as indicated by its visible spectrum, the very existence of antiferromagnetism pointed to a polymeric linear chain structure.

Unequivocal infra-red evidence for the participation of the alkoxy group in coordination was not found, but the spectra of some Of the nickel(II) complexes showed tentative evidence of methoxy group coordination. Because of the uncertainty of the bonding role of the alkoxy group, the number of possible structures for the complexes was increased.

The visible spectra and magnetic properties of the nickel(II) complexes were consistent with an octahedral stereochemistry. A comparison of the co|5per(II) and nickel(II) series of complexes is given.

Ac knowledgements

This work was carried out in the Chemistry Department of the University of Surrey during the period X$66-1969* The research topic was initially proposed, and financed by Unilever Limited, for which the author wishes to express his gratitude„

The encouragement and criticism given by Dr. J.I. Bullock during the course of this research has heen greatly appreciated.

C O N T E N T S

Page No,

-i. 0 jL 0

1.2.

Summary ... ... ... ... ...Acknowledgements.„ ... ... ... .Index of Tables and Figures ... ...

Section 1. Introduction ... ... ... ... .Historical Introduction ... ...Factors controlling the choice of ligand and method of preparation ... ... .

1.3. Reason Tor undertaking this research work •2.4. Nickel (II) complexes of 2 - a ikoxyph e n d s

Section 2. Physical methods employed in structuredetermination ... ... .Infra-red Spectroscopy.. .Magnetism.. ... ... .Visible Spectra.. ... .

Section 3. Experimental ... ...3.1. Ligands and Reagents ...

, Metal Analysis... ... ., Microanalysis „. „ ...

3.4. Stanton Thermobalance...3.5. Infra-red and Visible Spectra3.6. Magnetic Measurements... .3.7. Mass Spectra ... ...3.8. Microanalytical Results

2.1

2.3

3.23.3,

9. Preparation of Complexes

jl±

11

12131524373939424246464950 53

/contd.

Contents (contd.) Page No

Section 4. Copper(II) Complexes,, Results andDiscussion ... ... ... ... ... 6l

4.1. The Infra-red Spectra ... ... ... 694.2. The Diffuse Reflectance Spectra ... ... 734.3. The Magnetic Properties ... ... ... 754.4. Discussion ... ... ... ... ... 82

Section 5. Nickel(II) Complexes. Results,Discussion and Comparison ’with Copper(II). 88 The Infra-red Spectra.. ... ... ... 96The Diffuse Reflectance Spectra and Magnetism ... ... ... ... ... 99Discussion and Comparison... ... ... 104Concluding Remarks... ... ... ... 106

__________ Appendices ... ... ... ... ... 1076.1. Appendix 1. Magnetism ... ... ... 1086.2. Appendix 2. Infra-red ... ... ... m

Section 7. References ... ... ... ... ... 112

5.1.5.2.

5.3.e; h.

Section 6

■oOo-

Index of Tables and Figures

Tables

Number Page1. .. .0 . . 262 28 3.............. 33 4 .............. 385. .. .0 506. .. .. 60

Figures

Number Page1....... 252. .. .. .. 29

3.............. 31 4 .............. ZtO5. ............416. ............457........ 52

8. .. .. .. 57

9.............. 6210. .. .. .. 6511. O• « « 0 9 68

Number7.8.9.

10.11.

Sags.

7677 97100102

Number Page. 12...............70

13. .. .. 7114. ............ 72

> 15...............7916. .. ... «» 8017 . ............ 8918. .. .. . . 9 219.............. 9520. .......... 10121. .. .. ..103

Introduction

l.X. Historical Introduction

Although a great deal of attention has been given to complexes of copper(II) with ligands of the following type:

(where X is -CHO, ; - C =s A / 2* ; - H - R, ; - NHg, ;etc.); no work has been done with 2~alkoxyphenols. A cobalt(II)complex with guaiacol, X = -OMe, has been previously reported

( 5)by Cotton and Holm. Unfortunately, the described method of preparation failed when attempted in these laboratories. This has been rectified in light of the experience gained with the copper(II) system.

1*2. Factors controlling the choice of ligand and method of preparation

For complexation to occur the 2~alkoxyphenol has to be ionised. This was accomplished by reaction with alkali. It is well known that phenols aire susceptible to oxidation by air in neutral and alkaline solution, for example, an aqueoussolution of sodium 2-meihoxyphenolaie slowly darkens on standing in the atmosphere. The tendency for an aqueous solution of a sodium or potassium 2-alkoxyphenolaie to undergo aerial oxidation can be substantially reduced, if electron-withdrawing substituents such as -NO , and -CHO are introduced into the 4-position. Providing the sodium or potassium 2-alkoxy- phenolate solution was stable, the introduction of copper(XI)

into the system was never found to catalyse oxidation. The reaction of copper(II) with 2-alkoxyphenols can be written as:

2ML + Cu+2 = Cu L2 + 2M+,

where ML is the potassium or sodium salt of the 2-alkoxyphenol.A water milieu was always employed, and the copper(IX) saltsolution always added to the solution of the ionised ligand.If, however, sodium or potassium hydroxide was added to asolution containing a 2:1 mole ratio of ligand and copper(IX),complex formation did not occur, instead basic copper(IX)compounds were precipitated. This does not imply that theproblem of basic copper(XX) compound formation cannot arisein the former case, indeed it does if the 2-aIkoxypheno1 isa sufficiently weak acid, since the degree of base hydrolysisincreases with the decreased acidity of the 2-alkoxyphenol.The problem of unfavourable base hydrolysis can be controlledto some extent by using concentrated solutions and an excessof free ligand. The work of Dolique and Mestres illustrates

( 7)this point well. These workers studied the system copper(XI) -phenol, X = H.

In all the complex preparations described in this thesis a molar bias has been applied. Ideally we react LH (2 mole) and MOH (2 mole) irith copper(II) (1 mole), in practice a bias has been applied as follows:

LH MOH Copper (XI) .

The molar bias has ranged from 3% to 10%.With one exception, the complexes that have been

prepared cannot be washed once filtered. Washing with water, aqueous-organic, or organic solvents led to decomposition with the formation of basic copper(II) materials. Therefore in order to achieve a satisfactory level of complex purity, particular attention has had to be paid to preparative condition

Ligands

The ligands which we have considered are listed below.

Structure Name Trivial Name

2-me thoxy-4-formylpheno1 Vanillin

2-methoxy-4-nitrophenol 4-nitroguaiacol

2-methoxyphenbl Guaiacol

2-ethoxy-4-formylphenol Bourbonal

HO2 ,6-dimethoxy-4-formyl- phenol

Syringaldehyde

1.3. Reason for undertaking this research work

Copper(IX) compounds and metallic copper are widely employed as fungicides. At the present time the search for suitable fungicides for the protection of wood is still being actively pursued. The xylem cells of trees contain lignin, one of the oldest natural organic polymeric materials on the evolutionary scale. The occurrence of the structural entity

'Me

in lignin and the current interest in the development of coppe based fungicides for wood, led us to investigate the systems described in this thesis.

1.4. Nickel(XI) complexes of 2~alkoxyphenols

In order to extend our knowledge of the reactions of these ligands with the first-row transition metal ions, the reactions of 2-alkoxyphenols with nickel(XI) were investigated

Precisely the same considerations as discussed above under the heading, *Factors controlling the choice of ligand and preparative method1, hold.

Physical Methods Employed in Structure Determination

2.1. Infra-red Spectroscopy/ O \Prout et al. have unequivocally shown by X-ray

crystallographic analysis that the methoxy group indiaquobis(methoxyacetato)copper(II) is coordinated to copper(II)

(9)through oxygen. Similarly, Orioli and Di Vaira have shownthat the niethoxy group in dibromo-[1~(o-methoxyphenyl)-2, 6--diazaocatane]nickel(II) is involved in coordination.

Little has been done to establish the participationof an alkoxy group in coordination from consideration ofinfra-red spectroscopy. Some effort has been made in thisarea by B u t c h e r / w h o considered complexes of nickel(II)

(11)with ortho-methoxyaniline, and by Watt and Drummond whostudied the di-iodo(N~mefhyl-ortho-methoxyaniline)palXadium(IX)

-1c oiiip lex < The former worker looked at the 1200-1300 cmregion of the infra-red spectrum, which contains the aryl-oxygen stretching vibration, whilst the latter considered thepalladium-oxygen stretching region.

Examination of the part of the infra-red spectrumassociated with copper(II)-, and nickel(II)- oxygen vibrationswas not attempted, since the complexes already contained acopper(II)-, and nickel(II)- phenolic oxygen bond.

The aryl-oxygen stretching mode for niethoxy-, and*“X ( X2 X3)ethoxy-aromatic compounds occurs at about 1250 cm” • ’

However, in the case of substituted 2-methoxy, and 2-ethoxy-(13)phenols this region is complex, principally because the

aryl-phenolic oxygen stretching mode also occurs here.The aliphatic-oxygen stretching mode appears in the

range 1050-1000 cm and for all the complexes and ligandsdescribed in this work, unambiguous band assignment was possible.

The participation of* an alkoxy group in coordination:

(ll)/2

with the resultant drainage of* electrons from the ethereal oxygen, would be expected to lead to a weakening of the aryl-, and aliphatic-oxygen bonds. Hence a lowering of the stretching frequencies for these bonds might be taken as a criterion of coordination. However, there are difficulties in selecting a suitable reference for measuring the shift to lower frequencies. The free-ligand is unsuitable since it has to be ionized for complexation, and this in itself might produce the desired shift in frequency. Consequently, wherever possible,the spectrum of a complex was compared with that of the alkali metal salt of the ligand, and with the free ligand.

2.2. Magnetism

Copper(II}9The magnetic behaviour of copper(II), which has a 3d

electron configuration with one unpaired 3d electron, has been extensively studied.

Magnetically, the complexes of copper(XX) can be divided into two main classes. Firstly there are those which possess an essentially temperature independent magnetic moment of about 1.9 B.M. , of which about 0.2 B.M. arises from spin-orbit coupling, in which the coppei'(II) centres can be considered to act independently of one another. However, at very low temperatures weak magnetic exchange may be observed. The second class comprises complexes for which the magnetic moment is close to* or below, the spin-only value of 1.73 B.M. at room temperature, and which have markedly temperature dependent moments. When this is the case the complex is said to possess a subnormal

(14)magnetic moment, A subnormal magnetic moment is indicative of strong antiferromagnetic exchange between copper(II) centres. When encountered, information relating to the number (n) of interacting copper(II) centres, and to the magnitude of the antiferromagnetic exchange, can be obtained from a study of the dependence of the magnetic susceptibility on temperature. This is achieved by comparing the observed magnetic results to those predicted by theory. The theoretical treatment for copper(II) dealing with antiferromagnetic exchange between two or three centres within the same molecule (intramolecular antiferro­magnetism) is well established, but in the case of exchange

between magnetic centres being transmitted essentially overthe whole of the crystal (interraolecular antiferromagnetism)

(15)the theoretical work is less clear.

most cases only the spin contribution to each dipole involved is considered, and the coupling is taken to be between the spin angular momentum vectors on the centres. The interaction energy is

where J is the exchange integral, with positive values indicating that the spins tend to align parallel to the same direction (ferromagnetic exchange), and negative ones that they tend to align antiparallel (antiferromagnetic exchange).

Binuclear compoundsIn general, providing the interacting ions are identical

and their ground terms are effectively S terras, the magnetic,(16)susceptibility of an isolated dimeric aggregate is given by;

where N = Avogadro1s number.Net = temperature independent paramagnetic susceptibility, g = Lande splitting factor.P = Bohr magneton, k = Boltzmann1s constant.T = temperature.

S* = quantum number specifying the total spin of the systemx See page 48.

Magnetic exchange between two centres i and j arisesbecause of interaction between the magnetic dipoles. In

-2J (Si.Sj)

)+ Na

which can take values 0 to 2S in integral steps, S being the quantum number specifying the spin angular momentum of each ion.

ESf = the energy of a level specified by S f in the absence of a magnetic field.For a dimeric copper(II) aggregate, S = -J-, and

*Xa = ssfd x ----1-----+ N0U3kT 1 + 1/3 exP<||)

(17)This is the well known Bleaney and Bowers equation whichhas been extensively applied to susceptibility data for numerous!binuclear complexes in which antiferromagnetic exchange occurs,in for example, the a c e t a t e s , c a r b o x y l a t e s , 1:1pyridine-N-oxide complexes, bidentate and tridentate Schifffsbase complexes.

The theory has been extended by Earnshaw, Figgis and ( 22 )Lewis to isolated n-meric (n=3 to 10) aggregates of

equivalent electron spins with nearest neighbour exchange being considered. The resultant susceptibility expression is similar to that given in the dimeric instance, but incorporates the term X^l.(Sf), which specifies the number of times a given S* value may occur. S* is the quantum number specifying the total spin of the system.

Y_ S>(S'+l)(2S>+l)f^(S')exp(-|^-)Xm _ Ng3B2 SJ________________________ ^

3kT L (2S’+l)0.(s')exP(-i|->S' KThis expression is valid so long as the spread of the energy levels of any state which is appreciably occupied is small compared with kT. This condition arises because in the development of this equation use is made of the Van Vleck xSee page 48.

( 23 )susceptibility formula in which it is assumed that,

^ > > E *1 / / 1,kT > > E ^ H + E*2^H2 +m i. m(x)where E. is the xth order Zeeman coefficient for the energy i, m

of a level in a magnetic field* H; i and m specify quantum numbers. Upon application of a magnetic field the energy spread for a state specified by S* is (2S* + l)g(3H. Importantly, as n becomes large a value of S* which is sufficient to give (2S* + l)gJ3H ~ kT may occur and invalidate the expression. However, if J is large the expression may still be valid since the energy of states with large S* should be thermally inaccessible: the energy of a level specified by S1 is,

ES' = + 1) (22).nIn the present work there is interest in the theoretical

interpretation of the magnetic susceptibility data for polymericchain systems in which n tends to infinity. The theoreticalproblem has been tackled in two ways; firstly the spins have

(ph.)been treated as vectors (the Heisenberg apj>roach ), andsecondly they have been considered as scalar quantities (the

(2 5)Ising approach ). Of the two methods the first is morerealistic, but does not yield explicit functions for thesusceptibility, although Fisher and Bonner have performed

( 26)a numerical analysis on this basis. ' On the otherhand, the Ising approach does yield explicit functions.

The one-dimensional Ising modelThe crystal is considered to be made up of a number of

infinite chains of magnetic centres. Nearest neighbour

exchange occurs between magnetic centres in the same chain, but no interaction occurs between the chains. An axis of anisotropy, s, is defined, along which the spins tend to align parallel or antiparallel. The susceptibility along such an axis is and perpendicular to it, X j . In this work theresult for the a direction is taken to hold for the perpendicular direction.

Each magnetic centre in the chain is assigned a scalar spin coordinate, 0, which can take the value _+!. If 0 = +1, this corresponds to the spin state with the spin in the direction of s, then 0 = -1 corresponds to the spin in the opposite direction.

zAto = +1

1° = -1

The interaction energy between two centres, i and j with respective spin coordinates 0^ and CK is defined as,

Eij = - J 0. 0 if i and j are nearest neighbours,^ Jotherwise,

Eij = 0.

J is a measure of the strength of the coupling. Eij = -J if the nearest neighbours have the same spin, and Eij = +J if they are unlike. J is negative for an antiferromagnetic.

In addition, a magnetic moment, m, is assigned to each centre and the energy of interaction of the ith particle with an external magnetic field chosen to be,

E_j. = -mHO. .H x

(27 28)The problem is now one of statistical mechanics, * and the following expression results,

kiHM = Nm sinhrkTz-4Jx . , 2oH

eX?(kT“ ) + 3111,1 kTJL ~ 2

where M = magnetic moment per mole at T°K in an external magnetic field H. For copper(II), S = \ and gP/2.

If mK/kT is <^0.1, which is the case for T >5°& and K = 6,000 oexsteds, then

Xa •KT- P// " 4kl

, /gpHv 2exp( kT) + (2&t*

A field H of about 6,000 oersteds was employed in the experimental determination of susceptibility, consequently exp (~~~) ,and can be neglected. Allowing for temperature independent paramagnetism,

Xa = + Na . (A).

The subscript j j is dropped since we have assumed that the susceptibility is the same in the perpendicular direction as in the parallel direction.

An explicit function for X | has been derived by Fishex;^2^ Adams et al. have utilized it with x// in theform

X(T) = 4x /7(T) + fx , (T),powdor 3 '' 3 -1-

for interpreting the powder magnetic data obtained for anhydrous copper(II) c h l o r i d e . ^ A computer programme written in Elliot 503 Algol was used to check this claim. Unfortunately no

,2

reasonable agreement was found, neither was it found to satisfy the magnetic data obtained for the complexes described in this work.

Expression (A) has been applied with sonie success byInoue, Emori and Xubo in interpreting the magnetic data fordichloro(l, 2 ,4~triazole) copper (IX) and copper(II) benzoate

(31)trihydrate. The crystal structures for both compoundshave been determined, and the structural units in each case areinfinite chains. In addition to the linear antiferromagneticcontz'ibution to the susceptibility, an anomalous paramagneticeffect is observed at low temperature. Kobayashi et al. havealso utilised the expression in considering copper(II)

( 32)quinone complex salts, here an anomalous paramagneticeffect at low temperature was ascribed to the end effect andthe odd or even number effect of spins in the chain.

(2ft)Barraclough and Ng have proposed a theory based onexpression (A) and applicable above the temperature of maximum susceptibility, Tmax, but below which a modification was required.The analysis was later criticized^0 since Tmax had been taken

(33)to be For example, tetraamminecopper(II) sulphate(34)mnnohydrate, a well established case of a complex displaying

linear antiferromagnetism, displays a broad susceptibilitymaximum at about 3*5°K, but the Neel temperature is 0.37°K.This theoretical analysis has been applied to copper(XX) oxalate

(3 5)by Dubicki et al.

Curve FittingIn order to obtain an estimate of the interaction energy J,

-6the parameter Na in (A) was set at 60 x 10” cgsu. Using twosets of Xa and T values the equation was solved simultaneously forg and J 0 Xa for the remaining temperatures was then evaluated.Calculations were continued until the closest fit was obtained,,

(31)Alternatively, the graphical method of Inoue et al. couldhave been used.

The problem of obtaining the optimum value of J fromcurve fitting procedures has been discussed by Hyde et al.,

(37)and Ginsberg et al. The last group of workers put forwarda method for evaluating the temperature independent para­magnetism of the complex (t.i.p.). This method is not general and requires the compound to contain a small percentage of a paramagnetic copper(II) impurity, so that at low temperatures the entire susceptibility can be considered to be due to the impurity and t.i.p. of the complex.

Mechanisms of magnetic exchange(14)Two mechanisms of magnetic exchange are recognized.

Firstly, direct exchange between copper(II) centres, which isbelieved to be exemplified by copper(II) carboxylate complexespossessing a copper(II) acetate monohydrate structure, andsecondly indirect exchange (superexchange) which occurs by wayof some diamagnetic intermediary ligand atom Z, as is thoughtto be illustrated by copper(II) 1:1 pyridine-N-oxide complexes,

( 37)and iridentate Schif^s base complexes.The factors which affect the magnitude of exchange

(14 38)have been the subject of a great deal of discussion, ’and generally speaking the way in which the magnitude of exchange

is influenced can be summarized as follows. For both mechanisms the degree of interaction is dependent upon the magnitude of the relevant orbital overlap, and the greater this is, the greater the interaction. For the direct exchange mechanism the nature of the attached ligand is important. An electron donating ligand will Increase the interaction, and vice versa.

In the case of superexchange occurring by way of Z, themagnitude of exchange is increased if the electron density onZ is increased, and vice versa. Furthermore an increase inelectron density at the copper(II) centres will decrease the

(37 39)exchange through Z, and vice versa,

The one-dimensional Ising model does not specify whether the nearest neighbour interactions are direct or indirect, and as far as the statistics of the problem are concerned it makes no difference to a first approximation,

Nickel(II)The magnetic moment for nickel(II) is more sensitive

to stereochemistry than in the case of copper(II).Moments for octahedral complexes ( ground term)

occur between 2.9 and 3*4 B.M,, depending on the spin-orbit coupling contribution. The moment should be independent of temperature and of small departures from octahedral symmetry.

Tetrahedral nickel(II) with an orbitally degenerate 3 ^

ground term should possess moments lying between 3*6 and 4.0 B.M. ,

and decrease markedly with ■temperature. However, departurefrom tetrahedral symmetry or electron delocalization results inthe moment becoming closer to the spin-only value and to varyless with temperature.

Square planar nickel(IX) should be diamagnetic, however,if* the square planar crystal field is very weak a paramagneticcomplex may be produced. This appears to be the case for theanhydrous nickel(II) ortho-hydroxyarylcarbonyl compounds in

(41)hydrocarbon solvents.

2.3. Visible Spectra

Copper(II)In order to relate electronic properties to structure

a background of information based on compounds of known crystal structure is necessary. Considerable effort has been made in this area by Hathaway and co-workers, who have concentrated their attention principally upon copper(II)-ammine and ethylenediamine complexes (42, 43, 44, 45, 46). Even with this background information the situation is still not clear, and this is particularly well exemplified by the work of

(47)Tomlinson et al. on the tetrammine copper(II) complexes.Of interest in this work are the stereochemical deductions

which can be drawn from the study of the electronic spectra ofcopper(IX) complexes when there is a local oxygen environment.Hathaway et al. have recently turned their attention to this

(48)area of copper(II) chemistry, cThe 3d electron configuration gives rise to only one

2free-ion term, D, which is ten-fold degenerate in spin and

One positron energy level diagram for Copper(II).

zxyz

dz

xy

Octahedral Increasing tetragonal distortion

FIGURE j.

Electronic Spectra

- Powder

Resu

lts

03i*I03HIn

CQN

i s03

103

J*•O

*0N

*4A103>I

03}T3

£■PtoP!q

H

"0PI ^ O 03 & ^H(3•H

CN O VO C \ in cn 03 VD« « • • • • • •

03 03 H o o o H OH H H H H H H H

CO O vo Os in cn 03• • • « • « • •

co cn H O o O ON coH H H H H H

co O cn vo O r - 03 cn• • • • • » • •

in in in cn 03 cn cnH H H H H H i- i H

VO VD H ON O H VO COin in in cn cn H

• • • • • • • •i>~ 03 03 03 03 03 03 03

•

CO

flo•H•POS•H•OPioooIn0PIC3HPiQPi<3CO 03w

CO

*3Pio*p01poH<0•Hridac8O-PaoPi10©PiPiOOO+>TJO■Po0PiPioo5a0H

*0ao,pHOH(3•H

p« 03 oo 0 • 03© - C3 <H 03 W

« 03 • W 0 O H 03H 0 O o Si • m O 03*0 0 +> 03 03 r» o < <*>a a •H K « • W VO a Ha H Pi cn /—s o O 03 w o

0 O © 03 © O 0 03 u •cH o toa +> 03 •P 03 O •H 03 < */a a 03 •H W •H O to O a a0 o © O 0 in 03 W *0 a £» o oo •H N o H " ** a O 0 o w 03 03

-p i a 0 a O a H (3 a ,P aPi a Pi O o © K o •H CQ o Ph 53in +> ’W ' c3 CO o 03 Pito 0 a •H a h a a aC3 o Q o o CO 0 • • • •o H 03 cn ci-*

orbit. All or part of the five-fold orbital degeneracy is removed upon the application of a crystal field, the precise extent being dependent upon the field symmetry.

CuO^ chromophoreRegular octahedral stereochemistry is not found for

copper(II) and this may be attributed to the Jahn-Teller effect, however, crystal lattice packing forces and repulsions between neighbouring ligands may be responsible. The one positron energy level diagram for copper(II) in a normal tetragonal environment is given in Figure 1. The order of levels is dependent upon the degree of distortion. Table 1 summarizes the electronic data known for the CuOg system, which was compiled from reference. Copper(II) in diaboleite has symmetrywhilst in kroehnkite, copper sulphate, benzoate and copper formate this is approximately the case. Dibarium copper(II)

(49)formate tetrahydrate has been discussed in terms ofsymmetry.

The results of Table 1 show a clear correlation between the degree of distortion and the energy of the d^2_y2 — d 2 transition. Furthermore, for a copper(II) complex of unknown stereochemistry but having a local oxygen environment, the observation of a main peak (12.7 - 15.8 lcK) and a low energy shoulder (10.3 - 12.9 kK.) in its electronic spectrum may be taken as tentative evidence for the tetragonal CuOg chromophore.

CuO^ chromophoreRegular tetrahedral complexes of copper(II) are not found,

03mǤ

<3U+>o0aw : o•Hsr;o•poi ©HP

e+•+o V0 o oK-l « • • •Mr-3 CTs r- H Hx 03 03

a A rPp ta <Qf-i ca 03 V0 VD+> • • # •0 CO r- <A CAo H Hawo•H A .3P ta CQ0 IA «sQ 1A IAf*i 0 • ? •-P CP CO ODO0Hft}

0Op <—* /~N0 03 «A H HU IA IA IA IA0 s«* w w<H0ft

T3I0Pia0O

HOPO03tao

«at<U

O03mo

HHw0PiPi00H{»•PPrQ1+>!

WHHw£h0PiPi0OH>>a0 k Pi 0 n •H1&

U0<wCQfl(3+>0tou(3.SO

Hrs HH V-'H 0U PhO Pift 0ft O0o 0

/-v -P0 (3+> fi(3 •HS3 B•rt (3n*»* 0(3 fl© OP •d0 •HT) H•H >»H 0!» •Ho H*H (3H Wd H0 t>»H Pi>> 0•P UPi.a 0i W•P •HI

*-r1

f-* wta W•H

h0

e0•H-P(3

1 § fS(3

2Splitting of the D spectroscopxcterm of Copper(ll) in a crystal field of D?d symmetry.(ref.5l)

FldlTRE 2.

this may be ascribed to the operation of the Jahn-Teller effect.On the other hand, spin-orbit coupling may produce sufficientsplitting of the T0 ground term to overshadow the Jahn-Teller

( 23 )effect, and this appears to be the case in Cs^CuCl^.Distortional influences can be minimised if copper(II) issubstituted into a tetrahedral lattice. The spectrum ofcopper(II) substituted into ZnO has been recorded, an(jroom temperature a single broad band at 5.8 kK is observed.Table 2 summarizes the electronic data for copper(II) complexeswith approximate tetrahedral stereochemistry. Figgis et al.have discussed the assignment of the spectral data for

(51)the Schifffs base complexes in terms of symmetry(Figure 2), similarly, Hatfield and Piper have done so forCSgCuCl^ (52*) ancj Karipides and Piper for CSgCuBr^.^"^ Onthe basis of this data the appearance of d-d transitions in therange 5*0- - 13*8 kK may be taken as evidence for a distortedtetrahedral CuO^ environment.

Bis(salicylaldehydato)copper{II) and bis(acetylacetonato)copper(II) are examples of the planar CuO^ chromophore. Usingthe solid state reflectance spectra of these compounds as areference,Graddon and Mockler have proposed a planarstereochemistry for the anhydrous ortho-hydroxyarylcarbony1

(54)copper(II) complexes. In the solid state these spectraare characterised by high intensity bands in the ultra-violet region, the tail of which overlaps a broad absorption envelope^ in the visible region, extending from 20.0 - 14.0 kK. This partial obscuring of the vis.ible region does not occur in the case of the synthetic silicate BaCuSLjp^g, Egyption Blue, which

T

20,0

10.0

0

10.0

20.0

io.o0 15.0 20 .0

2p**v O\

'g

Ig'ig2 -

Ig.A.X

10

A,

Energy as a function of crystal fieldgparameter for. the 3d system nickel(ll).

IEnergies are in kK and the S state of the free ion is not shown-.

FIGURE 3.

is an example of a square planar CuO^ system. Clark and Burns(55)have studied the electronic spectrum of this material in detail,

and the absorption maxima data is given in Table 1.In conclusion, this section has attempted to show that

a set of criteria can be established for distinguishing betweenthe tetragonal CuOg, distorted tetrahedral C u O a n d planarCuO^ chromophores.

Nickel(II)In this section emphasis is placed primarily on the use

of electronic spectra as a guide to stereochemistry rather than upon detailed interpretation.

Six coordinate complexesThe energy level diagram for nickel(II) in an octahedral

crystal field as a function of the crystal field parameterAo(Ao = E(e ) - E(t? )) is given in Figure 3* Six coordinate S "Scomplexes usually exhibit a simple spectrum involving spin-

3 3 3allowed transitions to the Tn ... T. (F) and T_ (P) states,2g lg Igoccurring with low intensities in the ranges 7.0 - 13.0 kK,11.0 - 20.0 kK and 19.0 - 27.0 kK respectively. In addition, absorption bands arising from spin-forbidden transitions,AS ^ 1, may arise. The origin of spin-forbidden transitions

/ / pq Nhas been discussed by Dunn, and Ballhausen and(57)Griffith have considered this subject in depth. A

2S+1 2S *+1transition between a state X and J Y is made allowed2£>+l 2S * +1by the admixture into X of a state Z with the same

Diffuse

Refl

ecta

nce

Data

fab9HEhca

Cca

to03

toH

<ca

to03

to03£h

H

bO03

ca

tomH

cca

fa

bO03

Ehca

toH

ca

toca

Ehca

to03

bO03<

ca

•O§o&oo

NO ca A - A - co CO IA« « • e • « • 0o CM IA IA ca ca •cH LA

03 03 03 03 ca 03 03 03

ON•

Hca

03 IA co LAt a • • • • • •o ca ca H ca 03 ca03 03 ca 03 ca 03 ca 03

LA 03 03 in• « • • v • • • • •

A- ON H CO rri ON ON O HH H 03 H 03 H H ca 03

03 CO •cM LA CO A- LA« • • « % • • • •o ca 03 LA LA 03 ca ca

H H H H H H H H H

CO 03 m LA vo CN ca A . H♦ • • • • • • • •

VD A- A - CO CO VD A - A* CO

03uCQ•riS3

03HO•ri

fa•ri

IA

©OPS0Ci0<H0bSoPi<HTS0H

•d0ci•H0PPiOo

rdPto0od0Ci0<w0Ci03cl0oLA

rNP•H©cs0pd•H•ri

Cl0idr*400*0Ci0Ja©3PP0Apflo

•riCi0P•HCiO0tdPbOPS•HW2rdoPSto•ri

CO LA LA LA vo A- •ri w• • • « • % • • « • 0 «

H H 03 ca H 03 03 ca u 0H H H H H H H H H ON 0

©PSo•riP•H©PS0CiP

PS0'd*0

©PSo•Hp•H©PS0CiP•d0oHH003 H •ri Srf* •H+ O fa & & PSVO ca . ca s 0 U •ri

/—s ca 0 w 0 Pio Pi H W o H ©03 W-H O o o 0su o § K 0 PS PSv_x O O 03 ca p 0 •ri 0•ri •ri •ri •ri •H •ri 0 0 fa &S3 a $5 & Q to to p

2S*+1spzn as Y under the influence of spin-orbit couplingenergy, or by the admixture of 2o+1W into 2u?+1Y. The transition is now allowed, but the intensity is proportional to the square of the coefficient of admixture. This coefficient is dependent on the proximity of these states,and if close large intensities can result*

1 1 Figure 3 shows that the E state arising from the Dgstate is close to the T_ (F) state for Ao = 7,0 - 9,0 kK,IgFor six-coordinate complexes possessing crystal field

3parameters of thi3 size the transition to T^ (F) appears asa well defined doublet. This has been considered to be due to

3 1the spin-forbidden transition A_ - E arising from the spin-O SX 3orbit coupling interaction betxveen the S and T„ (F) states*“ g i-gJorgensen supports this interpretation and has investigated the

/ j - q \ ivariation in position of this band as function of Ao.Kfcrwever, Liehr and Ballhausen hold the view that the double­peaked absorption band arises from the spin-orbit fine structure

the (23)

3of the T, (F) state, the spread of which is too large forvibrational bridging.

Table 3 summarizes the diffuse reflectance electronicdata for weak field nickel(II) compounds having octahedral anddistorted octahedral coordination polyhedra. The table

(59)was compiled from the work of Ludi and Feitknecht, in whichthey satisfactorily interpreted the data for nickel bromide, chloride, fluoride, and hydroxide using the expression derived

o / ey \by Griffith for the energy of the two T 1 states, namely.E(3T ) = i(I5B - Ao) + %(225B2 - l8BAo + Ao2)^,xg3where £( A0 ) = -2Ao,

and B is the Racah parameter for the nickel(II) ion in the complex. From Figure 3 Ao is given by the energy separation between and . One interesting point to emerge from the2g 2gwork of Ludi and Feitknecht was that for compounds in whichthere was moderate or strong distortion no additional statesplitting was observed. This was in contrast to the resultsfor the corresponding cobalt(II) compounds. For a ligand

3 3field the T^ and states of an octahedral complex3 3 3 3split into Bn + E and A„ + E states. This splitting1 2g g 2g g ^

has been observed by Goodgame et al. for the complexes ofnickel(IX) with heterocyclic amines. The spectra are generally characterised by four absorption bands in the range7.0 - 17.0 kK.

Four coordinate complexes( /T *i \Cotton and Wilkinson have discussed the spectral

features of tetrahedral nickel(II) complexes. In general they display a multiple band near 16.0 kK assignable to the 3 3T^(F) - T^(P) transition with weak bands on the low and highenergy sides assigned as spin forbidden transitions to components of the and "G states respectively. A nearinfra-red band at about 8 kK is assigned to the ^T_ (F) - ^A01 £

3 3transition, but the T^(F) to T^(F) transition occurring in the range 3 - k kK can rarely be unambiguously assigned as this region can contain vibrational bands due to C-K fundamentals as well as 0-H bands due to any moisture. In contrast to octahedral complexes,tetrahedral compounds display

/ c n\relatively high intensity bands ' and it has been shewn2that the difference is of the order of 10 . Furthermore,

3 3the intensity of the T^(F) - T^(P) transition has been shewn3 3 (62^to be 10-20 times more intense than the T,(F) - transition.i 2

Planar complexes of nickel(II) can readily bedistinguished from octahedral and tetrahedral compounds in asmuch as no electronic transition occurs below 10.0 kK. Thelowest energy band is commonly attributed to theb0 (d ) - b, (d ?- p) transition, however, Lever has2g xy Ig x^ y^ * ’questioned the validity of this assignment and also generally considers the factors which influence band assignment.

The lowest energy band in planar nickel(II) complexes increases in energy approximately in the sequence,

NiS/i < N i N 202 ~ NxN2X2 ~ NiC>4 <NiP,,X2 < N i N 4

In the case of the square complex bis(2 ,2,6 ,6-tetramethyl-31( 6 )heptanediono)nickel(II) and the diamagnetic square form pf

bis(3-phenyl-2,4-pentanediono)nickel(II) the lowest energyband occurs at 18.69 kK.

Experimental

CO

©HA©H

©Pi 033 COP 1© HPi CO©

P «*•H ON►Q IN

I• r^-

P r -Pi

•s

/-No H

0 CO1

• IAp •Pi O

• COS

'd33o&ao

o

H0Pi©.3P.Hb£4fH0

1

1

©P(8

HO3©r3Pt0up•r43!-4«1Jo0AP©a103

oAp©03

•H©©©POPi

HO3©f3

0AP©a103

CO[N. CO I H VO H |n

% • 3© ffi 53 ©0 ♦ <rs • a3 Q LA Q p3 ft VO • ©0 PQ w PQ ©

w «

& I J?

P»3to

♦HAIAoo

HO3©

CDOH

•cHICO

OH

VOHHI

iaHH

M3o

*dJo

voON

W

PQ

ft3 ft353 © ©» a aQ VO p pft VO © ©PQ V © ©PQ PQ

H03APiHJo © ©

©3©3© a •3 'd HA 0 Jo Jo 0Pi u »3 .3 PH © © 0Jo I 'd *d Ua 41 H H P0 J © © •H H3 bri N N 3 O3 3 i 3I O © © in ©-4* A •p A i AI5o p©a. &0 S?0 k00AP

*ri•0I ■s ■sJou*db

0,3P©1 VO fES1 A• «3I ©102 02

1a Pi 03 03

HO3©>3r'l0 A p©a•H•01

VO

03

tom CO 03 VO• ON O VO IA ONVO O H H • HIN1 H1 ILA HI HON 1 H I *LA IN • m 1 ©• O 02 H *dLA H O H ON 0IN H 8

©•H-PPO

©PH

©P©0 •ri H •H W

1•H©©©POP<

<H• • Pi 0u *d W 1© © ©p 3 73 m o3 •ri 3 CO 3P © © ©o P <w ©© »Q © 0 ©<H 0 u3 So © Pi3 © PQ •H© © © ©a 3 •» Jo A

3 — H P© to © 0A •ri •d UP 3 -d o

3 rNJo H 0 r3 ©rP © P(

O a © p'd •H 0 © o© P o © 3•ri Jo AH H o •dPi © •H 73 •Hft 3 3 © »d3 © © tow 0 to 3 s3 0 p© o o H f-i© •ri O pa <w U oJo 0 Pi ©rri Jo PiP 3 Jo Jo ©0 0 u .P© p © TJU o 3 *d ©•ri © 0 © Pi*0 •ri 3 I

© P •H ©•d •H 0 © u© P •H P Vhw © Q .a 3& W W» o •H

• « ft ftH 03 CO

3*1. Ligands and Reagents

The purity of the ligands was checked before use and some were recrystallized (see Table 4.). All other reagents were of A.R. quality. Standard I.N* potassium and sodium hydroxide solutions were obtained from B.D.H.

The use of cumbersome systematic names will be avoided. The following abbreviations based on the trivial names of the ligands (parentheses) will be used.

Systematic Name

2-methoxy-4~formylphenolato 2-methoxy-4~nitrophenolato 2-methoxyphenolato 2-ethoxy-4-formylphenolato 2 ,6-dimethoxy-4-formylphenolato

3*2. Metal Analysis

Copper(II) was determined volumetrically using EDTA and Fast Sulphon Black F as the indicator. This indicator was found to be specific in its colour action towards copper(II) in ammoniacal solution and its end-point sharp and stable.The organic matter in 0.1 gm. samples was removed by heating with a mixture of concentrated nitric acid (8 ml.), perchloric acid (4 ml.), and concentrated sulphuric acid (2 ml.). It was absolutely essential that evaporation of the fuming mixture was completed to dryness. The optimum working pH for the indicator was 11, and if acid was still present after the

Abbreviat ion

V. (vanillinato)N. (4-nitroguaiacoiato) G. (guaiacolato)B* (bourbonalato)S. (syringaldeliydato)

FIGURE 4.

SCHEMATIC DIAGRAM O F'STA N TO N *' AUTOMATIC DIRECT READING THERMO BALANCE. •

'JWHBLE HGH LOW SWITCH IMG

H>~0«ELECTRONIC

Schematic Diagram of Thermobalance.

Silica Sheath

Asbestos .Collar

Gas Inlet/Outlet

W Pt-Crucible

■Thermocouple

•Rise Rod

O-ring Seal

Gas Inlet/Outlet

Furnace arrangement for controlledatmosphere work.

•FIGURE 5.

organic matter was removed sufficient ammonium salts remained in solution to set up an adverse buffer action.

Nickel(II) was determined gravimetrically using d i m e t h y l g l y o x i m e f T h e organic matter in 0.1-0*15 samples was removed by heating with concentrated nitric acid (8 ml.), perchloric acid (6 ml.), and concentrated sulphuric acid (2 ml.).

3.3. Microanalysis

The majority of microanalyses were performed by Dr] A. Bernhardt (5251 Elbach uber EngelskirChen, West Germany), and the remainder (marked asterisk in Table 5) with a Perkin-r Elmer 240 Elemental Analyser. Under optimum running conditions the absolute accuracy of this machine with respect to carbon, hydrogen and nitrogen determinations was +, 0.3%. All anhydrous samples were analysed by Dr. A. Bernhardt, and quantitatively dried prior to analysis.

3.4. Stanton Thermobalance

Lukaszewski and RedfernJ^^ and Coats and Redfern^*^ have reviewed the subject of-thermal analysis. Figure 4 is a detailed schematic diagram of the thermobalance and no further comment as to its mode of operation will be given here. The machine (type HT-D) employed in this work was sensitive to +_ 0.1 mg., and the maximum temperature attainable was l400°C.The furnace set-up in Figure 4 was not employed and instead the arrangement shewn in Figure 5 adopted. This enabled work

to be carried out in a controlled atmosphere by either pumping gas upward through the furnace sheath (inlet volume rate,400 ml./min.; outlet, 200 ml./min.), or by passing gas downward (inlet volume rate, 4-00 ml./min.). Nitrogen was the only artificial atmosphere employed. The gas was dried by bubbling it through concentrated sulphuric acid and then with an anhydrous magnesium perchlorate column. The temperatures at the furnace wall and at the base of the crucible were recorded by means of platinum, 13% rhodium/i)latinum thermocouples, and at the latter point this was done to an accuracy of 3°C. Platinum crucibles were always used except when macroscale work was undertaken, and in these circumstances aluminium foil crucibles were employed.

When dealing with a new hydrated complex compound the theroobalance was found to be an extremely useful tool: firstly,it was possible to determine that the drying conditions adopted in the preparation of the compound, yielded a material which was stable when exposed to the atmosphere at room temperature; secondly, the degree of hydration was accurately determined; thirdly, the thermal stability of the anhydrous compound was appraised; and fourthly, the machine was used for the macroscale preparation of the anhydrous complex.

In this study the following routine preliminary operations were carried out when dealing with a new hydrated complex. The complex was dried in vacuo over anhydrous calcium chloride. A small sample was taken (*-100 mg.) and its behaviour, over a period of 1-2 hours, observed on the thermobalance at room temperature under normal atmospheric

conditions. This was then repeated under an atmosphere of dry nitrogen. If the mass-gain/loss base line was constant under both conditions then this method of drying was satisfactory. However, if there was a mass increase (moisture uptake) with time for the air run, and a mass decrease (moisture release) with time under dry nitrogen, then the drying procedure initially selected was unsuitable* In such cases, and in those in which the complex was unstable under dry nitrogen conditions alone, air drying was employed.

The bouyancy correction for the crucible used, under the conditions employed, was determined. This correction was especially large when dry nitrogen was passed downward through the furnace. As a general rule small samples (100-150 mg.), finely divided and loosely packed, were used and a heating rate of about 2°C/min. applied. After each run the infra-red spectrum of the material was recorded to checlf that the mass loss was due to the removal of water.

Macroscale preparations of anhydrous complexes were always performed with dry nitrogen passing downward through the furnace, which was operated isothermally on reduced power (only possible for temperatures below 200°C). This method of preparation was satisfactory only if the anhydrous complex was thermally stable, and if the rate of loss of water was greater than 0.1 mg./rain., otherwise vacuum assisted conditions were considered. Once dehydration was completed the product was cooled to room temperature on the machine with dry nitrogen flowing through the system.

w

VOl

CM

fHCO

CMMM

( H OU tfl Hr-, W «M

VO MCM

PFI

GURE

*i-0.

3*5# Infra-red arid Visible Spectra

Infra-red mull spectra were obtained on a Unicam SP.200, and Grubb Parsons (model MK2E2), *Spectromaster1, spectro­photometers. Sodium chloride windows were employed for the former instrument, and potassium bromide windows- with the latter* With potassium bromide Windows the Grubb Parsons machine gave unsatisfactory results below 500 cm. ^ since it was operated at a high gain setting. This was obviated by using matted polythene plates, matted to reduce interference fringe bands, with the gain at the normal setting. Hygroscopic compounds were mulled in a dry-bag filled with dry nitrogen.

Visible diffuse reflectance spectra were recorded ona Unicam SP.700 spectrophotometer using lithium flupride asthe reference. The normal cell supplied with the instrumentwas satisfactorily used for samples which were moderatelyhygroscopic, provided it was packed in a dry-bag and theseams Sealed^ For very hygroscopic materials the cellshewn in Figure 6 (3), which was sealed under vacuum, was used.Spectra were recorded at liquid nitrogen temperature using

(10)the attachment previously described by Butcher.

3.6. Magnetic Measurements

Molar magnetic susceptibilities, X^, were determined by the Gouy method, and

v - fk »v + B *w ) m *M = ( W

where k = volume susceptibility of air. v = volume of the specimen •vr = force exerted on the specimen by the magnetic field.W = weight of the specimen.M = molecular weight of the compound calculated on the

basis that there is one paramagnetic metal ion per •molecule*.B = constant involving the dimensions of the specimen

and the magnetic field.B was determined by calibration with tris(ethylenediamine)

nickel(II) thiosulphate. In order to calculate Xa, the magneticsusceptibility per gramme atom of the paramagnetic metal ion,X^ was corrected for the diamagnetic contribution of all atoms,including that of the paramagnetic metal ion, by using the

(72)corrective values given by Lewis and Wilkins. The t.i.picorrection was not included. The effective magnetic moment of the paramagnetic metal ion at the temperature of measurement was calculated from the equation;

[Ieff = 2.84(Xa.T)^.B.M.

For the measurement of magnetic susceptibility overa range of temperature (~80-300°K) the Gouy balance systemdevised by Earns haw was used. Earnshaw has described and

(73)discussed the merits of this particular system in detail.(74)The apparatus was calibrated with nickel chloride solutions.

As measurements were made in an atmosphere of dry nitrogen the equation for the molar magnetic susceptibility reduces to,

where g = acceleration due to gravity at the place of measurement.1 = length of* specimen.H = magnetic field calibration value for the length of

specimen, 1.The susceptibility notation used in this work is now

reviewed. The following terms are defined:

= the observed molar magnetic susceptibility, calculated on the basis of one paramagnetic metal ion per •molecule 1.

Xa = the magnetic susceptibility per gramme atom of the paramagnetic metal ion.

Xm = n.Xa, for an n-raeric aggregate of identical paramagnetic metal ions.

Hygroscopic compounds were handled in several ways.For large quantities of material the Gouy tube was packed in a glove-bag filled with dry nitrogen. When quantities were small and the compound moderately hygroscopic the miniture ’pig’(Figure 6(2)) was used,which could be connected directly to the Gouy tube or via a short length of polythene tubing.The •pig• and its contents were heated in a vacuum drying pistol, to remove last traces of moisture, immediately before connection. For extremely moisture sensitive compounds the apparatus shewn in Figure 6(3) was employed. The compound was dehydrated in bulb B, and with the apparatus still under vacuum, packed into the Gouy tube, which was sealed off.

3.7* Mass Spectra

Mass spectra were obtained on an A.E.X. MS#12 machine,, The standard direct insertion lock was used to introduce solid samples directly into the ion source region*

3.8. Analytical Results

The analysis figures for the prepared complexes are catalogued in Table 5*

Table 3

Found % Required %Compound c. H H2° N Cu C H H 2° N ! Cu

CuV2,2H20 47.68 4.85 - - 15.87 47.82 4.52 “ - :15.81

CuV2 1- 52.34 3.80 - - - 52.53 3.86 - - -CuN ,4H„0 2 * CuN 3* 4-1.94 3.14

15.057.10

13.4842.06 3.03

15.277.01

13.47

Gu G2 54.15 4.70 - - 20.72 54.28 4.56 - - 20.52CuB2,2H20

CuB2

50.4054.75

5.334.73

— — 14.66 50.2954.89

5.164.61

— — 14.78

CuS2,3H2CK CuS2 6-

44.9650.63

4.884.20

11.61 --

45.0550.77

5.044.26

11.26-

*

C H h 2° N Ni C H H 2° N Wi

NiV2t2H20K 48.63 4.68 ~ - 14.71 48.41 4.57 - ' - ;14.79

m V 2 1 * 53.19 4.02 - - - 53.24 3.91 “ ... -NiN2,4H20x 36.11 4.12 - 6.08 12.46 36.OO 4.32 - 6.00 12.57NiN0,2H00 8* 38.88 3.89 - 6.66 - 39.01 3.74 - 6.50 -NiNg 9* 42*38 3.22 - 7.10 - 42.57 3.06 - 7.09 -NiG„,2H 0® 49.20 5.24 10.37 - 16.90 49.31 5^32 10.57 - 17.22NiGg 10' 54.96 4.47 - - 55.14 4.63 - - -IWS2,3H20X

11NiS045.3351.26

5.044.46 - -

12.17 45.5151.35

5.094.31 - -

12.36

x Analysed with a Perkin-Elmer 240 Elemental Analyser. The percentage deviation for duplicate analyses was always well within the specified accuracy of the machine (Section 3.3.)*

/continued....

Se x ing conditions and % drying loss prior to analysis

1. 2mm., 80°C; 1.40%.2. Difficult to microanalyse for C, H and N since the compound

readily partially dehydrates when subjected to slight mechanical pressure and mild desiccation.

3. Dx-ied 75°C.4. 2ram., room temperature 0 .00%.5. 2mm.f 90°C; 1.77%.6 . 2mm.f 85°C; 2.47%.7. 2mm., 100°C; 1.82%.8 . 760mm., 8l°C; 2.31%.9. 2mm., 136°C; 2.26%

10. High Vacuum, 80°C; 0.00%11. High vacuum, ll4°C; 7.94%

All lower hydrate and anhydrous compounds were prepared directly from the pure, fully hydrated compound, except for CuG-^

VO

+»

o

+»

o

M

CM

+»O

- P •'

CD +»

O *HO rHO rH£3 *H

EH *H M

Tim

3.9* Preparation of ComplexesDrying in vacuo was performed over anhydrous calcium

chloride. The yields of the complexes were essentially quantitative„

CuVr,2Ho0___

2--methoxy~4~formylphenol, vanillin (6.3D gm. , 0.04l4 mole), was added to water (750 ml.) containing sodium hydroxide (40.5 ml., l.N., 0.0405 mole) and sodium acetate trihydrate (0.4 gm.). The mixture was vigorously stirred and the solution filtered. Copper sulphate pentahydrate solution (4.6? gm., O.CT87 mole) in water (200 ml.) was added dropwise to the filtrate which was stirred. The yellow-green microcrystalline complex was filtered off and dried in vacuo*

CuV0

This was prepared by dehydration of CuVo,2Ho0 on a Stanton Thermobalance under an atmosphere of dry nitrogen.The thermogram for the isothermal dehydration process is given in Figure 7.

CuN_,4Ho0£.~J i-J

Potassium 2-methoxy-4-nitrophenolate (7.35 gm. ,0.0355 mole) and sodium acetate trihydrate (0.4 gm.) were dissolved in water (300 ml.). To the stirred solution was added dropwise a solution of copper sulphate pentahydrate

(3.93 gm., 0,0159 mole) in water (250 ml.). The bright yellow- complex which separated from solution was filtered off (Whatman 542 filter paper), washed with water (*.250 ml.) until the washings were free of sulphate ions, and air dried for two months.

When ground or subjected to mild desiccation the complex partially dehydrated giving a brown material.

CUN2Prepared by ’isothermal’ dehydration (58-86°C) of

CuN^, 4lig0 on a Stanton Thermobalance under an atmosphere of dry nitrogen. The thermogram showed no evidence for the existence of an intermediate hydrate.

Although the instrument was set for isothermal working on reduced power, its design did not allow satisfactory control below about $Q°C+

CuG22-methoxyphenol, guaiacol (2.00 gm., G.0l6l mole), was

added to a stirred aqueous solution of sodium hydroxide (14.51 ml., l.N., 0.01451 mole, diluted to <)Q ml.)* The ligand slowly dissolved and the solution became darker in colour with time due to oxidation. Copper sulphate pentahydrate solution (I.63 gm., 0.00653 mole) in water (90 ml.) was added dropwise to the stirred solution. The light-brown material which separated from solution was filtered off, press dried between filter paper, and dried in vacuo.

2~eihoxy~-4-*foririyXphenoI, bourbonal (3-00 gm,, 0.0181 mole), was added to water (750 ml.) containing sodium hydroxide (i?. 18 ml., Ilf., 0.01718 mole) and sodium acetate trihydrate (0.4 gm.). After vigorous stirring the solution was filtered and copper sulphate pentahydrate solution (2.04 gm., 0.00817 mole) in water (200 ml.) added dropwise to the stirred filtrate. An orange coloured microcrystalline solid separated from solution. This was filtered off, sucked dry for three hours, and finally dried in vacuo.

CuB2This was obtained directly from CuB0,2Ho0 by isothermalu tU

dehydration (94°C) on a Stanton Thermobalance under an atmospher of dry nitrogen. Last traces of moisture were removed by heating in vacuo at 90°C, 2 mm.

CuS 3Iio0<r.\ u-j

2,6-diniethoxy-4~formylphenol, syringaidehyde (1.00X2 gm 0.005492 mole), was added to water (240 ml.) containing sodium hydroxide (5«32 ml., I.N., 0.00532 mole) which was vigorously stirred. The solution was filtered and copper sulphate pentahydrate solution (0.6444 gm., O.OO258O mole) in water (60 ml.) added to the stirred filtrate. The buff coloured solid which separated from solution was filtered off, sucked dry for three hours, and dried in vacuo.

Isothermal dehydration (87-*92°C) of CuS^, 3HpO on a Stanton Thermobalance under an atmosphere of dry nitrogen produced 0uSo. The thermogram shoi^ed no evidence for there being an intermediate hydrate. Last traces of moisture were removed by heating in vacuo at 85°C,2 mm.

NiVrt,2Ho0

2-m@thozy-4-formylphenol, vanillin (3-00 gm. , 0.0197mole), was added to a stirred aqueous solution of sodium hydroxide (l8„7 ml., I.N., 0.0187 mole, diluted to 750 ml.).The solution was filtered and a solution of nickel nitrate hexahydrate (2.59 gm., 0.00891 mole) in water (200 ml.) added dropwise to the stirred filtrate. On standing overnight a bright green microcrystalline solid separated from solution.This was filtered off, sucked dry for three hours, and dried in vacuo.

!£s.The Stanton Thermobalance was not used for the

preparation of NiVV, from NiV^, 211 0 , since the rate of dehydration was too slow. Dehydration of NiVot2H 0 was accomplished by heating in vacuo at 100°C, 2 mm.

NiN0, 4H„,Ctii Cli

Potassium 2-raethoxy-4-nitrophenolate (3*00 gm., 0.0145 mole) was dissolved in water (720 ml.). Nickel nitrate

FIGURE

8

S3•HS3

tCMK

CO

CO

•H

A

OCOwM-CM

OCD

COMCM

O O roVO.mHMM

O OM

CM

CMCM

O o

om

CM, MW

0}eo•H+*•Hr—:COo

oo

orH<H>c'C?©u©.£3p<woG-p

CM

WO s

rH •CjD W G . ra v_^0}

Air

run

temp

erat

ures

given

in pa

rent

hese

s Sample

v/t. =

0.1141

gm.

hexaliydrate (2.05 gm. , 0 .0Q705 mole) in crater (96 ml.) was added dropwise with efficient stirring. Oh standing overnight a yellow solid separated from solution. This was filtered off, sucked dry for three hours, and air dried for three weeks.

NiN„,2Ho02 1 2This compoundwas best prepared by isothermal*

dehydration (44-8l°C) of NiN^^iH^O under an atmosphere of dry nitrogen on a Stanton Thermobalance (Figure 8), or alternatively by heating in air at 8l°C.

NiNp

Obtained by heating NiNo,4H00 or NiNo,2H_0 in vacuo at 136°Cj 2 mm.

N 1G2 ,2H 20

2-methoxyphenol, guaiacol (2.10 gm., G.CI69 mole), was dissolved in water (60 ml.) containing sodium hydroxide (1.53 ml., l.N., 0.0153 mole). To this was added, slowly, and with efficient mixing a solution of nickel nitrate hexahydrate (2.11 gm., 0.00725 mole) in water (60 ml.). Alight-grey solid separated from solution, which was filtered off, press dried bet’ween filter paper, and dried in vacuo.The dry product was treated with hexane, to remove free ligand impurity, and dried in vacuo.

N i G 2

Prepared by dehydrating NiG^t^H^O isothermally (87°C) on a Stanton Thermobalance under an atmosphere of dry nitrogen.

N±S2 ,3H20

2,6~dimethoxy~^-formylphenol, syringaldehyde (X.2500 gm., O.OO686O raoXe), was added to a weXX stirred aqueous solution of sodium hydroxide (6.66 ml., l.N., 0.00666 mole, diluted to 300 mi.). The solution was filtered and a solution of nickel nitrate hexahydrate (0.9380 gm. , 0.003225 mole) in water (80 ml.) added dropwise to the stirred filtrate. After standing for five hours a light-green solid separated from solution, which was filtered off, sucked dry for two hours, and dried in vacuo.

HiSg

NiS^ was prepared by dehydrating NiSg,3HpO in the apparatus shown in Figure 6(3) at 11^°C under high vacuum.The compound was extremely hygroscopic and immediately rehydrated on exposure to the atmosphere. Thermogravimetric analysis (t.g.a.) of NiS^^H^O failed to show the existence of intermediate hydrates.

Table 6 catalogues the complexes along with their colours.

Sodium 2-methoxy-4-foraylphenolate2-methoxy~4"*formylphenol, vanillin (0.5 gro.), was

dissolved in a minimum volume of sodium hydroxide (2.N.). A clear

yellow solution resulted, from which, on vigorous shaking, the sodium salt separated. The sodium salt was filtered off, washed with a small quantity of water, and dried in vacuo.

Sodium 2 -ethoxy-k -formylphenolat e2~ethoxy~4~formylphenol, boiirbonal (0.5 gm.) , was

dissolved in a minimum volume of warm ethanol. Sodium hydroxide (0.75 cal. , 4.N.) was added, whereupon the sodium salt separated from solution. It was filtered off, washed with ethanol and ether, and dried in vacuo.

Sodium 2,6-dimethoxy-4~formylphenolate2,6-dimethoxy~4~formylphenol, syringaldehyde (0*2025 gm.

was dissolved in a minimum quantity of warm ethanol. Sodium hydroxide (1*10 ml* * l.N.) was added* and on cooling the sodium salt crystallized from solution. It was filtered off, washed with absolute ethanol and ether, and dried in vacuo.

TABLE 6The Complexes and their Colours

CuV 2K„0 Yellow-green NiV2,2H20 GreenC u V z Brown NiV2 YellowCuN2,4H20 Yellow NiN2,AH?0 YellowCuN„

UsBrown NiN2,2H20 Yellow

CuGS Light-brown NiN2 BrownCuBg,2HgO Orange NiG„,2Ho0 2* 2 Light-greyCuB2 Brown NxG2 YellowCu S2,3H20 Buff NiS2,3H20 Light-greenCuS2 Brown NiS2 Orange

Copper(II) Complexes* Results and Discussion

FIGURE 9.Infra-red absorption in the range 9*00 - 10*00 microns (lIII - 1000 cm for the copper(ll)

complexes

Notes*(i) Transmittance increases in the direction

(2) The scale, I---- 1--- ! , is linear in micronsand covers the range 9*00 to 10,00, reading from left to right,

(3) Specific band positions are given in units of reciprocal wavelength (cm ).

( k ) #, see Appendix 2,

I

J 1

£ u Y.2jl

1021

*J£a3Li

JliL.

1028

£uH,

1020

1017

IiN.

IO jl

1038

. HG.

IO^Oj 3j022A

SJ\t . \ »

oil1037

103 6I04J|

IOIO

J

J.043j jE002

HH*

10^1

JL

L

«“V

*NaS«

1038

HS ,

FIGURE 10.Infra-red absorption in the range 7• 50 - 8*50 microns (1333 “ 1176 cm"^) for the copper(ll)

complexes.

Notes .(1 ) Transmittance increases in the direction •

■ Y(2) The scale, I L 1 , is linear in microns

and covers the range 7*50 to 8 .50, reading from left to right.

\3) The 1250 cm- reference line is sbown. See Table 9»(4) *, see Appendix 2.

r

* KN 1250

HN

1250

1250

1250

IHL*.

121

50

J

HS1250

inro

cu

c JrO

C\J

oo

in

MM

rOON

in

cu

cu

o

4.1. The Infra-red SpectraFigures 9 and 10 show the infra-red absorption for the

— 1 -1copper(II) complexes in the range 1111-1000 cm and 1333-12.76 cm .The latter region contains the aryl-ethereal oxygen stretching

-1mode (about 1250 cm , see Section 2.1.) and the former the aliphatic-oxygen stretching mode (1050-1000 cm’***', see Section 2.1.) for the 2-ethoxy-, and 2-methoxyphenols.

In Figure 9 there can be little doubt concerning the assignment of the indicated absorption band(s) to the aliphatic- oxygen stretching mode. Only in the case of the copper(II) complexes with 2-ethoxy-k-formylpheiiol (CuB ,2H„0 and CuB^)C* dU Ciis the situation less clear. From Figure 10 it is apparentthat any attempt to attribute one particular band to the aryl-ethereal oxygen stretching vibration would be unwise on accountof the complexity of the spectra. It is not possible to decidewith any certainty the role of the alkoxy grou|> as infra-reddata for compounds in which an alkoxy group is known tocoordinate is not available. None-the-less,a comparisonof the spectra of the free ligand, its alkali metal salt,

-1and the corresponding complexes in the 1000-1111 cm region,which is given here, may point to the non-involvement of thealkoxy group in coordination.

Figure 11 compares the region of the infra-redspectrum associated with the 0-H stretching mode (at aboutthree microns) for the hydrated copper(II) complexes. Ofparticular interest is the infra-red absorption of the complex,

— 1CuNo,^ H o0. Two distinct bands occur at 3509 and 3367 cm ,£i dtaccompanied by a split 1650 cm band assignable to the

C > CO CM CM

COCM o

to,COU~\ CM

LO,

M ON s— voM

HHCQ fQ oM

MM M

cm! o oCM CM CM

COrHCM CM CM

MON

M

CO

{Piuose o u o q . io s q . Y

o

CMIT\ roO

FIGURE

13

CM

CO

ONON

COON

rH

CM

CM MWrHCM

WwM

(©X^oe iCjejq.xqa:v) ’eousqjcoeay^

coo VO

oUN

Orn CM M

FIGURE

14

CO

M

CMM

O'M

O o o

•H

COM

(oi^osoouTjqjOGqy

CM

oo CM CO

CM M

H-O-H bending vibration. This points to the presence of* twotypes of bonded water molecule within the complex. The similarity

-1 , - • 1in profile of the 35^9 cm bond to the 3 -3& cm band ofCuVc,2Ho0 is to be noted, and in both instances the water molecules

/Cj a j

giving rise to these absorptions do not appear to be involved(75 )in appreciable hydrogen-bonding. In contrast, the broad

absorption profiles of CuB^SHgQ and CuSo,3Hg0 shows the presence(75)of nydrogen-bonded water,

k o 2. The Diffuse; Reflectance; Spectra

2n Section 2.3* a set of criteria to distinguish different stereochemistries was presented. Figure 12 shows the diffuse reflectance spectra obtained for CuV ,2Ho0, CuNo,4Ho0 and

m Cj £i Cj

CuSo,3Hr,0 at room temperature. The completes display a mainCi tit “absorption band at 13.8-1^.3 k£ (l.kll = 100G cm *) and a low energy shoulder at 10.1-11.7 kK, which points to the presence of the tetragonal CuO^ chromophore (section 2.3 end Table 1).The main absorption band in the spectrum of CuS ,3Ho0 shows

eC* C.t

extensive asymmetry on its low frequency side, but resolution of this into component bands was not observed when the spectrum was recorded at liquid nitrogen temperature (l*w*. 0 isll (0 .32) and 10.3 (O.23)). The figures in the inner set ofparentheses refer to the absorbance on the arbitrary, 0-2,Unicam S?.700 scale.

The diffuse reflectance spectra of CuBo,2H„0 (recorded£1* C*

at 3-iquid nitrogen temperature) and 0uGo (room temperature) are shewn in Figure 13* Absorption bands ascribed to charge

transfer occur respectively at -21 liK (1,1) and -23 kX (0.75).In contrast to the spectra in Figure 12, the most striking feature is the low energy at which absorption occurs in the visible region. The presence of three bands below 12 kX in each case points to a quasi-tetrahedral stereochemistry (section 2.3, Table 2). If correct, these are rare examples of the quasi-tetrahedral CuO^ chromophore. The resolution of the OuG? spectrum is improved at liquid nitrogen temperature, the bands occurring at -24 kK (0.75), 11.5 kK (0.43), 8.6 kK sh. (0 .53) and 7.5 kK (0 .6l).

Figure 14 shows the diffuse reflectance spectra of CuV , CuNOJ CuB0 and 0uSo recorded at room temperaturej Alld* Cj G* Caare brown (Table 6) on account of the tailing-off of strong ultra-violet bhnds into the visible region. The spectra are distinguished from those in Figure 12 in that absorption in the visible region is much more intense. The tailing-off of the strong ultra-violet bands into the visible region may account for this. Furthermore, d-d transition absorption bands may be swamped, resulting in a single broad absorption envelope being observed. This appears to be the case for CuV and CuN0, but

u Cj '

for CuB0 and CuSr, the existence of two absorption bands is clearly seen at 14.6 kl (1.18), 11.4 IcS sh. (0.95); and 12.5 ks (0.93), 10.6 kK sh. (0.78) respectively. This indicates the presence of the tetragonal CuOg chromophore (section 2.3 and Table 1), which may also be tentatively inferred for the CuV^ and CuN^ complexes. The extent of absorption in the visible region for the complexes in Figure 14 is -8~1‘6 kK, and

consequently the presence of the planar OuO^ chromophore is unlikely. (Section 2.3., extent of absorption envelope for the planar CuO^ chromophore is 14-20 kK.)

4.3. The Magnetic Properties

The effective magnetic moments of the copper(II) complexes at room temperature are given in Table 7. The hydrated complexes have normal moments (-1.9 B.M.) commonly observed for compounds having magnetically isolated copper(II) ions, whereas those for the anhydrous compounds are considerably smaller indicating spin interaction between the copper(ll) ions. Magnetic susceptibilities for the anhydrous complexes, CuB^^H^O and CuVg^SHgQ were recorded over the temperature range 80-30G°K.Table 8 summarises this magnetic data, and with the exception of CuGg and CuB^ the Curie-Weiss Law is obeyed. The small Weiss constants for CuBp,211^0 and CuVg^H^O (0,-20 and -12 °K respectively) coupled with the small variation in magnetic moment (-1.9-1.8 B.M.), confirms the absence of significant magnetic exchange. Similarly, although CuS possesses a low moment, the small Weiss constant indicates no substantial magnetic exchange. The data in Table 8 is meaningful only in a relative sense, since the Gouy balance system was. operated at the limits of its sensitivity, and the hygroscopic compounds were always contaminated with a small amount of hydrated compound. Magnetic field dependence was checked at room temperature. The variation of reciprocal magnetic susceptibility with temperature for the anhydrous compounds, along with the calculated one-dimensional Ising

<H©=L03OS Os 03 IA Os in H CO00 OS OS in VO vo VO VO• • • « • « « •

H H H H H H H H

©Hfita

©

■P©©£1a©aooP5■p(3

a0aw-pfi©a05So•H•p©pjfee(3a©>•rl-PO©

HSOOHXU

tQ•P•rl

.p

HOX

OSOs03

WbO

Os«VOOSH

ITSco•HO•4*

inosC3

VOCO•HOs.

IN 03 in CO H OS OOs Os Os Os OS Os O03 03 03 03 03 03 on

H vo co OS H VO CO» • • • « * • •co O 03 O VO 03 CO •cn 03 -tf* CO vo Os p rs03 01 03 H H H H 03 03

rH 03 03 o O 00 COOs OS 00 co co co CO« • • « • • •OS Os in OS Os cn in03 IN vo os O OS 03<0* cn cn cn cn -=l*

o o O o03 03 03 03a a a a

03 03 cn•*03%

03 03

?03w

o O O o03

%o03

<w©UPIo•H-PO©uuooo•ri+»©fl60(3a(3•H*0!!O

o03 03 03

P a W3 5*o o O H

H

HJ

Ocn

*ON

*“3020203•

03

w II003 •O oH H

dII oCD to

03o0 O o •CO 03 03 HCO CO CO03 03

cno in• H cn •o H vo H03 -4* 4H 02 CM

03O VO• VO 0! «r- H Hin O HH 03 03

Os•4 VO

• CO VO ••4* H cn HON co coH H H

cn4 0 -« CO cn •ON O cn H03 VO vo02 H H

CO•VOCO

oCOVO

II©

03 03r- n- n- c03 02

H /—N v-/H HHI ■r I t-4

a s Oo O 0302 e• oIs- 03 o02 H 03

. 1! 11 tli*~> ►0

03 02 039* *» «*

CO H n-ON in H• • •H 02 03

i; •... „II• « «o 0 O

H H Hd d do 0 o

to to to

H(

OSCO

H• •4' CO

ONcn

•O VO o HOs VO r -

02 03m

» 03 cn

vo•4

•CO cn ON HH 03 01H OJ 02

ON« H o

4in•

VO VO ON Hin CO COH H H

vo

o r-o 00 O VO• CO ON • « cn r^- •

03 02 03 H O in cn HO cn cn o H Hcn H H cn H H

ir\ 03coCO

COinCO

o

ocnON

h*.vo

m%voH

ON

ON0 -

o

ino

« CO CO CO •r - in in CO HH H H OH H H H

VOin 030 O 03 O •

cn co VO Is- Hin ca oa OJH H H H

ON r^. m VO in• r* in « <» 03 co e •in in VO H m 00 in H C2vo 4 4 vo 02 03 VO

02 H H 03 H H 02

cnvoHH

HO03H

cncoHH

n -in

in

co•4co

m

ON Hm in H •4* cn« in in • • 02 vo H • •

cn H H H O o Is- ON H 03ON VO VO Os cn 03 03 ONH H H H H H H H

cn HCO vo ON in H• cn Is- • • 0- Is- O • •CO -4 H H in •4 -4 in H CO02 -4 4 03 03 02 03 0302 H H •• ' Cv3 H H H 02

H•4HH

E3o

H

in03HJII

»”302H•4

03

OH(3O

HI

•4 cnIV 02IV co

00 voON VO H H H H

HH

£oco•

HONH

II03

ino

■03

OH<0O

bO to

• ON ON VOr - m in inH o o oH H H H

CO• CO in 03-4 ON in Oin H H 03H H H H

in• 02 O VOcn ON o COvo o H o03 H H H

H* in CO 03

VO•

cn• r - Oi VO

ON 02 •4 H H o 03 in HON H H H o O o O02 H H H cn H H r*l

TfO

g•rl■PFJOO\

tv- cn o Is*^ o

oo

0202

HO03H

OinHH

VOcn

in•4*

03in

coin

d W

\4 vo o oJH Hoa

%O

oHdod

t-'iVOoH

ft<!-!

0r l

v,<i

o£-i

03

O

dr"i

OHdodk*jri

VO VOOH

OH

• • • 0O o O O

* H rH • H H 0£ d d kr-t d d •ST1■ • o o • o o •

f t f td d f t

d d d d

i4r i kjf"! Mf 1 k i M M ><! ^ i

VO VO VO VO VO vo0 0 o o o 0 0 O o O 0

=L Ch H H H =L H H H r l

C2CD2

Q

H•H 1(0 a+> o.a0 •

bO 03o g H+> *H 1

11+» oH G 030 Pi •O•h a Os<H k •<H 0 H•H in'd -H <v«* IIH 0 0>> 3 o • C\3k cn O HG) H> II G 11

OK © • ©W CO• • • in

CO •ci4 •4* •• n- tv. H

cn cbco cn cn

vo• Os inTJ • 03 4 * •■P Os CM O H

H in vo0 H Ol 030

OCO VO

c0 • in in •VO H n - H

o in O oH H 03 03

G <4*6-i n- VO

• cn in •cn 03 03 HOs n - n -H H H

VOH VO• cn cn eO Os oo Hcn d 403 H H

VO•4* vo• *4' IN •in OS o Hvo 03 cncm H H

CO03 VO

• H H •o r - r - Ho H Hcn H H

cn03CO

« Os •VO cn HCO rv.

O 00• CO •

o OS H03 -ShH cn

invoCO

• in •N cn Hin tv-H 03

inCOCO

• cn ••41 vo Hos 03r * 03

oOSCD

• in •o cn Hcn os03 H

OOOs

• •44 •vo tv. Hvo VO03 H

OHOs

• •4* •03 Os HO •4cn H

in

•v*4IN0o03

1!

©

osIN• cn •rv. •4 1 HCO in

in voCO

e VO •Os 0- HH inH cn

COCOCO

• H •VO O Hin COH 03

Hcnoo• vo t

cn co HOs 03H 03

OsCO« •41 •O 03 H

cn OsOo H

CD03Os• CO «

•4* H HVO rv.03 H

in -41OS• cn •os vo Hos in03 H

G\A vo o OfH H

cm

oHGo e«

I03 03fQPO

CQG 0 GM o

p-ik-itr% O inr-i MVO 03 VO 03 vo

O 0 X 0 o O W 0 o 0H 03 H rL CM -e-i H rl

bOti

•HUH

IIH (D

a)

= Dimeric

aggr

egat

e

<£ > l

12.0

ro.o

' O T10 (Xa)

i.o

6.0

ILmSL

\ CUB2.o\

\

— observed, calculated 0 •

Temperature,° K

0 50 100 150 200 250 >00

alculated ACuV observed observed, calculated 0 observed, calculated X

EO.O CuNCuS

10** (Xa)

i.O

2.0

Temperature, K

25050 100 200

.160)Tnonnorwit** b tderJafe.

./6a). i&ot rniwtc-TaoaodcrWtiie,

Ib(H-)GaB^/tH^O .morioraerHc- ’tnoaaclcribafceN

-HxO

/ 6C3)CuB^ZtfjO|>ol mwi't-,moriodicni'ab2.

- H x OH/

+ tatfcfce.HxO

lb(S)

! TnoiwmoHt - bidentate,

0 < i

Ml)GuBxdiW C-bicWlilte..

M b )CuB^.

imo<tc - Taoaodeata tc,

C H O

0 K > |G(S)GuB^poUjrafiC - bidefttate *

v°

HjiO H*Q HjO

o o

tix.0

fbuo) (ny it cm)CuN^lf.H^O rionrvornmc- bidenfeitc.faonomoWc- rqonodtababe.

I b ( l X )

jjp o feymerric - too acrtk niafce<

16(13)

IbliH)poLpe 'c -'monodcntatc*

variation, equation (A), Section 2.2., is shewn in Figure 15- Clearly Cu6n and CuB0 display 1 strong* antiferromagnetic interaction, whilst for the remaining anhydrous compounds it is weaker. Using the Weiss constant as a guide to the magnitude of interaction, it appears to increase in the order,

CuV2(0 = 102°K) CuN2 (6 = 68°K)^>CuS2 (0 = 30°K).

k0k. Discussion

In the absence of detailed X-ray crystailographic information any discussion can only be of a general nature. The ligands may be monodentate (coordinating through phenolic oxygen) or bidentate and chelating (coordinating through the alkoxy and phenolic oxygen)i As shewn, the infra-red investigation provided no unambiguous evidence in support of either a monodentate or bidentate-chelating coordinating action. (In this discussion ’'bidentate* means *bidentate and chelating1, unless otherwise stated.)

The anhydrous CuG^ complex (no hydrated form was obtained) provides the best starting point for discussion.The visible spectrum showed copper(II) to be in a quasi- tetrahedral environment. Two reasonable structures are possible, Figure 16(1) and 16(2). 16(1) is monomeric with the ligandbidentate, whilst 16(2) is polymeric with the ligand monodentate. In view of the antiferromagnetic interaction displayed by this complex, structure 16(2) is assigned.

The visible spectrum of CuBot2H 0 also indicated a quasi-a* u

tetrahedral environment for copper(II). On this basis possible

structures are 16(3), 16(4), and 16(5) . If 16(3) is correct, then the absence of significant magnetic exchange could be due to the presence of an asymmetric phenolic oxygen bridging network. The long Cur-0 bond, and/or large Cur-Cu distance would prevent magnetic exchange. Dehydration might produce 16(6) in which the copper(II) stereochemistry is planar within the chain, and the bridging symmetric. The shortening of the long Cu— 0 bond, and/or reduction in the Cu— Cu separation associated with symmetric bridging, could account for the appearance of magnetic exchange in CuB^. Axial perturbation from oxygen atoms in adjacent chains, giving a tetragonal CuOg environment, could also account for the observed visible spectrum. If the structure of CuB2,2H20 is based on 16(4) and 16(5), then in order to account for the antiferromagnetism of CuBg, these structures would be expected to undergo rearrangement on dehydration with the formation of bonding links between copper(II) ions so that magnetic exchange can take place. The nature of CuB0 will then be dependent on the juxtapositionC4

of the monomeric molecules within the crystal lattice, and perhaps the most readily visualized d e h y d r a t i o n Process for 16(4) to-the linear polymeric structure 16(6). If the monomeric CuB2,2H20 molecules of 16(5) are suitably positioned within the crystal lattice, then removal of water may be accompanied by a twisting of the chelate rings into the same plane,followed by dimerization 16(7 ) or polymerization 16(8) of the planar molecules, with the intermolecular Cu-0 linkage between them providing a means for magnetic exchange between copper(II) ions.

The NN* ~ ethylenebis (salicylaldimine) copper(II) complex isdimeric in the solid state with a similar bonding link, 16(9)*

(77)between the monomers as in 16(7) and 16(8). Lewis and Walton showed, however, that this linkage did not give rise to any significant magnetic exchange between the copper(II) ions, and suggested that this might be due to the length of the inter- molecular Cu— G bond (2. 4iX) or the Cu—0— Cu bond angle.

It is to be noted that if substantial rearrangement does take place on dehydration, then the magnetic properties of CuB0 may well be dependent on the method of dehydration.tL

( 7ft )This has been found for copper benzoate trihydrate.Of particular interest was a theoretical interpretation

of the observed dependence of magnetic susceptibility on temperature which would differentiate structures 16(6) and 16(8) from structure 16(7) for CuB^. As seen in Table 8 and Figure 15, the magnetic data for CuB^ can be interpreted in terms of the infinite one-dimensional Ising expression, but it can also be explained by the expression for a dimeric aggregate of interacting copper(XI) ions (Table 8 , section 2.2.; Appendix 1.), therefore differentiation is not possible on this basis. Similarly for CuGn, but the dimeric aggregate equation does not give as good an explanation of the magnetic data (Table 8). The calculated Ising g-values for CuB^ and CuG^ are rather high. For copper(II) compounds g-values usually fall in the range 2.10-2.Ao. 79)

CuVg, 2IIo0, CuNp^H^O and CuSp,3HgO have diffuse reflectance spectra which are indicative of a CuO^ tetragonal environment. Possible structures starting from this standpoint are 16(10),

16(11) and 16(12). The infra-red spectrum of CuN o ,4Ho0,C4 tZiand the ease with which it readily partially dehydrates favours structures 16(11) and 16(12) for which the presence of 1lattice* water, as with CuS0 ,3H^Q, is required. The dehydrated complexesCfj &show a variable degree of antiferromagnetic interaction, andif derived from 16(11), it would he expected to arise from theinteraction of planar molecules bonded as in 16(7) and 16(8),If the hydrated complexes are based on structure 16(12), thenthe observed normal room temperature magnetic moments couldbe explained if an asymmetric phenolic oxygen bridging network,16(13), existed (see discussion on CuB^SKgO). Dehydrationof 16(12) with the formation of a symmetric bridging network(see discussion on CuBg,2HgO) could explain the appearance ofmagnetic exchange. The observed visible spectra can beaccounted for if axial Cu— 0 bonds exist between the chains,16(14). A structural and magnetic analogy is provided byCuCl^aKgG and CuCX^. So far no comment has been madeconcerning the stereochemistry about the phenolic oxygen atomin polymeric structures where the ligand is monodentate.In structures such as 16(6), 16(12) and l6(l4) it wouldKprobably be approximately trigonal planar with the benzene rings ofthe ligands twisted by some angle 0 with respect to the Cu— 0— Cuplane so as to minimise steric interaction (16(15)). Examplesof such twisting are afforded by di-(l-(pyridine N-oxide) -bis(di-chlorocopper(XI) ) , 0 = 70.0°, and polybis(p,-(2-picoline H-oxide)-

o C S X )-chlorocopper (II) di-^,~chloro-) diaquocopper(II) , 0 - 88.8 «If CuV^,CuNg and CuS^ have structures based on l6(l4), then the3£More precisely trigonal monopvramidal for 16(6) and l6(l4), probably with the phenolic oxygen sp^ hybridised and the (76'unhybridised p orbital taking part in the long inter-chain bond. J

slight decrease in antiferromagnetic interaction in going from CuVo to ®uM0 can be ascribed to the greater electron withdrawing£U m -

power of -NG0 over -CHO. Steric interaction between the ligand would be expected to increase in CuS^ and this could be alleviated by increasing 0 and/or increasing the Cu—G bond lengths and Cu—Cu distance. The observed decreased magnetic interaction is best accounted for by the latter. Good curve fitting was achieved with the one-dimensional Ising expression (Table 8 and Figure 15), but it can only be considered as tentative evidence in favour of 16(14) (see Appendix 1.).

Attempts to complex copper(II) with 2-ethoxyphenol yielded products contaminated with copper hydroxide, whilst the rapid aerial oxidation ot 2,6-dimethoxyphenol on treatment with alkali, which is greatly enhanced when copper(II) ions are introduced, prevented its complexation.

To clarify the structural situation attempts were made to react copper(II) urith 4-hydroxy-, and 3-hydroxyben- zaldehyde. No complexes were obtained. If the 2-aIkoxy- phenols are chelates in their coordinating action then this is understandable, but not conclusive, since in structures such as 16(2), 16(3), and l6(12), although the all-coxy group is non-coordinating, its importance may lie in other roles, e.g. in the formation of a quasi-chelate species in solution prior to the separation of the solid. The failure of 2-hydroxy-5- -nitrotoluene to complex can similarly be considered.

A mass spectroscopic investigation of CuVo,2Ho0,& sit

CuNo,4Hg0 , CuS^^Hg®, CuBg, and was uswtez’taken. All

complexes other than the latter dehydrated in the source region and decomposed with the expulsion of organic fragments. The spectrum of 0uGo (source temperature 222°C) showed the presence

<*j f

of copper containing species in the vapour state. The highest mass peak occurred at 311 mass units, accompanied by one at 309 mass units. The ratio of the peak height of the latter to the former was 3:2, which is correct for the presence of

63 65Cu and Cu isotopes. The 311 mass peak corresponds to 65Cu and the fact that a mass spectrum is obtained shows that

the Cu—G bond(3) possesses substantial covalent character. The detect.ion of only a monomer Species j CuG0, does not favour structure 16(1) over 16(2), since results for the vapour state cannot be automatically taken to hold for the solid state.

Nickel(II) Complexes. Results, Discussion, and Comparison with Copper(XX)

FIGURE 17»Infra-red absorption in the range 9*00 - 10,00 microns (IIII - 1000 cm-1) for the nickel(ll)

complexes.'

Notes.(1) Transmittance increases in the direction

(2) The scale, I-----1-----1 , is linear in micronsand covers the range 9*00 to 10.00, reading from left to right.

(3) Specific band positions are given in units of*■*1reciprocal wavelength (cm )•

(4) *, see Appendix 2 .

/

__1

■MYg*

I0I.4

*NaV •

1024

1_____1

HV.

102,8

— 2 *

* NaS.

1038

HS ,

1038

^^•2 * 1027

1037

HG.

1040, ,1022

.II

2m L ^ Z H 2.Q.

10

2UM,

1017

Jin,

1021

1010

FIGURE 18. ,Infra-red absorption in the range 7*50 - 8 .50 microns (1333 - 1176 cm"1 ) for the nickel(ll)

complexes•

Notes.(1) Transmittance increases in the direction •

(2) The scale, 1----—1---- 1 , is linear in micronsand covers the range 7*30 to 8 *50, reading from

’ Vleft to right*(3) The 1250 cm"’'1 reference line is shown. See Table 9*

see Appendix 2*

I

NIV12

* J i a Y . . 12^0

HV

NiS2 .

* E a S .

h s .

t

msi"2

TIG .

I

12510

J L L 2 J .

Infr

a-re

d abso

rpti

on

in the

range

2.50-3*50

micr

ons

_1T\

CMCM

CMCM

CM

CMCM

•H

r\M (HM

rO

(—IM

CMIPsinrO

CM CM

CM CM

CMCM

CM

5•1« The infra-red Spectra-1A£5 with the copper(II) complexes the 1111-1000cm and

-11333~ H 76cm regions of the infra-red spectrum were examined in an attempt to decide upon the bonding role of the 2-methoxy group in the nickel(II) complexes (see Figures 1? and i8 respectively)*, The aliphatic-oxygen stretching mode gives rise to an absorption band(s) at 1000~1050cm A, and a band for the aryl-ethereal oxygen stretching vibration occurs at about 1250cm~JL (see Section 2.1.) for the 2-methoxyphenols.

Comparison of Figure 17 and Figure 9 shows that the absorption pattern is not different from that for the corresponding co£3per(II) compounds. However, for the nickel(II) domplexes of 2-methoxy-k-formyXphdnol (IlV) and 2-methoxy-4™nitrophenol (HN) the absorption band associated frith the aliphatic-oxygen stretching mode is shifted to lower frequencies relative to the band position for the corresponding copper(II) complexes. Furthermore, the shift is to frequencies lower than those observed for the ligands and their alkali metal salts (Appendix 2). This is tentative evidence for methoxy group coordination (Section 2.1.). No significant change in band position was found for the remaining nickel(II) compounds compared with the copper(II) counterparts.

In Figure 18 no distinct difference in the absorptionpattern for the nickel(li) and corresponding copper(II) complexes

-1was observed in the range 1333--1176cm For the same reasonsas given in Section 2.1. and 4.1. no attempt at band assignment is made. The absorption band positions for the copper(II) and

Infr

a-re

d ab

sorp

tion

in

the

range

1333-1176

cm

(7.5

0-8»

50

micr

ons)

for

the

Co

pper

(II)

and

Nick

el(I

I)

Comp

lexe

s

a03ONHH

•£ a* 0A03 A

VI•B

•a

O 03 CO osOS O IN. r -H 03 H HH H H H

•a•

w •* •*& 03 a 03 03

IA O COO O 03 H H03 03 03 03 03H H H H H

* V*03 B• •

v. t *. £ A « •ft •»n a 03 03 03 £ a ao o IA CO VD OS IN- A-

co o 03 OS CO r-. •4*Ol 03 03 H 03 H C3 OJH H H H H H H H

•» r* •»m 03 03• • »•* A «. •' *> A A

B 03 03 03 B B 03 03 03CO CO OS C Os VO 03 •4* 03VO co IA C-3 OS IA 03 CO CO03 01 03 03 H 03 03 03 03H H H H H H H H H

•*n*-4•

«h W* •ft *c •ft •* •ft f«m 03 03 to 03 a £ 03 WIN H co O os IN- O oO In . IN- A IA r- Os OSCO 03 03 03 03 03 03 03 03H H H H H H H H H

•• •ft03 03 03 03• * • •

•ft •ft •» •• •• £ Ata 03 03 03 03 03 03 03 03

VO 03 OS IA VO OS IA LA IN.H H OS CO VO o OS O oCO CO 03 03 03 CO 03 CO COH H H H H H H H H

•ft •ft •ft03 03 03• • •

A A A •■* *- •* •ft *»03 03 03 03 03 s 03 03

1 . VO CO O OS H 03 03 OCO 03 CO CO OS 03 CO CO COCO CO CO CO 03 CO CO co COH r t H H H H H H H

• • % • *O O O O O

03 03 03 03 03w w S3 S3 &03 03 «sF 03

•ft •* • 0 W •»' •». «*03 03 03 03 « 03 03 03> > > > • > s f-J* tr=*•H 2 •H t> (3 ♦H •rlO o $3 fa £; o J25 &

TJQ>•H+»SO

•>>

0to

•> •ft A

03 03 03VO VD VDCO IN- OSH H HH H H

e* ••03 a• 0

•* A r»03 03 03 03CO O H 03OS O 03 03H 03 03 03H H H H

•ft •03 £• 0

* « A £03 03 a 03 03

03 IN- o VD O•4< H CO LA o03 03 03 03 03H H H H H

03,

• r. •• « •fta 03 03 a 03 a 03IA IN 03 H IN CO OIN- VD VD Cs COH 03 03 03 03 H OlH H H H H H H

•ft •»a 03• •

• •* A •*. •»£ a 03 03 03 03 03 03OS LA VO O IN. 03 CO OsIN. O VO CO VD H IA IAH 03 03 03 03 CO 03 03H H H H H H H H

** •ft •ft •fta" 03 B 03 03 03 03 03IA O! OS Jj* CO VD LA

H 00 O OS O H SS- O03 OS CO 03 CO CO CO COH H H H H H H H

1 03•

•ft •ft •** A «•03 3 03 03 a 03 03 ►ft*

H 03 CO CO o 03 OS COo - CO CO CO CO CO 00 0303 CO CO CO CO CO 03 COH H H H H H H H

•o03

ss03

• « •ft • ft03 03 V>4»K 03 03 •

fe & • ft •p •H }2 3 0 (0

O & fa fa O o fa

ta toVOis. iv-H HH H

TJ4JrtOoON0)H,Qrt

*» •tO W *CM H HH H CO03 CM HH H H

r*ft to% •A rrt • •*« to £ £

O CO ON COCM H CO 00CM CM H HH H H H•»

to•• A *• • « •* «to to to * V S W aCM co cn co cn H VD inIs- in in ON CO CM O 00H CM CM H H CM CM HH H H H H H rrt H

*•to to

r»© •to

•Se S

#*S*in cn cn VO CO ON CMH vo vo 0 ON CM CMCM CM CM CM H CM CMH H H H H H H

rt©£uV*nrt•rt*0©2

a* *T £ bO• • • rt

.C A A Oto to to u

CO O O pin CM ON 0CM CM HH H H 11

r* •* •* t*to to to 10 SB * a to a PH m O CM O O VO CM in •rtvo CO co CM cn IS. in IN in H toCM CM CM CM CM CM CM CM CM CM rtH H

s'•

H H H H H H

«r»to•

H H ©prt•H*0,rt «. « •» ,rt rt

to to to to 5* to SB to £ s rt•c* On in H ON in VO ON ACO ON CO VO VO CM Is- H Is- COCM CM CM CM CM cn CM cn CM CM <HH H H H H H H H H H 0

torto•rtP

0* •* «* •• m rta to a a to to to to w a 0CO H H 0 H cn cn VO CO cn •rtH CM CM 0 CM cn CM CM CM CM V •

<n cn cn cn cn cn cn cn cn cn • rtH H H trt H H rrt H H H CM

iN•rt©

©•0H• • • •rt Is rtO O 0 •a •rt 0CM CM CM rt P irt

S3 ►r»K’-l S3 © rt wCM cn cn S3 p. p•* % • •V « « Ph •rt 1!CM CM CM CM CM CM CM • << HCJ CJ « t/3 to to • w rt «•rt •rt rt eJ r- •rt rt •rt w rt rtS3 S3 O W O S3 0 S3 S3 S3 w O' to

nickel(II) complexes in this range are catalogued in Table 9*Figure 19 shows the infra-red absorption for the hydrated

nickel(II) complexes in the 0-H stretching region at about three microns* Like CuN ,4 h o0 ? NiNo,4HL0 displays an

M Cil ^interesting absorption profile in this region, the sharp highfrequency bands at 3552 and 3484cm and the broad band centred

-1at 3195cm indicating water molecules of two types in the compound(compare CuN_,4Ho0. Section 4.1.)* Thermogravimetricd* 2analysis established the presence of two types of water molecule(Figure 8), but rather surprisingly the infra-red spectrumof NiNp,2HgO is almost identical to that of NiN^, 4&g0 *.Consequently? the assignment of two classes of water molecule inCuNg^HgO on the basis of the infra-red evidence alone, canonly be tentative. The remaining hydrates display broadabsorption bands in the three micron region characteristic

(75)of hydrogen-bonded water.

5*2. The Diffuse Reflectance Spectra and Magnetism

The spectral data for the nickel(II) complexes is collectedin Table 10 and the spectra shewn in Figure 20. In light ofthe discussion given in Section 2.3* on six coordinate nickel(II)complexes, these spectra are consistent with an octahedralstereochemistry (weak ligand field ) . It is doubtful whetherthe low energy absorption band in the spectrum of NiSg(«*5kK,Figure 20) arises from a d-d transition. The band can be

3 3assigned to the A„ (F) - T_ (F) transition and the split or2g 2g3 3asymmetric V band to the A (F) - T^ (F) transition. Possible

Diffuse

Refl

ecta

nce

Spectra

(kK)

of the

Ni

ckel

(IX)

Co

mple

xeXo «oca

b

IQ*dPidm

03

*d§o$oo

X• 03A •tQ COX C3O• •ft

ft o ISA CA LAb A-r» •• X VO

IA CO CA 03*4< •41 03 •« • CO •»

A- A - 03 •03 03 • A

,p •» tQ•• «* £Q X Xco a- LA in- CO• « « • •

vo vo A- vo LA03 03 CQ 03 OJ

tQ •VO ON O .P

• • • • IQA CA CA -4< COIQ 03 03 03 •VO H

• C* *• 03LA • • •03 rP .P A 9s

I tfl CQ B LA XVO VO vo H CQ O CO

• • • • • • ••4* -41 H 03 H H «4*03 03 03 03 03 03 03

HI X x x X00 vo LA o o« « ft ♦ ft : ftCA CA CA CA o o03 03 03 03 CA CA

•A• ft m,P >p X

tQ tQ OH •* * LA Ho CA CO 0303 03 •

03 •*ft* 03 •• • • • • *P

»P .P »P A A 9s tQCQ CQ tQ tQ X X m LA X X•4* 03 o -4* O CO vo CA O -4<• • • • • • • • * •ON ON ON ON 03 H H H O ONH H H H 03 03 03 03 03 H

,pto

03 CA • • 0 • » •• • A ,P ftp rP rP rP e •

LA LA n IQ B B tQ tQ -4 1 •41H H o 03 03 03 CO co O vo O A*- H H O CAs • • • t » * ft * • • •■„ *» -4* *4« -4< CA CA LA LA LA LA •* 9s IA IA• • H H H H H H H H H H • ft H Hr'»■-« A rP J5ftM

tQ tQ r» •*■ •ft 9s 9s 9* *» 9* •» tQ tQvo VO <*< co 03 (A A - vo LA O CA CO ON LA vo CA• • • » • • « • • • « • « • t • •CA CA 03 03 CA CA 03 CQ CA CA CA CA 03 03 03 CM 03H H H H H H H H H H H H H H H H H

03H

LAIN. vo CA LA co O LA tN- 03 03 •4< -4< o CA COft • • » • « • • • • • • • • • ft ft

00 CO A- r- A- co IN- (N- co ON ON ON A- r- ON ON IN- tN-

O O O O O03 03 03 03 CQK X w . w

03 03 -4* 03 CAws #• f* *• «ft03 CQ 03 03 C3 03 CQ CQ CQCJ CJ > > fe; X hr W CO

•H •H •H •H •H •H •H •H •H& 15 !5 15 Hr#£-l 15

S3s+3ooPitQG>PiP•PdPiOPiae■P

aooPi

Pio*dHdo*Ptn

•Pto

as■po0PitQoPi0•PdPiePiao+>p0bOOPi+3•HP'd•HPOr1•HH

II II15

tQTJPd»Q©«P©•{■3pH

X

.0-51

okJCOs °J_5<£uloz<CQa:oC/1c£)<

/•X

10

*Lo- fe

o>8.

2.0

Effe

ctiv

e Ma

gnet

ic

Moments

(B,M.)

at Room

Temp

erat

ure

102

H O IA 03 CO CO in CIN03 03 03 CQ 03 03 CA CC 03

« • • • • • • • •CA ca CA CA CA CA CA CA CA

IA n - ON CA 03 C3 C3 o0 I CA ON co ON ON ON ON ON ONH 03 03 03 03 03 03 03 03 03

VO

ON ON H H H 00 00• t . • • • • • • ♦00 CCS VO o CO 03 vo ca CACO vo ON CN CA H CO O «H H H H 03 03 H ca 03 /—N

03r-

«HO

•+> o IN- 03 ON CA O IN ON IAl-r. o ON o ON O o ON o O

« • « • « « • • •• H 1^ o n - H -tf1 IA H

H O ON vo VQ CA ON r - caO

SCA CA CA CA •4"1 CA -=h

•H•PO©uuooo•H+>0PiM<2a(tf•H

•0 o O O O Opi 03 ca 03 03 03$ w X « « W0 03 ca ca CAPi •» •*S 03 ca ca 03 03 03 oi ca0 C3 O > & X X too •rl •H •H •rl •H •rt •ri *H& J2S

03WJ2S

«OStQ

O = Ni

21(1 )

21(2)

21(3)

2 1 ( 4 - )

causes for this splitting and asymmetry were considered in Section 2*3# The substantial lowering of ligand field strength, as indicated by the position of (V^oct. = Ao, Section 2.3#)» in the NiG , NiN and NiS complexes,is to be noted.Cm cL Ci

Supporting evidence for the assignment of octahedral stereochemistry comes from the effective magnetic moments at room temperature (Table 11). The magnetic moments fall within the range normally observed for octahedral nickel(II), 2.9 - 3#^ B.M

5.3. Discussion and Comparison

Unlike copper(II), nickel(II) produces octahedral hydrated and anhydrous complexes with 2-methoxyphenol. Furthermore, nickel(II) fails to yield a pure complex with 2-ethoxy~A— formylpheno

The hydrated copper(II) complexes of HV, HIT, and HS were assigned a distorted octahedral stereochemistry on the basis of their visible spectra, and in addition, this stereochemistry was tentatively inferred for the anhydrous complexes. A parallelism can be drawn with the nickel(II) complexes of these ligands, but here the evidence for an octahedral environment in the hydrated and anhydrous complexes is more definite.

In contrast to copper(II), there is tentative evidence for methoxy group coordination in the nickel(II) complexes of 2-methoxy-4-nitrophenol and 2-methoxy~4-formylphenol (Section 5*1#)# A possible structure, if this is correct, for the hydrated complexes of these ligands is Figure 21(1), with the addition of •lattice* water for NiN^, 4HgO. To account for the observed effective magnetic moments and visible spectra of the anhydrous

compounds MxV and NiNA., dehydration must be accompanied by polymerization. A possible structure is 2 1 ( 2 ) . To attempt to explain the large reduction in ligand field strength which accompanies the complete dehydration of NiN^,^R^O,' compared with the smaller reduction in ligand field strength observed on dehydrating NiV^j 2K/>G, would be unwise, especially in the absence of detailed structural information.

In the remaining nickel(II) complexes the ligands may well be bidentate and chelating (see Section 4.4.), and if so, the same considerations as given above apply. However, if the ligands are monodentate (see Section 4,4.) then a structure based on 2l(3) is possible for the hydrates (plus 'lattice1 water in the case of NiS^^H^C). Dehydration with the formation of inter-chain axial Ni-0 bonds would account for the observed

/ O o \visible spectra' ; and magnetic properties of the anhydrous complexes (21(4)). As each phenolic oxygen atom is shared between three nickel(II) ions in 21(4) (in structure 21(2) the unspecified oxygen atom involved in the intermoiecular linkage is shared between two nickel(II) ions) this might account for the reduction in ligand field strength, since the availability of lone pair electrons to the nickel(II) ions should be reduced, A similar structural argument holds for the nickel(II) complexes of 2-Eiethoxy~4-formylphenol and 2-methoxy~4~nitrophenol if the ligands are monodentate, but the problem of the small change in ligand field strength in going from hiV^, 21- 0 to NiV^,, as compared with the large decrease in ligand field strength observed on the complete dehydration of the other nickel(II) complexes, would again be difficult to reconcile.

5ok . Concluding Remarks

The crux of the structural problem has been whether or not there is alkoxy group coordination. CuVA,2'Ho0« CuV ,

C m d» c m

CuB , 2 Ho0 , CuB and NiV , 2 Ho0 are mierocrystalline, and ifQ M U A M

suitable single crystals could be obtained,an X-ray crystallo- graphic examination would resolve the structural problem. A structure determination coupled with a single crystal magnetic anisotropy study would be immensely helpful in trying to understand the magnetic exchange found in the anhydrous CuV and CuB^ complexes*

«*« 4. i

Appendices

6.1. Appendix 1, Magnetism

No reasonable interpretation of* the magnetic data TorCuB£ and CuG^ was achieved using the results of Earnshaw et

(22)al. (n=10). However, this data could be satisfactorilyconsidered in terms of the magnetic susceptibility expression for a dimeric aggregate of interacting copper(II) ions, vis.,

Xa = - & -§■■■ x r 73-7 + Nog (See section 2.2.)3kT 1 + 1/3 oxp(2J >!r.T

The following curve fitting procedure was adopted. Usingthe observed temperature of maximum susceptibility, 2J was

(72)evaluated; 2J = 8/5Tmax. g was calculated from themagnetic susceptibility equation using Xa.max., Tmax., 2Jcal<p.,

and Na = 60 x 10~^ e.g.s. units. If satisfactory curve fitting with these values was not obtained, g was varied.(S ee section 4.3-).

Reference was made in Section 2.2. to the fact that Adams et a l . ^ 0 had utilized the one-dimensional Ising expressions for the parallel and perpendicular susceptibility of an infinite chain of interacting copper(II) ions for interpreting the magnetic data of anhydrous cupric chloride, vis. ,

X(T) = 1/3 X--(T) + 2/3 X- (T) ................... (B)powder

The expressiven for X1 in reference (30) is wrong. It should be

X .( T ) = f f j f “ ( t a n hJ « JISr T S r sech8 ^p)

Sinn has confirmed t h i s . ^ ^ Using the literature magnotic

/Or O (L \data * for anhydrous cupric chloride (observed temperature of maximum magnetic susceptibility is 69°X) reasonable curve fitting was achieved using expression (B), ( g = 2.13;2J = -110.0 cm **■; temperature range considered, 20-300°K; percentage deviation from literature data +5.5%). The magnetic susceptibility data for CuB^, cou**-c* no^ ^e i't’ked toequation (B), but the data for CuV , CuN_, and 0uSo was

Ca tL cj

satisfactorily fitted. The calculated g* and J * values are compared with those Using only the parallel esq>ression, g and J (Table 8).

s' 2 J 1 cm g 2J cnCuV2, 2.13 -88.0 2.22 -39.3CuN 2 , 1.96 -72.0 1.98 -27.2

CuS2* 1.96 -30.0 1.94 -12.4

Only with additional magnetic data extending below the temperature of maximum susceptibility might differentiation between these two approaches be achieved, but a$shewn for CuB^ and CuG^, any structural deduction based on the criterion of reasonable curve fitting must be tentative. Logically, theoretical discussions of magnetic data should follow an X-ray structure determination and a single crystal magnetic anisotropy study (see tetraammine-copper(II) sulphate monohydrat e .

Computer curve fitting procedure for equation (B).Equation (B) was solved for g using an esq)erimentally

observed value of Xa at temperature T, tisually the lowest~6temperature at which measurements were made, Noc = 60 x 10 e.g.s.

units, and an arbitrarily chosen value for J. Xa was then calculated using g-calculated and J, at various values of T. This represented one cycle of the programme, which was then repeated for a whole series of J values. The computed and observed magnetic susceptibilities were manually compared.

(Units; for J in cm \ k = 0.69499* NB^ = 0.26072)

6.2. Appendix 2. Infra-red,

All band positions reported in this thesis were takenfrom the spectra recorded on the Grubb Parsons spectrophotometer,

(75)Nujol bands were used for internal calibration, and theassignment of these bands confirmed by comparison with the spectra obtained using hexachlorcbutadiene as the mulling agent.

The sodium salts of HV, HS, and KB could not be obtainedpure because of contamination with the free acid or solvent, andsodium carbonate. Although the sodium salts of HV and HS areformulated as NaV and NaS they are in fact solvated (infra-redspectroscopy). In the case of NaV it was found that onstanding under absolute ethanol overnight the infra-red spectrumwas unchanged except for the appearence of a secondary peak at1032 cra~\ but if ground under absolute ethanol, dehydrationoccurred, accompanied by a considerable simplification of the

-1spectrum. A single band now appeared at 1032 cm in the region 1000-1050 cm

The weak sharp infra-red absorption band observed for KN at 3559 cm was initially attributed to a small amount of

z 0 (19Kpotassium hydroxide impurity (H-Q stretching vibration ), but a combined t.g.a. and infra-red study showed that this

-1band originated from water of hydration. The weak 3559 cm and 1658 cm*’’1" (H-O-H bending vibration) absorption bands observed in the hydrated compound were absent in the anhydrous compound.

(Required for ICM,H20: H 20 8.0%; Found 7.6%)

7. R E F E R E N C E S

(1) C.Ne Tyron and S„C. Adams, J.Am.Chem.Soc., 1940,62,1228.(2) C.M. Harris and E. Sinn, J .Inorg.Nuc1.Chem., 1968,^0,2723.(3) D.P. Graddon and G.M. Mockler, Aust.J.Chem., 1968,21,617.(4) P.R, Shukla, J .Inorg.Nuc1.Chem., 1967,.29,1300.(5) F.A. Cotton and R.H. Holm, J .Am.Chem.Soc., I960,32,2979.(6) W. I. Taylor* and A.R. Battersby, *0xidative Coupling of

Phenols *, E. Arnold (London), M. Dekker Inc. (N.York), 1967.(7) R. Dolique and R. Mestres, C„A., 48:8008c.(8) C.S. Prout et al., J .Chem.Soc., (A), 1968, 2791.(9) P.L. Orioli and M. Di Vaira, J .Chem.Soc., (A), 1968, 2078.

(1G) A.V. Butcher, Ph.D . (Thesis), 1968, University of Surrey.(11) G.W. Watt and F„0. Drummond, Inorg.Nuc 1.Chem. Le 11 er s, 1968,4^551(12) A.R. Eatritzky and N.A. Coats, J.Chem.Soc.. 1959, 2062.(13) L.K, Briggs et al., Anal.Chem., 1957,29,904.(14) M. Eato, H.B. Jonassen and J.C. Fanning, Chem.Rev., 1964,99-(15) B.N, Figgis and J. Lewis, 1Progress in Inorganic Chemistry1,

1964,6^, Interscience.(16) A. Earnshaw, introduction to Magnetochemistry1t Academic

Press, 1968.(1?) B. Bleaney and K.D. Bowers, Proc.R.Soc.. 1952,214A,4gl.(18) D.M.L. Goodgame and D.F. Marsham, J .Chem.Soc.. (A), 1966,1167.(19) J. Lewis et al., J.Chem.Soc.. 1965, 6464.(20) W.E, Hatfield and J.C. Morrison, Inorg.Chem. , 1966,^,1390.(21) C.M. Harris and E. Sinn, J .Inorg.NucI.Chem., 1968,30,2723;

G.A. Barclay et al., Proc.Chem.Soc., I96I, 264.(22) A. Eamshaw, B.N. Figgis and J. Lewis, J.Chem. Soc., (A),

1966, 1656.

(23) C.J. Ballhausen, 1 Introduction to Ligand Field Theory1, McGraw-Hill, New York, 1962.

(2(0 S. Itatsura, Phys.Rev., 1962,127 % 1508 ; and references therein contained.

(25) E. Ising, Z.Physik, 1925,31,253.(26) J.C. Bonner and M.E. Fisher, Phys.Rev., 1964,135«A64G.(27) G.F. Newell and E.W. Montroll, Rev.Mod.Phys., 1953,25,, 353 .(28) C.G. Barraclough and C.F. Ng, Trans.Farad.Soc., 1964,60,836.(29) M.S. Fisher, J.Math.Phys., 1963,4,, 124.(30) R.¥. Adains et al., Aust. J.Chem., 1967« 20, 235i»(31) M. Inoue et al. , Inorg. Chem., 1968, 7,, 1427.(32) H. Kobayashi, J .Phys. Soc .Japan, 1963,18,, 349.(33) R.W. Adams, Private Communication, 1968.(34) R.B. Griffiths, Phys.Rev., 1964,135,Ao59.(35) 33* Dubicki, Inorg.Chem., 1966,5., 93.(36) IC.E. Hyde et al., J. Inor g. Nuc 1. Chem., I96S, 30* 2155*(37) A.P. Ginsberg et al., J.Inorg.Nuc1.Chem., 1967» 29,353»(3B) E.A.V. Ebsvrorth et al., !New Pathways in Inorganic Chemistry1,

Cambridge University Press, 1968, p.175*(39) ICato et al. , Bull.Chem.Soc. Japan, 1968,41,1864.(40) M.E. Fisher, Private Communication, 1968.(41) D.P. Graddon and G.M. Mockler, Aust.J.Chem., 1967,20,21.(42) H. Elliot and B.J. Hathaway, Inorg.Chem., 1966,^,885.(43) I.M. Procter et al., J.Chem.Soc., (A), 1968, 1678.(44) A.A.G. Tomlinson and B.J. Hathaway, J.Chem.Soc., (A),

1968, 1685.(45) A.A.G. Tomlinson and B.J. Hathaway, J.Chem.Soc., (A),

1968, 1905.

(46) A.A.G. Tomlinson and B.J. Hathaway, J.Chem.Soc., (A), 1968, 2578.

(47) A.A.G. Tomlinson et al., J.Chem.Soc.. (A), 1969, 65.(48) D.E. Billing et al., J.Chem. Soc. , (A), 1969, 33-6.(49) D.E. Billing and B.J. Hathax^ay, J . Chem. Soc.« (A), 1968, 15X6(50) A.B.P. Lever, 1Inorganic Electronic Spectroscopy1, EXesevier

1968.(51) B.N. Figgis et al., J .Chem.Soc., (A), 1968, 2028.(52) W.E. Hatfield and T.S. Piper, Inorg. Chem., 1964,3., 84l.(53) A.G. Earipides and T.S. Piper, Inorg.Chen., 1962,1., 970.(54) D.P. Graddon and G.M. Mockler, Aust. J . Chem. , 3.968, 21, 617.(55) M.G. Clark and R.G. Burns, J.Chem.Soc., (A), 1967, 1034.(56) T.M. Dunn, 1Modern Coordination Chemistryt, edited by

J. Lewis and R.G. Wilkins, Interscience, 1964, p.278.(57) J.S. Griffith, fThe Theory of Transition-Metal Ions11

Cambridge University Press, 1961.(58) C.E. Jorgensen, Acta.Chem.Scand., 1955,2,1362.(59) A. Ludi and ¥. Feitknecht, Helv.Chim.Acta. , 19631 46, 2226.(60) D.M.L. Goodgame et al., J . Chem. Soc. , (A), 1966, 1769*(61) F.A. Cotton and G. Wilkinson, 1 Advanced Inorganic Chemistry1

2nd edition, Interscience, 1966.(62) D.M.L. Goodgame et al., J.Am.Chem.Soc., 1961,83,4l6l.(63) F.A. Cotton and J.P. Fackler, Jr., J.Am.Chem.Soc.,

1961,83,2818.(64) J.P. Fackler, Jr., and F.A. Cotton, J .Am.Chem.Soc.,

1961,83,3775.(65) F. Pollecoff and R. Robinson, J.Chem.Soc., 1918,645.

(66) D.A. Shirley, fPreparation of Organic Intermediates1,J, Wiley and Sons, Inc., New York, 1951* p.203.

(6?) T.S. West et al., Chem. and Ind., 1957, 1647.(68) R. Belcher and C.L. Wilson, *New Methods of Analytical

Chemistry1, Chapman and Hall, 1964.(69) A.I. Vogel, 1 Quantitative Inorganic Analysis*, 3rd Edition,

Longman s, 196 2.(70) G.M. Lukaszewski and J.P. Redfern, Lab.Practice, 1961, 469.(71) A.W. Coats and J.P. Redfern, Analyst, 1963,88,906.(72) J. Lewis and R.G. Wilkins, 1Modern Coordination Chemistry1,

Interscience, 1964, Chap.6.(73) A. Earnshaw, 8Introductinn to Magnetochemistry*, Academic

Press, 1968, Chap.VI, p.90, Figure 43.(74) H.R. IJettleton and S. Sugden, Proc.Roy.Soc. , (A) , 1939* 173* 313«■(75) S. Nakanishi, 8 Infra-red Absorption Spectroscopy1 ,

HoIden-Day, 1964.(76) D. Hall and T.N. Waters, J.Chem.Soc., i960, 2644.(77) J« Lewis and R.A. Walton, J.Chem.Soc., (A), 1966, 1559.(78) C.S. Fountain and W.E. Hatfield, Inorg.Chem., 1965,^4,1368.(79) D.J.E. Ingram, 1 Spectroscopy at Radio and Microwave

Frequencies 8, Butterworths, 1967.(80) A.F. Wells, 8Structural Inorganic Chemistry1, Oxford

University Press, 1962.(81) R.S. Sager and W.H. Watson, Inorg.Chem. , 1968,7., 2035 •(82) D.A. Langs and C.R. Hare, Chem.Comrnun., 1967, 853*(83) L. Sacconi, ♦Transition Metal Chemistry1, 1968, _4, 199*

(84) S„ Sinn, Private Communication, 1969..(85) W.J. De Haas and C.J. Gorter, Conm. Phys. Lab. Univ.

Leiden, 1931, 215a, 3«(86) C. Staar et al., Phys. Rev. , 194-0, 58, 977 •