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So1vat~on. React~v~~y and Spectroscopy
of Co:m.p1e:xes of
SO:J:D.e F~rs·t Ro,,", Transit:ton M:e~a1s
A Thes~s sub~~tted by
, STOJAN RADULOVIC
for
the degree of
DOCTOR OF PHILOSOPHY
:In the
FACULTY OF SCIENCE
of the
UNIVERSITY OF LEICESTER
Inorganic Chemistry Laboratory. Department of Chemistry, The University. LEICESTER LEI 7RH
May 1988
Solvation, Reactivity and Spectroscopy of Complexes of Some First Row Transition Metals
Stojan Radulovic
ABSTRACT
The work in this thesis is mainly concerned with the discussion of effect of solvation on the reaction kinetics of inorganic complexes. Rate constants for chemical reactions in various aqueous cosolvent systems have been measured and analysed.
Crystal structure of several Fe(II) complexes are analysed from crystallographic data for possible structural parameters which might have effect on solvation. Solubility data, for a range of inorganic salts containing simple and complex ions, are reported for aqueous solutions and for solutions in aqueous cosolvent mixtures. Transfer chemical potentials for single ions in aqueous i-PrOH and t-BuOH solvent mixtures are determined using solubility data for salts in conjunction with TATB,tetraphenylarsonium tetraphenylboronate, assumption and are compared with those in corresponding aqueous methanol, ethanol and acetone solvent mixtures.
Kinetic data are reported for reaction between [Fe(gmi)3]2+ and hydroxide ions at atmospheric and elevated pressures in above binary aqueous mixtures. Initial state and transition state analysis of reactivity trends for hydroxide attack on other Fe(II) diimine complexes in aqueous methanol solvent mixtures are reported.
Dependence of visible absorption spectra on solvent has been examined for a number of Fe(II) and Fe(III) mixed ligand complexes. Preliminary redox study of the Fe(II) and Fe(III) complexes is also reported.
Ocu. :m.a.~ ki •
i
supr-ugi
cer-kici Rusan.di
"Tvrd j e orab voc.ka cudnova ta -
De salomi ga, a1 I zube po1o.mJ.!"
Petar Petrovic Ijegos
'Gorski Venae'
ACKNOWLEDGEMENTS
I would like to thank my supervisor, Dr. John Burgess, for
guidance and encouragement throughout the period of this research. Also
I would like to thank Dr. Michael J. Blandamer for assistance and
discussions.
Further, I would like to thank the British Council and Accion
Intergrada for the travel grant to the University of Seville (Spain).
STATEMENT
This thesis is based upon work conducted by the author, in the
Department of Chemistry of the University of Leicester, during the
period January 1984 and September 1986.
All the work recorded in this thesis is original unless otherwise
acknowledged in the text or by references. This work is not being
presented for any other degree.
May 1988 stojan Radulovic
University of Leicester
LIST OF CONTENTS
CHAPTER 1 - INTRODUCTION
1.1 1.2 1.3
Introduction Properties of binary aqueous mixtures Analysis of medium effect an reactivity 1.3.1 Initial state trends and solubility data 1.3.2 Initial state-transition state salvation REFERENCES
CHAPTER 2 - EXPERIMENTAL
2.1 2.2 2.3 2.4 2.5 2.6 2.7 2.8
Introduction Solubility measurements First order rate constant SP 800 Spectrophotometer SP 8-100 Spectrophotometer SP 1800 and HP 8451A Diode Array Spectrophotometers High pressure kinetics Atomic absorption spectrometry REFERENCES
CHAPTER 3 - CRYSTAL STRUCTURE OF IRON(II) COMPLEXES
3.1 3.2 3.3 3.4
Introduction Experimental Results Discussion 3.4.1 structure 3.4.2 The coordination polyhedron 3.4.3 The imine moiety REFERENCES
CHAPTER 4 - SOLUBILITY OF SALTS AND DERIVATION OF TRANSFER CHEMICAL POTENTIAL IN SEVERAL BINARY COSOLVENT SYSTEMS
4.1 4.2
4.3 4.4 4.5 4.6
Introduction Experimental 4.2.1 Preparation of compounds 4.2.2 Solubility measurements Results and discussion for aqueous methanal Results and discussion for aqueous i-PrOH and t-BuOH Calculations of 6m~a(OH-) Discussion REFERENCES
Page No.
1 4 9
11 13 16
18 18 19 21 21 23 24 24 26
27 28 30 36 36 37 38 43
45 46 46 48 48 65 75 78 87
CHAPTER 5 - AMBIENT AND HIGH PRESSURE KINETICS OF REACTIONS BETWEEN HYDROXIDE IONS AND IRONCII) DIIMINE COMPLEXES
5.1 Introduction 5.2 Experimental 5.3 Results 5.4 Discussion 5.5 High pressure kinetics 5.6 Experimental 5.7 Results 5.8 Discussion
5.8.1 Reaction in water 5.8.2 Reaction in binary systems REFERENCES
CHAPTER 6 - SOLVATION AND REACTIVITY OF SEMIAROMATIC IRONCII) COMPLEXES
6.1 6.2 6.3
Introduction Experimental Results and discussion 6.3.1 Transfer chemical potentials 6.3.2 Reactivity and initial state-transition
state analysis REFERENCES
CHAPTER 7 - SOLVATOCHROMISM AND SOLVATION OF IRONCII) AND IRONCIII) TERNARY COMPLEXES
7.1 7.2 7.3 7.4
7.5
Introduction Experimental Results Discussion 7.4.1 FeCII) dicyano and tetracyano complexes 7.4.2 FeCIII) dicyano and tetracyano complexes Transfer chemical potentials of ternary complexes REFERENCES
CHAPTER 8 - REDOX REACTIONS OF IRON(II) AND IRON(III) TERNARY COMPLEXES
8.1 8.2 8.3
Introduction Experimental Results and Discussion 8.3.1 Oxidation of catechols by [Fe(bipy) (CN).]-
89 90 91 91
108 109 110 116 116 117 122
123 124 125 127
133 144
145 146 150 156 157 159 160 168
169 170 171
and [IrCl s ]2- 172 8.3.2 Peroxodisulphate oxidation of Fe(II)
ternary complexes 181 8.3.3 Ligand oxidation in Fe(II) hexadentate complexes 184 REFERENCES 188
APPENDICES 189
SOXE LIGAID ABBREVIATIOBS
bipy
cxcage
en
mrl
phen
tsbXe
2,2'-bipyridyl
biacetyl-bis-B-methylimine
1,2-cyclohexanedione-bis-l-methylimine
5,6,14,15,20,21-triscyclohexane-1,3,4,7,8,10,12,13,16)7,19,22--dodecaazatetracyxlo-[8.8.4.13 . 17.19 • 12] tetracosa-4, 6, 1 3,15, 19, 21-hexaene-J4,I',113,116,119,B22
ethylene diaDine
glyoxal-bis-B-methylimine <2,5-diaza-2,4-hexadiene>
3, 14-dimethyl-4,7, 10, 13-tetraazahexadeca-3, 13-diene-2, 15-dione dioxime
methylglyoxal-bls-J-methylimine
methyl 2-pyridylketiDine
1, 1 O-phenanthroline
phenyl 2-pyridyl ketimine
2,6-diacetylpyridine-bis-B-methylimine
OTHER ABBREVIATIONS
DIfSO dimethilsulphoxide
HtOH ethanol
i-PrOH iso-propanol
IS initial state
l.p solubility product
XeOH :methanol
pL ionization constant of water
TA typically aqueous
t-BuOH tertiary butanol
TS transition state
CHAPTER
1
Introduction
1.1 INTRODUCTION
There has been much interest for quite some time in the influence
of the solvent on the rate constant of a chemical reaction. In solvents
comprising binary aqueous mixtures c1 -.) the rate constant dependence on
the solvent composition, whether small or dramatic, is often complicated
with only a relatively few patterns emerging. In terms of transition
state theoryC&>, the complexity reflects the separate influences of
changes in solvent composition on initial and transition states.
Therefore, independent estimates are sought of the effects of solvents
on initial states, the corresponding effects of solvent on transition
state are calculated using both kinetic and initial state data.
However, the difficulty arises in analysing solvent effects on rate
constants for chemical reactions involving ions, as in such cases
extrathermodymanic assumptions have to be invoked in order to calculate
single ion properties from measured properties of salts.
The work in this thesis discusses solvation of transition metal
complexes and their reactivity with ions in aqueous and binary aqueous
mixtures. The inorganic complexes studied are all ideal for
conventional kinetic monitoring due to the chemical inertness of the
electronic configuration of the central metal atom, low-spin d6• The
reactions were all bimolecular which enabled parameters, such as
solubility, to be measured on individual solutes prior to kinetic
determination. Derivation of transfer parameters provides an insight
into solvation of simple·and complex ions in binary aqueous mixtures.
Solvation in terms of charge, ligand structure and
- 1 -
hydrophobic/hydrophilic character is hoped to provide a basis for
analysis of kinetic data for reactions involving ions in such mixtures.
Discussion is primarily concerned with low-spin (t206) Fe(II)
complexes containing ligands which have a large crystal field effect.
Amongst such ligands are cyanide ion and a series of organic nitrogen
bases, e.g. 1,10-phenanthro1ine and 2,2'-bipyridyl, 1 and a
respectively. The existence of stable low-spin Fe(II) complexes with
these hetero-aromatic ligands has been known for almost a century.
1
The characteristic of these ligands is the diimine moiety~. This
structural element is also present in the aliphatic<6) ! and
semiaromatic(7,9) ~ Schiff base ligands which also form stable low-spin
Fe(II) complexes of the general formula (Fe(LL)aJ2+,
R
(q> < B-R
The ability of the above ligands to form very stable complexes with
Fe and other metals is associated with their visible spectra
similarities which is the result of the electron delocalisation(9)
within the five-membered che1ating ring Q.
- 2 -
c-c ! \.
J J
~/ Fe
Electron delocalisation within the Fe<II) diimine chromophore may be
more clearly described in terms of n-back bonding of electrons from
filled metal d-orbitals into vacant n' ligand orbitals 'L <synergic
effect). This back bonding (t2g to nf) results in stronger metal-ligand
bond than the sum of isolated ligand to metal ~-binding and metal to
ligand n-bonding effects.
@~d x* orbitals
~ ~ bond Fe 2+ ~ N
@ ~<10n .. palr)~ 'L
~ t t2g orbital
Iron complexes with above and related ligands exhibit intense
charge-transfer absorptions in the visible spectrum. The frequencies of
maximum absorption have been found to vary extremely little with solvent
composition. However for complexes containing both two or one of
these ligands and two or four cyanide ligands, the frequencies of
maximum absorption of the charge-transfer bands have been found to vary
considerably with the nature of the solvent(10-12>. This solvatochromic
behaviour for such mixed-ligand Fe(II) and Fe(III) complexes is
described in Chapter 7.
- 3 -
General structure of Fe(II) diimine cation, [Fe(LL)3]2+, is
depicted in a below. It resembles a three bladed propeller
structure('3.1.', where the blades are the planar diimine ligands. The
six nitrogen atoms form a distorted octahedron(lS' around the iron atom
(Chapter 3 deals with the structural properties of Fe(II) complexes).
Several reactions of Fe(II) diimlne complexes, which proceed via a
known mechanism are studied in aqueous binary mixtures. The discussion
involves a link between, on the one hand, thermodynamic and kinetic
data, and the other hand, interaction at a molecular level between
solute and solvent molecules.
1.2 PROPERTIES OF BINARY AQUEOUS MIXTURES
Insights into component interactions in binary aqueous mixtures can
be obtained by considering the thermodynamic excess functions XE. These
parameters express the extent to which the properties of the given
mixture differ from those of the corresponding ideal mixtures at the
same temperature, pressure and mole fraction. A mixture of two liquid
non electrolytes can be defined as ideal if the chemical potential of
each component obeys the following relationshlp:-
_ " -
..... <1.1)
where pOi = the chemical potential of pure i
and Xi = the male fraction of component i
For a real solution the chemical potential is given by:-
..... (1. 2)
where a;= activity of i = xif i
and fj = activity coefficient such that fi ~ 1 as Xi ~ 0
The excess thermodynamic functions of mixing for binary aqueous
mixtures are calculated as the difference between the real and the ideal
mixtures. Thus the Gibbs free energy of mixing n, moles of component 1
and n2 males of component 2 is given by:-
..... <1.3)
Since GE = 6Gm (real) - 6Gm <ideal) ..... <1.5)
It leads to ..... (1.6)
This can be positive or negative depending an the signs of the
activity coefficients. Negative values of GE indicate favourable
interactions between the components of the mixture. Nan-aqueous
component/water interactions are stranger than water/water interactions.
Positive values of GE indicate that intercomponent interactions are
weaker than those between two cosolvent molecules or two water
molecules.
For an ideal mixture the excess molar enthalpies, HE; volumes, VEj
and heat capacities, CpEj of mixing are zero. Therefore these
- 5 -
quantities directly measure deviation from ideal. The excess molar
entropies of mixing, SE, are calculated from:-
..... (1.7)
On the basis of these properties mixed aqueous solvents can be
divided into two main groups(lSl,
(1) Typically Aqueous TA
(2) Typically Non Aqueous TNA
The typically aqueous mixtures are characterised by having positive
GE values, where the sign and magnitude of GE are determined by entropy
term rather than enthalpy term, i.e.
ITSEI)IHEI
Examples of cosolvents forming TA mixtures with water include
monohydric alcohols and acetone. The plots of excess functions versus
mole fraction for some of these cosolvents with water are shown in
Figure 1.1. The striking feature is the HE curve which indicates how
the interactions between components in these mixtures change as the mole
fraction is varied.
The dependence of XE on mole fraction can be analysed to obtain the
corresponding partial molar quantities (Xl - X,.) and (X2 - X2-).
Figure 1.2 shows the variation of the partial molar volumes for some TA
systems, and this diagram identifies three types of behaviour:
(i) negative slopes at low mole fractions, this implies "structure
making". Water/water interactions are enhanced producing a clathrate
type structure around the hydrophobic group. An exothermic HE can be
identified with this behaviour.
- 6 -
GE
10001 + cr 1 /\ J.1000 ~
GE
500 500
~ ~
I I ~ ~
i 0.4 i ...., 0 0 o ...., ....... .......
-..J I III III >cI >cI
-500 -500
-1000 -1000
-1500 (a) (c) -1500 (b)
FIGURE 1.1
Kolar excess thermcdynaDdc functions for TA mixtures, at 298.2 K. (a) ethanol/water, (b) t-butanol/water, (c) acetone/water (from ref. 17,18)
~
I ...-4
i 1:1 ()
...... ...... • (14 I>
(14 I> ~
o
-4
-8
-12
0.2 0.4
(a) = ethanol (b) = acetone (c) = n-propanol (d) = t-butanol
FIGURE 1. 2
Dependence of relative partial molar volumes, (V2 - V2-) for cosolvents in some TA Dixtures as a function of cosolvent DOle fraction at 298.2 K (from ref. 17)
- 8 -
(ii) zero slope with (V 1 - Vl.) negative, no structural effects. The
point is reached where the proportion of water in the mixture is
insufficient to maintain the clathrate structure.
(iil) positive slope, "structure breaking" as X2 is increased the
extended structure is broken down. The tendency for phase separation is
most marked at these mole fractions.
Many properties of TA mixtures show extremes at the points at which
they occur on the mole fraction scale. These can act as "signposts" in
the analysis of kinetic data (as will be seen in Chapter 5).
The typically non aqueous mixtures are characterised by IHEI)ITSEI,
where the sign and magnitude of GE are determined by enthalpy term.
Thus GE can be positive, TNAP mixtures, as in the case of aqueous
acetonitrile mixture, or negative TNAN as in the case of aqueous
hydrogen peroxide and aqueous dimethylsulphoxide mixtures.
1.3 ANALYSIS OF MEDIUM EFFECT OR REACTIVITY
Consider a single step reaction (1.8) where the reactants are A and
B and (AB)~ is a transition state.
A + B (AB)* ~ products .. ... <1.8)
At fixed temperature and pressure transition state theory(S) provides
the link between the rate constant (which describes the irreversible
process) and 6G* (which can be treated using the principle of reversible
thermodynamics) in the manner shown in equation 1.9.
k .. :: (kT/h)exp(-6G*/RT) .. ... <1.9)
where k .. :: rate constant
k :: Bol tzman constant
h :: Planck's constant
- 9 -
If the solvent is a binary mixture where the mole fraction of one
component is X2, then in this solvent 6G*(X2) is given by equation 1.10.
6G*(X2) = ~. - [~e(A) + ~e(B)] ..... (1.10)
Here, ~e(A) and ~e(B) are the standard chemical potentials of the
reactants A and B (make up the sum of the initial state), and ~* is the
chemical potential of the transition state. Thus the 6G*(X2) is given
by the difference between the chemical potential of the transition and
initial states in their respective solution standard states, i.e. if the
rate constant is expressed in dm3 mol-1s-1, the relevant standard state
is the hypothetical solution where the concentration is 1.0 mol dm-3 and
the corresponding activity coefficient is unity. The standard state is
symbolised by e Because the transition state is a composite quantity
and cannot be measured directly it can be calculated from equation 1.10.
The Gibbs function is however just one of a group of thermodynamic
variables; therefore chemical potential in equation 1.10 may be replaced
by H (enthalpy>, S (entropy>, V (volume), Cp (heat capacities) .... ,
these being derived from the appropriate dependence of the rate constant
on temperature and pressure<19'. Though the major part of this thesis
only concerns itself with G, the analysis has been done from volume data
(see Capter 5).
To avoid any confusion which might arise on using the symbol 6 to
indicate the changes resulting from both chemical reactions and medium
effects, it is convenient to use a solvent operator Om for the
latter c2o >. Thus, if the Gibbs Free Energy of activation of a certain
reaction in a reference medium (1) is 6G*(1) and in a different medium
it is 6G*(2) then
..... (1. 11>
- 10 -
Linking equation 1.11 with equation 1.8 gives,
oMAG* = -RTln[k(2)/kc1 )1 ..... (1. 12)
where kl and k2 are the rate constants in medium (1) and (2)
respectively. Throughout this thesis the reference solvent has been
water.
From equation 1.12 therefore, equation 1.10 may be rewritten thus:-
oMAG* = Om~* - [oM~e(A) + om~e(B)] ..... (1.13)
If the rate constant increases on going from the pure solvent to a
mixture then the term OmAG* will be negative, ie the activation barrier
is reduced. Conversely if the rate constant decreases there is a
positive effect on OmAG*. Similarly if solute (initial or transition
state) is stabilised then Om~e is negative but if it is destabilised
then Om~e is positive. The transfer chemical potentials of solutes in
this thesis are derived from solubility data.
1.3.1 Initial State Trends and Solubility Data
Consider a saturated solution of a compound which is in equilibrium
with a solid material. The chemical potential of the solute and solid
must be equal:
~(solid compound) = p(solid compound in solution) .... (1.14)
For a solid in solution
~ = ~e + RTln(mf) ..... (1.15)
~e = the chemical potential of the solute in its solution standard
state
f = the activity coefficient such that .f ~ 1 as m ~ 0
If the solubility in solution one is S(1) mol dm-3 and solution
two is S(2) mol dm-3 then
- 11 -
..... <1.16)
The usually reasonable assumption is made, for dilute solutions
that the ratio of activity coefficients, f(2)/f(1), is unity. Thus
equation 1.16 gives the required quantity solely in terms of
experimental solubility data.
The above applies to uncharged solutes. For ionic compounds the
same analysis applies to obtain transfer chemical potential for the salt
except that S is replaced by solubility product, K.p •
The absorbance of a saturated solution of a particular species is
generally directly proportional to its concentration and hence
solubility.
Absorbance = c x E x 1 ..... <1.17)
where c = concentration, E = extinction coefficient and 1 = path length.
More generally for a salt A which forms n ions on dissociation,
equation 1.16 can be expressed in terms of absorbance, i.e.
. .... (1. 18)
where n = 1,2,3 etc for neutral, 1:1, 1:2 or 2:1 salts respectively.
However, for the analysis of kinetic data the required parameter is
the contribution of the individual ions. Hence single ion transfer
values must be estimated, i.e. in the case of a 1:1 salt, AB .
. . . . . (1. 19)
There is no single experiment by which the absolute free energy of
transfer of a single ionic species from one solvent to another may be
determined, therefore an extrathermodynamic assumption is applied.
Several methods are known(21.22>. The most popular one in recent years
1s the assumption that certain ions are only lightly solvated for an
- 12 -
anion and cation of similar size, charge and exterior and thus have
equal transfer parameters, i.e.
. .... <1.20)
In this equality assumption large organic ions of the type R.N· and
BR.- are used as reference ions. The most commonly used ions are PhAB-,
Ph4As· and Ph4P· as reported recently by Abraham(23), where
..... (1.21)
There are a number of variants to this general approach, earlier
studies using the tri-isoamyl-n-butylammonium cation. The analysis
built round the phosphonium boronate salt [the tetraphenylphosphonium
tetraphenylboronate (TPTB) assumption] has been used by Popovych(2.) while
Popovych(2S) and Tissier(26) have used the arsonium boronate salt (the
TATB assumption). In Chapter 4 the TATB assumption is used for
derivation of single ion transfer chemical potential in aqueous
i-propanol and aqueous t-butanol solvent mixtures.
1.3.2 Initial state-transition state solvation
As mentioned, transition state theory provides a route for analysis
of solvent effects on a rate of reaction. Some of the complexities
which can arise in solvent effects on 6G* are shown in Figure 1.3 where
a summary is made as to what might happen to 6G* for a hypothetical
reaction, initial state ~ transition state. The central vertical line,
in Figure 1.3(a), represents the activation process (in terms of 6G*)
for a reaction in the reference solvent. The diagram shows how on going
to a second solvent an increase in rate constant (i.e. a decrease in 6G*
at fixed temperature and pressure) can result from either (i) a
destabilisation of both states with initial state being stabilised to a
- 13 -
~
.::--
(a) (b)
TS TS
~G· ~G·
IS IS s _ ,
(i) new solvent. reference .(ii) new solvent (i) new solvent. reference .(ii) new solvent solvent solvent
FIGURE 1. 3
Representation of the effect of changing the solvent on the activation Gibbs function, ~G·, and the initial and transition states, (from ref. 2).
larger extent, or (ii) a stabilisation of both states with the
transition state being stabilised to a larger extent. Two cases, in
Figure 1.3(b) are shown where on changing the solvent, initial states
and transition states are affected differently, leading to either (1) an
increase, or (ii) decrease in ~G*. These examples only indicate trends
which may occur and not all the possibilities are covered. These
solvent effects give indications of changes in solvation of the initial
state and transition state individually on transfer between media. The
importance of solvation changes in kinetics of reaction can be
dramatically demonstrated by trends(27.2e> in activation volumes, ~V*.
Activation volumes, derived from pressure dependence on rate constant
give an indication of changes in solvation on going from initial state
to transition state. Therefore the two approaches, solvent effect and
pressure effect on the rate of reaction, provide complementary
information on the role of solvation.
- 15 -
REFERENCES
1. M. J. Blandamer and J. Burgess, Pure Appl. Chern., 51(1979)2087.
2. M. J. Blandamer and J. Burgess, Coord. Chern. Rev., 31(1980)93.
3. M. J. Blandamer and J. Burgess, Pure Appl. Chern., 54(1982)2285 55(1983)55.
4. M. 1. Blandamer, J. Burgess and J. B. F. N. Engert, Chern. Soc. Rev., 14(1985)237.
5. S. Glasstone, K. J. Laidler and H. Eyring, 'The Theory of the Rate Processes', McGraw-Hill, New York, 1941.
P. J. Robinson, J. Chern. Ed., 55(1978)509.
6. P. Krumholz, J. Am. Chem. Soc., (1955)137.
7. D. H. Busch and J. C. Bailar, J. Am. Chern. Soc., 78(1956)1137.
8. P. Krumholz, Inorg. Chern., 4(1965)609.
9. P. Krumholz, O. A. Serra and M. A. De Paoli, Inorg. Chim. Acta, 15(1975)25.
10. J. Burgess, Spectrochim. Acta, A26(1970)1369, 1957.
11. H. Kobayashi, B. V. Agawala and Y. Kaizu, Bull. Chem. Soc. Jpn. , 48
12. J. Burgess, J. G. Chambers and R. 1. Hains, Transition Met. Chem. , 6 (1981) 145.
13. D. L. Kepert, Progress in Inorganic Chemistry, Volume 23 <1978 )
14. L. Jonansson, M. Molund and A. Oskarsson, Inorg. Chim. Acta, 31(1978)117.
M. E. G. Posse, M. A. Juri, H. A. Negri, P. J. Aymonino, O. E. Piro and E. E. Castelano, Inorg. Chem. 23(1984)948.
15. K. R. Dymock and G. J. Palenik, Inorg. Chern., 14(1975)1220. P. Comba, A. M. Sargeson, L. M. Engelhardt, 1. M. Harrowfield,
A. H. White, E. Horn and M. R. Snow, Inorg. Chern., 24(1985)2327.
16. M. J. Blandamer, Advances in Phys. Drg. Chern. , Academic Press, 14(1977)203
17. F. Franks and D. J. G. Ives, Quarterly Reviews, XX (1966) 1
18. M. A. Vi llamanan and H. C. Van-Hess, J. Chern. Eng. Data, 29(1984)429
19. S. J. Dickson and 1. B. Hyne, Can. 1. Chem. , 49 (1971) 2394.
-16 -
20. J. E. Leffer and E. Grunwald, "Rates and Equilibria of Organic Reactions", Wiley, New York, 1963
21. R. Alexander and A. Parker, J. Am. Chern. Soc., 89(1967)5539. M. H. Abraham, J. Chern. Soc., Faraday Trans. I, 69(1973)1375.
22. C. F. Wells, J. Chern. Soc., Faraday Trans. I, 71(1975)1868.
23. M. H. Abraham, T. Hill, H. C. Ling, R. A. Schulz and R. A. C. Watt, J. Chern. Soc., Faraday Trans. I, 80(1984)489
24. P. J. LaBrocca, R. Phillips, S. S. Gobdberg and O. Popovych, J. Chern. Eng. Data, 24(1979)215.
25. O. Popovych and A. J. Dill, Analyt. Chern., 41(1969)456. A. Berne, B. Wajsbrot, P. D. Klahr and O. Popovych, J. Chern. Eng.
Data, 28(1983)316. 26. J. Jillard and C. Tissier, Electrochim. Acta, 27(1982)123
27. G. A. Lawrance, D. R. Stranks and S. Suvachittanont, Inorg. Chern. , 18 (1979)82
F. :M.. M.ikhai 1, P Askalani, J. Burgess and R. Sherry, Transition M.et. Chern. , 6 (1981>51
28. J. Burgess and C. D. Hubbard, J. Chern. Soc. , Chern. Commun. , (1983) 1482
J. Burgess and C. D. Hubbard, J. Am. Chern. Soc. , 106(1984)1717.
-17 -
CHAPTER
2
Experimental
2.1 IBTRODUCTIOI
Experimental details and the equipment required to collect the
necessary data for work in this thesis are briefly described in this
Chapter. Absorption measurements and kinetic data collection were made
using a double beam SP 800, SP 8-100 and SP 1800 spectrophotometers,
made by Pye-Unicam, and a single beam HP8451A Diode Array
Spectrophotometer, made by Hewlett-Packard. The block silica UV/visible
cells <of path length 10mm) with reagent volume of approximately 3 cm3
were used in all instruments. All experimental data were collected at
298.2 K.
2.2 SOLUBILITY MEASUREKENTS
Saturated solutions were obtained by vigorously shaking sealed
vessels containing the solvent and a large excess of the solute. The
vessels were then thermostated by placing them in a constant temperature
bath maintained at 298.2K, where they were left for a period ranging
from 2 - 48 hours. Periodical shaking of the solutions was maintained
to ensure the solvent was completely saturated. Aliquots of the
saturated solution were centrifuged to separate undissolved solids and
then thermostated again to maintain a temperature of 298.2K. The
saturated solution was carefully decanted from the solid and diluted, as
necessary, with a known volume of the pure solvent until it was of
sufficient concentration to be monitored spectrophotometrically.
To calculate the solubility of the solid in a particular solvent
mixture the absorbance and Xmax were followed using the SP 800 or SP8-
100 spectrophotometers. The direct concentration of metal was
determined using flame photometer or an atomic absorption
- 18 -
spectrophotometer. For most of the compounds which absorb in the
visible region the extinction coefficients were determined by measuring
the absorbance of a known concentration (in moles per litre) of the
solid in solvent mixture. The extinction coefficients were calculated
using the Beer-Lambert equation, equation 1.17. It is important to note
that all work in this thesis is done on the molar scale and hence when
calculating solubilities the concentration is expressed in moles per
litre. For each solvent composition two saturated samples were made
and each monitored at least three times to check for consistency.
2.3 FIRST ORDER RATE COISTANT
Consider a reaction that proceeds to completion in which the
concentration of only one reactant, A, changes appreciably during the
reaction. Such a case may arise because:-
(1) only one reactant, A, is involved.
(2) all other reactants are in a much larger concentration than A.
(3) concentration of all the reactants is kept constant by buffering.
In this thesis only case (2) is relevant where one other reactant is
involved, ie
A + B ~ C, where [Bl » [Al .••.• (2.1)
The course of the reaction, equation 2.1, may be monitored by measuring
the changes in the concentration of A with respect to time, the change
which may often be expressed by
d[Al/dt = k[Al- ..... (2.2)
A first order dependence occurs when a = 1. The rate of loss of the
reactant A decreases as the concentration decreases. The differential
form of equation 2.2 leads to several equivalent integrated expressions
- 19 -
and
[Al~ = [Al o exp (-kt)
In([Alo/[AJ~) = kt
-dln[Al~/dt = k
Where [Al o = initial concentration of A
[Al~ = concentration of A at time t
..... (2.3)
· .... (2.4)
..... (2.4a)
k = rate constant under the given conditions of temperature(T),
pressure<P) and ionic strength(Xi).
A quantity characteristic of first order reactions 1s t~, the half
life of the reaction which 1s the value of t when
[Al~ = [Alo/2 or
[Al ~+~.. = [ Al ~/2 · .... (2.5)
t.. = In2/k · .... (2.6)
From estimating the half life one can thus calculate the approximate
value for the first order rate constant. The half life of a first order
reaction remains constant aver the complete reaction. The observed
first order rate constant, k(ob.), in many of the systems varied as the
concentration of reactants ather than A was varied. A cammon pattern
which emerged was dependence where:-
kCOb.> = k, + k2[BJ
where k, = first order B-independent path
k2 = second order path A + B ~ C
· .... (2.7)
The second order rate constant may thus be measured from variation
of kobe with concentration of B. This type of pattern is illustrated in
Chapters 5, 6 and 8 where it was found that k, is effectively zero. In
all systems studied here, kinetics of reaction were monitored by
observing changes in concentration of either the reactants or the
- 20 -
products. Concentration changes of any species were measured
spectrophotometrically by measuring the absorbance arising from a single
chemical species, A, in the solution. For most of the reactions the
absorbances were followed over 2~ lives of the course of the reaction.
2.4 SP 800 SPECTROPHOTOMETER
This instrument was used to produce scan spectra in the range
from 200 - 700nm. The cell compartment could hold four samples (and
four reference) cells in a cell block which was kept at constant
temperature by circulation of water from a constant temperature bath.
The instrument has a SP 825 timer attachment fitted which allows spectra
to be scanned repeatedly on the same chart paper at any preselected time
interval up to 15 minutes. This is a very useful facility as it enables
preliminary investigations into particular kinetic systems to be carried
out.
2.5 SP 8-100 SPECTROPHOTOMETER
This instrument is capable of measuring absorbance over the range
of 200 - 800nm and gives an accurate digital reading from 0.000 to 2.000
+0.001 at any particular wavelength. The four-cell block holder is
thermostated by circulating water from a built-in temperature controller
which may be set to give any temperature within the range 273.0 K to
381.0 K to.l K. The temperature is monitored by a platinum resistance
thermometer placed in the cell holder and gives a digital reading of the
temperature. Apart from preliminary kinetic investigations, the
kinetics of fast reactions were monitored on SP 8-100 Spectrophotometer
using the 'SFA-ll Fast Kinetics Accessory' unit which is depicted in
- 21 -
I
N N
STOP WITH 0 MICI!OSWIT( II
MANUAL DRIVE
DRIVE SYRINGES
\ SOLUTION RESERVOIRS
THERMOSTAT flUID IN
flEXIBLE UMBIliCAL (600mm long)
THERMOSTAT FLUID OUT
J - PORT VALVES
SPEC TROPHOTOMET E R THERMOSTAT TEO Cell HOLDER
~ LIGHT PATH!
Cell
-I I I I
-£3- t--, I ,
L- __ .-J
FIGURE 2.1
2mm PATH , /
lOmm PATH // /'
DETAIL OF CELL
The 'SFA-ll Fast Kinetics Accessory' unit and details of the cell
! l /
-1 Smm
/}smm /
Figure 2.1. This unit, with its specific cell, is easily fitted onto
the spectrophotometer enabling reactions with a half-life down to one
second to be monitored. The limiting factor is the response of the
plotter on the spectrophotometer.
2.6 SP 1800 AND HP 8451A DIODE ARRAY SPECTROPHOTOMETERS
Both the SP 1800 and HP8451A Diode Array Spectrophotometers were
used to monitor the dependance of absorbance on time at a single
wavelength. Central to the operations of these instruments were
microcomputers which recorded, displayed and stored absorbance as a
function of time.
The SP 1800 Spectrophotometer, which is similar to SP 8-100 in that
it possesses many similar features, has been equipped with a mini
computer(HP 9825A) with an attached graph plotter (HP 7245A). The
development work of interfacing the mini-computer with the
spectrophotometer was carried out by Dr. M. J. Blandamer and fully
described in literature<1,2).
The HP8451A instrument is operated by two microcomputers; the Z.80
which controls the internal hardware of the instrument and the HP85A
which deals with data and acts as an interface between user and
spectrophotometer. The wavelength range of this instrument is from 190
to 820nm. This instrument has one cell holder which is thermostated. A
scan of spectrum over the whole range required only a few seconds. The
detailed operational procedure for this instrument has already been
described elswhere<3).
- 23 -
2.7 HIGH PRESSURE KINETIC APPARATUS
Rate constants at high pressures were determined by the use of
apparatus which is depicted in outline in Figure 2.2. The reaction
mixture, approximately 150 cm3, was made up, from pre-thermostated
solutions, immediately before introduction into the pressure vessel.
Each run was monitored by expelling aliquots from this pressure vessel,
itself thermostated, and promptly reading off their absorbances on the
SP 8-100 spectrophotometer. The pump restored the pressure to its pre
set value within a second of taking a sample. The limiting factor on
rates which can be measured with this apparatus is the time needed to
fill the pressure vessel, re-assemble the container, and re-thermostat
the reaction mixture. These operations take a few minutes so the
apparatus is only suitable for reactions whose half lives are many
minutes, or hours.
2.8 ATOHIC ABSORPTION SPECTOPHOTOMETRY
This technique was employed in the determination of solubilities of
Fe, Co and Rb salts in pure water and in binary aqueous solvent mixtures.
The instrument used was 'Perkin-Elmer 360' Atomic Absorption
Spectrophotometer, using air-acetylene flame. The absorbance of the
radiation was monitored at a given wavelength for each element employing
a hollow cathode lamp source. Calibration of the instrument was carried
out using the appropriate freshly made standard. Absorbance for all the
metal ions used were linear up to concentrations of approximately 5~g/ml
in aqueous solution; only a small increase in absorbance was observed
when organic cosolvents were used.
- 24 -
I\.) U1
FIGURB 2.2
Schem3tic Diagram of High Pressure Apparatus
HZ
TAP B GAUGE
WATER -..- :::=:::::t \' I \ J \ ..... ,.... ew.
WATER ~ BATH
/" o \ -
4U-+- CEll.
BOMB"
1-
TEfLON PLUNGER
REFEREN~
1. K. J. Blandamer, Computer Control
2. P. P. Duce, Ph.D. Thesis, University of Leicester, (1984)
3. B. Clark, Ph.D. Thesis, University of Leicester, (1985)
- 26 -
CHAPTER
3
Crystal Structure of Fe{II> Complexes
3.1 INTRODUCTION
An extensive study of thermodynamics of solvation of low spin
ironCII) diimine complexes in methanol-water mixtures(l) suggests that
ligand bulk and hydrophobicity playa key role in determining
preferential solvation patterns. Thus, for example, the relatively
small complex [Fe(hxsbh)]2+, where hxsbh = 1, has a rather small
transfer chemical potential from water into methanol-water mixtures(2).
The much bulkier complex [Fe (bsb-Me2)3] 2+, where bsb-Me2 = ~, which has
a totally hydrophobic periphery, is greatly stabilised on transfer from
water into methanol-water mixtures, indicating strong selective
solvation by methanol(3). The parent diimine complex, [FeCphen)3]2+,
and the cage complexes of the type ~ show an intermediate
behaviour< "'.6).
Of particular interest, for this study, are aliphatic diimine
complexes [Fe(LL)3]2~ where LL = i, and cage complexes of ironClI).
Complexes from both series are suitable substrates for solubility and
reaction kinetic studies. Further, in both series a varied ligand
hydrophobicity can be achieved by introducing an appropriate organic
substituent during the preparation. The aliphatic iron(II) diimine
complexes can be prepared through condensation of a-diketones (RCOCOR)
with methylamine in the presence of Fe2+, while a series of cage
- 27 -
complexes, (Fe(R2cage)]2+, can be prepared through condensation of the
dihydrazones of a-diketones with formaldehyde, or other aldehydes
(RCHO), in the presence of Fe2+. However, in the case of formation of
cage complexes, it is feasible that the products of such reactions might
be uncaged complexes containing three bidentate ligands(S). Before
assessing preferential solvation characteristics for the above mentioned
and other complexes, as is the case in Chapters that follow, in this
Chapter preparation of several complexes and their crystals for X-ray
diffraction crystal structure determination is reported. The X-ray
diffraction results are reported and structures described for iron(II)
complexes of gmi, bmi and cxcage.
R R' " / C-C
I \-){e-I I-Me
3.2 EXPERIMENTAL
Preparation of [Fe(gmi)3]2+ and (Fe(bmi)3]2+ complexes
Iron(II) tris-gmi and bmi complexes were prepared by the published
methods(6.7) through condensation of glyoxal and 2,3-butanedione with
methylamine for gmi and bmi respectively in 50% methanol solution. To
the resulting product one third equivalent of FeC12 solution was added,
resulting in formation of a purple red solution. The iron(II) tris-gmi
- 28 -
complex was precipitated as its tetrafluoroborate salt while that of bmi
was precipitated as the bis-perchlorate salt. Both complex salts were
recrystallised from 50% aqueous ethanol. The crystalsfor X-ray
diffraction were prepared by dissolving the complex salts in aqueous
acetone solution, then gradual evaporation of the solvent was maintained
whereby octahedral crystals for gmi were obtained, while those for bmi
were hexagonal. Both of the above complexes were also prepared as PF6-
salts for nmr purposes as was [Fe(cmi)3]2+ which was prepared from
condensation of cyclohexane-1,2-dione with methylamine followed by
addition of one third equivalent of FeC12 solution(7). Final products
were characterised by their ~M.X and (E), which are 554(8600),
564(10600) and 582(12467) for gmi, bmi and cmi respectively.
Preparation of [Fe(cxcage)]2+ complex
Cyclohexane-l,2-dione dihydrazone was prepared by condensation of
cyclohexane-l,2-dione with hydrazine by the standard method(e,. The
preparation of [Fe(cxcage)]2+ was attempted, from Fe(BFA )2 by the method
published for [Fe(Me2cage)]2+(9). In contrast to [Fe (Me2cage)] 2+,
[Fe(cxcage)]2+ could not be easily isolated from the reaction products
as its BF4- salt. Crystallisation after addition of KPF6 to the
reaction mixture produced a mixture of at least two constituents. A
good crystal was selected from this mixture for the X-ray diffraction
crystal structure determination. Subsequent fractional crystallisation
gave pure crystals of this product, which is the desired [Fe(cxcage)]2+
complex (E622 = 5074 M-1cm- 1). The more soluble product, similar but
not identical in colour and in E and ~(M.X)' could not be obtained pure
(£674 = 9587 M-l cm- t • This is probably an uncaged complex though the
- 29 -
simplest ligand formula does not seem altogether likely in view of the
very different reactivity of this complex. The perchlorate salt of this
presumed uncaged complex was prepared by direct reaction between iron
CII) chloride, formaldehyde and cyclohexane-l,2,-dione-dihydrazone in
acetonitrile with subsequent precipitation by sodium perchlorate.
Besides the difference in E and Xmax of the above two products in
acetonitrile, the cage complex was found to be stable in H20 while the
presumed uncaged complex was shown to be reacting with H20 when
dissolved, i.e. a change in Xmax from 574 nm to 642 nm. Further it was
realised that [FeCcxcage)]2+ in presence of hydroxide in solution
undergoes dissociation resulting in a new complex whose Xmax is similar
to that of the uncaged species.
The lH-nmr spectra of the gmi, bmi and cxcage ironCll) complexes
were run in d3-acetonitrile on 90KHz instrument. For better resolution
the proton nmr spectra of the cxcage complex was also run on the 300HHz
instrument. The proton nmr spectrum was run for the uncaged complex but
no useful results were obtained. The crystals of the [FeCcxcage)] (BF.)2
complex used in X-ray diffraction were black parallel piped blocks.
3.3 RESULTS
Tables with X-ray diffraction data for [Fe(cxcage)]2+, [Fe(gmi)3]2+
and [Fe(bmi)3]2+ complexes(10.11) are set out in Appendix 1 together
with 'H-nmr spectra for each complex. Selected bond distances and
angles for Fe environment are presented in Table 3.1 for all three
complexes. Due to the symmetrical nature of gmi and bmi ligands only one
Fe-N bond distance is found and reported. Further tabulated data of
bond distances and angles can be found throughout this Chapter. The
- 30 -
TABLE 3.1
The Iron Environment in (a) cxcage, (b) gDd and bud Fe(II) complexes
The second column in the matrix is the Fe-I distance in A. other
entries are the angles subtended at the iron atom by the relevant
ligand atoms at the head of the row and colum.
(a)
r
1(1) 1. 914 (6)
1(7) 1.896(6)
1(3) 1. 921 (5)
1(9) 1.930(5)
1(5) 1. 897 (6)
1(11) 1. 910 (6)
(b)
Complex
[ Fe (gmi) 3] ~l+
[Fe (bmi )3] 2+
1(7) 1(3)
78.3(2) 86.7(2)
117.1(3)
11-1 I A
1. 952 (2)
1. 956 (2)
- 31 -
1(9)
149.6(3)
86.2(3)
77.2(2)
1(5)
86.5(2)
150.7(2)
86.4(2)
117.6(3)
80.0(1)
79.5(2)
1(11)
117.0(2)
85.9(3)
150.5(3)
86.9(3)
79.1(2)
(b)
o
FIGURE 3.1
The structure of [Fe(cxcage)]2+ cation (a) and a view along the threefold axis showing the twist angle of 23- (b)
- 32 -
w w
~-G~ n
n~' Y ~/J (.G ).....
c''---{ ~ >-:\ '--~~J P Fe / ,,-I
~~) '---\c c!''--'~
(a)
H:s:C
H H
\ 1.423 / (\1> C C'-
1.459 Y ~N CH:::3
~ / ~2
Fe (b)
FIGURE 3.2
The structure of [Fe(gDd)3]2+ cation (a) and relevant bond lengths (A) within the complex (b).
Q ~~') '(~ G~ c~~) r
~Jc'l . 0~) n
o _I , I . (I J (
~ c-l-'~ a""''--{ 0 . -"-/A:J P Fe)-O 0
(~/i~ f) .j- ~ A~ ---, ~):J ~
-x--'i ~ ) (( ~, eA-~~ 'Y'/
(a)
H=\ 1. 477 lH= 1.474 yC ,
H~C N, /N---CH~ Fe
(b)
FIGURE 3.3
The structure of [Fe(bDd)3]2+ cation (a) and relevl bond lengths (A) wi thin the complex (b)
TABLE 3.2
Summary of lH-nmr data for Fe(II) complexes in d3-acetonitrile
Complex proton nmr data
[Fe (c][cage)] 2+ . II , II ppm 8 6 4 2 o
[Fe (Jfe2 cage)] 2+ I II , II I I I
ppm 8 6 4 2 0
[ Fe (cDl h.] 2+ • , I III1 .III ppm 8 6 4 2 0
[Fe(bDl)31 2 + , I , II I ppm 8 6 4 2 0
[Fe(gmi)a1 2 +
ppm 8 6 4 2 o
For further proton nmr of these complexes see Appendix 1
- 35 -
geometry of the complex ions is shown in Figure 3. l(a and b), Figure 3.2
and Figure 3.3.
3.4 DISCUSSION
3.4.1 Structure
The X-ray diffraction results prove the encapsulated nature of the
(Fe(cxcage)]2+ in the solid state, as shown in Figures 3. l(a and b).
The three cyclohexane rings are arranged in a propeller type manner
around the Fe, while the further two rings, C3N3 formed as a result of
capping by H2CO on cyclohexane-l,2-dione-dihydrazone, are found on
three-fold axes. Similar ligand arrangements are found in structures of
(Fe(gmi)3]2+ and [Fe(bmi)3]2+ complexes where the ligands are of
propeller type, two groups of three methyls on imine nitrogens are found
along the three fold axis as shown in Figures 3.2 and 3.3.
The evidence, that these structures are maintained in solution, is
provided by lH-nmr spectra, results of which are tabulated in Table 3.2
where they are compared with those compounds which are closely related.
The lH-nmr spectrum for (Fe(cxcage)]2+ complex is consistent with its
structure. The AB pattern of the resonances from the protons on the
capping, C3 N3, rings is analogous to that reported Cg) for the
(Fe(Me2cage)]2+ cation, while the cyclohexyl protons give rise to
unresolved multiplets at 1.75 and 2.95 ppm, corresponding to similar
signals at 1.85 and 2.92 ppm for the [Fe(cmi)3]2+ complex. The
symetrical nature of bmi and gmi ligands is also indicated by their
lH-nmr spectra. The protons from two methyl groups in [Fe(bmi)3]2+ give
rise to singlets at 2.5 and at 2.79 ppm, whereas a singlet for the
- 36 -
methyl groups in [Fe(gmi)3J2+ is located at 3.08 ppm while a singlet for
the ethylene protons is located at 8.56 ppm.
3.4.2 The coordination polyhedron
The six imine nitrogen atoms around the iron define a coordination
arrangement intermediate between that of an octahedron and of a trigonal
prism(12.13', as depicted in Figure 3.4. For [Fe(gmi)3]2+ and
[Fe(bmi)3]2+ the twist angles are 52.8' and 53.0' respectively which are
closer to an octahedron where a = 60', as in Figure 3.4(c). The same
twist angle, where a is close to an octahedron is found in other tris-
diimlne complexes of Fe(II), [Fe(LL)3)2+. Thus a = 51' in [Fe(4,4-
bithiazole)31 2+(14', 53 0 (16) or 550(16) in [Fe([phen)3J2+ and 55 0 (17.19>
in [Fe(bipY)31 2+. The small deviation from an octahedron in bmi and gmi
iron complexes can be attributed to the same factors as those in iron
complexes of phen and bipy: that is that the methyl groups, on imine
nitrogen, in axial position are in a close proximity.
(a)
FIGURE 3.4
Coordinntion arrangements defined by six i1l1ne nitrogen atoms ar-ound the iron: (a) trigonal prism, (b) intermediate and (c) octahedron.
- 37 -
On the other hand the [Fe(cxcage)]2+ complex exhibits a twist
angle a of 22.7· which is nearer to a trigonal prism (a = 0·) than to an
octahedron (a = 60·). It is interesting that in a closely related
series of complexes of the encapsulating ligand di(amH)sar(19) twist
angles vary between 25· and very nearly 60· depending on crystal field
factors. In the free ligand and in the complexes of metal ions with
zero CFSE a is 25·, but for high CFSE as for t2gS Co(III) a is 55· to
58·. The ligand di(amH)sar must be less rigid and demanding than
cxcagej t2gS Co(III) can force the di(amH)sar type ligand to adopt a
geometry that permits almost octahedral stereochemistry of the metal,
but the more rigid cxcage does not allow even t2gS Fe(II) to force
octahedral geometry.
The high rigidity within the cxcage ligand is primarily due to the
Sp2 hybridised nitrogen atoms in the C3N~ rings. Bond angle data
analysis shows a consistent difference in angles around all the
nitrogens from the C3N3 rings. These nitrogen atoms which are in the
plane of the cyclohexane dihydrazone di1mine moiety prefer trigonal
prismatic geometry, Figure 3.5(a), while the t2g6 iron prefers
octahedral geometry, Figure 3.5(b). The compromise of these two factors
results in 23· in the twist angle from that of a trigonal prism.
3.4.3 The imine moiety
It is of interest to compare the N-C and C-C bond lengths in the
imine moiety in the various Fe(II) compounds. The bond lengths in
conj ugated molecu les are 'chiefly dependent on (i) bond length, <i 1)
hybridisation, (iii) formal charge distribution, where the order of
importance, not always but often is (i) ) (1i) ) (i11)(20>. In the
- 38 -
I I
c_
I
I
(a) (b)
FIGURE 3.5
Part of the cxcage ligand indicating that angle ~ = ~' in trigonal prismatic geometry (a) and ~ < ~' in octahedral geometry (b).
- 39 -
.!> 0
TABLE 3.3
Bond lengths (A). bond angles (.) and twist angles (aJ·) associated with free
and coordinated diimines I-C-C-I.
ColIplex ](-I/A I=C/A C-C/A I-](-I/- aI-
I [Fe(gm)3]2+ 1.952 1.272 1.423 80.0 52.6
[Fe (bJd )3]2+ 1.956 1.292 1. 477 79.5 52.8
[ Fe (cxcage) ] 2+ 1.896-1.930 1. 262-1. 298 1.425-1.450 77.2-79.1 23.0
[Fe(bipY)3]2+ 1. 947-1. 964 1. 340-1. 350 1.420-1.480 81.5 54.6
bipy 1.360 1.50
[Fe(phen)3]2+ 1. 960-1. 980 1. 300-1. 360 1.373-1.389 82.9 53.0
phen 1.380 1.450
Ref.
this work
this work
this work
17, 18
23
15
24
.!> --
TABLE 3.4
Variation in second-order rate constant, k2. for hydroxide attack in water and redox
potentials for four iron(II) diiDdne complexes
CoDplex [ OH-] I Dll d.Dr3 k<o~)/s-l k21 d.JIl3 Dll- 1 E/V<a)
[Fe <gm.)3] 2+ 0.02 1. 14x10-4 5.'70x10-3 0.81
[ Fe hllli ) 3] 2+ 0.02 8.8'7x10-6 4. 44x10-4 0.62
[Fe (bmi) 3] 2+ 0.05 3.69%10-6 '7. 38x10-5 0.45
[Fe (Cm!)3] 2+ 0.05 2. 46x10-6 4. 92x10-5 0.42
(a) from reference 26
diimine complexes the C-C bonds are all longer than in benzene
(1.39A)(21) but much shorter than in ethane (1.54A)(22). The C-C bonds
in these complexes have slightly more double bond character in the
complex than in the free ligand as can be seen in Table 3.3 for bipy(23)
and phen(2A). The C-N bond lengths are slightly shorter than those of
pyridine(1.37A)(2S) indicating the higher bond order.
Although Fe-N bond lengths for gmi and bmi complexes are the same
and do not differ much from Fe-N bond lengths in other complexes, it is
evident that bond orders are higher in N-C and C-C bonds for the gmi
chromophore than that for bmi (see Table 3.3 and Figures 3.2 and 3.3).
This is probably due to the inductive effect from the methyl groups on
the imine carbon in bmi. Therefore for the aliphatic ligand Fe(ll)
complexes gmi, mmi, bmi and cmi it can be assumed that the bond order in
the imine moiety decreases from gmi to cmi. Replacement of hydrogen by
methyl groups on the imine carbon decreases the bond order; i.e.
increases the ~-bond character but decreases the rr-bond character in the
Fe-N bond. This variation in the bond order is also evident from the
results of hydroxide attack, Table 3.4, on these complexes where k2,the
second order rate constant, decreases from gmi to cmi. Further evidence
for this is the redox potential(26) for the four Fe(ll) complexes which
decreases by 0.18 volt per methyl group replacing hydrogen atom at the
imine carbon.
- 42 -
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10. J. Fawcett. Chemistry Department. University of Leicester
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12. K. R. Dymock. G. J. Palenik, Inorg. Chem .• 14(1975)1220
13. D. L. Kepert, Progress in Inorganic Chemistry. Volume 23 (1978)
14. T. A. Baker. and H. A. Goodwin. Aust. J. Chem., 38(1985)851
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19. P. Comba, A. M. Sargeson. L. M. Engelhardt. J. M. Harrowfield, A. H. White, E. Horn and M. R. Show. Inorg. Chem. , 24 (1985) 2325
20. C. A. Coulson, J. Phys. Chern. , 56(952)311
21. R. T. Morrison and R. N. Boyd, Organic Chemistry. All yn- Bacon, 1976
22. E. S. Gould. Mechanisms and Structure in Organic Chemistry. Holt-Dryden. 1960
-43-
23. O. P. Anderson, J. Chem. Soc., Dalton Trans., (1972)2597
24. s. Nishigaki, H. Yoshioka and K. Nakatsu, Acta Cryst., B34(1978)875
25. J. A. Dean, Handbook of Organic Chemistry, McGraw-Hill, 1986.
26. P. Krumholz, H. LiChum, M. A. D. Paoli and T. Rabockai, J. Electroanaly. Chem., 51(1974)465
CHAPTER
4
Solubility of Salts and Derivation of
Transfer Chemical Potential in Several
Binary Aqueous Cosolvent Systems
4. 1 INTRODUCTION
The influence of hydrophobic/hydrophilic character of the complex
on its transfer parameters from water into methanol has been documented
for [Fe(hxsbh)]2+ and [Fe(bsb-Me2)3]2+ complex cations(1.2). Further
investigation into factors which determine transfer chemical potential
and preferential solvation(3.4) initiated preparation of other iron(II)
diimine complexes(S) and complexes of other transition metals(S).
Single ion transfer chemical potentials are well established for
methanol-water mixtures, using TPTB assumption(7). However such
information is sparse(B.9) in other binary aqueous mixtures due to lack
of data necessary for single ion splitting.
The work in this Chapter deals with solubility and derivation of
transfer parameters of simple and complex salts in binary aqueous
mixtures, the results as such are used as the basis for examination of
the effect of charge, ligand structure and hydrophobic/hydrophilic
character on the solvation of inorganic ions. The importance of the
size of a complex and hydrophobicity of its exterior is provided by a
series of aliphatic [Fe(LL)3]2+ complexes, where LL = I to IV below.
R RI
'\ / I gmt R = RI= H C-C II mmi R = HI RI = 1(e
Jle-.f ~ III bmi R = RI =1Ie H-Me IV c.mi HRI =0
Single ion transfer chemical potentials in methanol-water solvent
mixtures, based on the TPTB assumption, for Fe(II) tris- gmi, romi, bmi
and cmi complexes are derived from solubilities of their moderately
soluble perchlorate salts. The results are compared not only with each
other but also with those of other bidentate, terdentate, hexadentate
- 115 -
and cage Fe(II) diimine complexes which are also derived from the
solubilities of their perchlorate salts. Further transfer chemical
potentials of Fe(II) diimine complexes are used in deriving transfer
chemical potentials for such anions as BF.-, PF6 -, SCN-, and
nitroprusside which do not form sparingly soluble simple salts.
The above work with a methanol cosolvent is extended into other
alcohols (EtOH, i-PrOH and t-BuOH) and also acetone; cosolvents which
are more hydrophobic and have larger effects on water structure(lO).
Single ion transfer chemical potentials have been derived for the whole
range in aqueous mixtures of i-PrOH and t-BuOH. These were obtained by
dissecting Om~e values of Ph4AsC10., [Co(en)3) (Ph.B)3 and
(Co(en)3) (C104)3 salts using tetraphenylarsonium/tetraphenylboronate
(TATB) assumption. From this, Om~e values for other simple and complex
ions have been obtained and are compared with corresponding values in
methanol, ethanol and acetone. The aim of this is to provide a detailed
picture for solvation of simple and complex inorganic ions which is used
in analysis of kinetic data for reactions involving such ions in binary
aqueous solvent mixtures.
4.2 EXPERIMENTAL
4.2.1 Preparation of Compounds
Most of the Fe(II) diimine complexes were prepared by introducing
some improvements to already established methods(1',12). The
preparation of tris-gmi, bmi and cmi Fe(II) complexes has already been
described in Chapter 3. For solubility purposes these complexes were
prepared as bis-perchlorate salts, as was the [Fe(mmi)3)2+ complex which
was prepared from pyruvic aldehyde and methylamine(12). Complexes were
- 't6 -
precipitated as perchlorate salts and recrystallised from the minimum
amount of aqueous ethanol. Fe(I1) tris-phen and tris-bipy complexes
were prepared by mixing, 3:1 ratio, ligand with FeCl2 in methanol
solution, while the bidentate (Fe(mpmi)3]2+ complex was derived from
condensation of methylamine with 2-acetylpyridine according to a method
by Krumholz(13). The terdentate Fe(11) complexes were prepared through
condensation of 2,6-diacetylpyridine(14.1S) with an appropriate amine
(ammonia, methylamine or aniline for [Fe(tsbh)2]2+, [Fe(tsbMe)2]2+ and
[Fe(tsbPh)2]2+ respectively). The two types of hexadentate Schiff bases
were prepared by condensation of 2,3-butanedione monoxime(16,17) or 2-
acetyl pyridine(ls> with 1,4,7,10-tetraazadecane (trien), The Fe(11)
oximes were prepared by mixing (one to one) FeC12 and ligand as in the
case of Ni(I1), while the Ni(IV) complex was prepared by oxidation of
li(I1) oxime complex by concentrated nitric acid(19.20). The
preparation of the cage complex, [Fe (Me2Bcage)], was according to the
published method(21), while that for (Fe(cxcage)]2+ has already been
reported in Chapter 3. The majority of the above complexes were
initially prepared as perchlorate salts but many of them were also
prepared as thiocyanates, tetrafluoroborates, hexafluorophosphates or
nitroprusside salts.
The cobalt(1II) complexes, (Co(en)a]3+ and (Co(NHa)6]a+, were
prepared as their chloride salts according to published methods(22.2a).
The former was also precipitated as a salt of iodide, perchlorate and
Ph4B-, while the latter was also prepared as a perchlorate salt. The
(Co(NH3 )4(COa)]+ cation was prepared as a perchlorate salt by a
published method(24), as was the Reineckate salt(2s>, K(Cr(NHa )2(SCN)4].
The tetraphenylarsonium perchlorate, Ph4AsCl04, was prepared by
- 47 -
metathesis from Ph4AsCl and NaCl04 , while KCl04 and RbCl04 were
prepared by metathesis from their respective chlorides with NaCl04 in
aqueous solution. All the complexes prepared were characterised by
their ~m.x and extinction coefficient values.
4.2.2 Solubility Measurement
Solubility measurements were carried out at 298.2 K in water and
methanol-water solvent mixtures. The sample tubes containing a range of
solvents were saturated with complex salt and left to equilibrate. The
solutions were regularly agitated in order to ensure complete
saturation. After a period of 4-5 hours for more soluble salts or
longer for less soluble salts the solubilities were, after appropriate
dilution, determined spectrophotometrically by measuring their
absorbance at appropriate ~max using SP 800 or SP 8-100
spectrophotometers for most of the salts. An atomic absorbance
spectrometer (Perkin Elmer 360) was used in solubility determination for
RbCl04 and [Co(en)3] (Ph4B)3, and flame photometer in the case of KCl04.
4.3 RESULTS AND DISCUSSION FOR AQUEOUS METHANOL
Solubility data together with derived transfer parameters for
complexes are presented in Tables 4.1-4.6. In all derivations of single
ion transfer chemical potentials calculations have been made by using
the molar (as opposed to molal) scale for the solute and volume percent
composition for the mixed solvent. Volume composition defines the
composition before mixing, therefore 40% by volume defines a solvent
prepared by mixing 40 cm3 by volume of methanol and 60 cm3 of water at
- 48 -
298K. Most Tables also contain corresponding weight percent and mole
fraction composition scale for the mixed solvent.
Table 4.1, which uses [Fe(bipY)31 (Cl04 )2 complex salt as an
example, summarises the normal procedure followed for determining
solubility and derivation of single ion transfer chemical potential. In
the tables that follow solubilities of salts will be presented in the
form of absorbance units, also included will be derived transfer
chemical potentials for the salt and anion or cation. Moderately soluble
Fe(bipY)3(CI04)2, like most Fe(II) diimine complexes, shows an increase
in solubility with an increase in methanol content of the binary aqueous
mixture (Figure 4.1a>. Maximum solubility in the methanol region (80%
vol MeOH) is also reflected in derived transfer chemical potentials for
the salt and the cation as shown in Figure 4.1b.
Transfer chemical potentials for the salts are derived using an
equation 1.18 with the assumption that, for a given salt, ratio of
activity coefficients (f(2)/f(1» is unity. This approximation is
acceptable for sparingly soluble salts whose solubility decreases with
an increase in methanol content. However for moderately soluble salts
of Fe(II) diimine complexes this assumption is less acceptable due to
ion pair formation in the methanol rich region. Some evidence that
derived single ion transfer chemical potentials are not seriously
affected by the above assumption is provided(2S) for [Fe(phen)3] (C104)2.
Here, the derived single ion transfer chemical potentials are affected,
if the activity coefficient correction is not taken into account, only
when the methanol content exceeds 60% by volume.
The obvious increase in ligand bulk (cmi > bmi > mmi > gmi) and
hydrophobic character of the periphery of the complex is shown in
- 49 -
Figure 4.2 which is in accordance with the earlier suggested
pattern(l) based on the hydrophilic [Fe(hxsbh)]2+ and the most
hydrophobic cation [Fe (bsb-Ke2) 3] 2+. The smallest complex cation
(Fe(gmi)3]2+ with a hydrophilic periphery shows little preference for
methanol in the water-rich region, but, unlike the hydrophilic
hexadentate complex, in the methanol region it shows some
destabilisation. The difference in the preferential solvation in the
high methanol region between [Fe(gmi)3]2+ and hexadendate complex is
that the former is preferentially solvated by water whereas hexadentate
which contains hydrophobic areas also shows some stabilisation,
indicating a preferential solvation by methanol. Increase in ligand
bulk, the complex cations of (Fe(pmi)3]2+ and [Fe(bmi)3]2+, leads to
preference for methanol until, with cmi ligand stabilisation of the
(Fe(cmi)3J2+ complex cation on transfer to methanol rich region is
nearly as high as that for the most hydrophobic cation [Fe (bsb-Ke2)3] 2+.
The downward turn of the plot against solvent composition of
6M~e{[Fe(cmi)3]2+} towards 100% methanol may reflect significant ion
pairing, [Fe(cmi)3]ClOA+ will be more favourably solvated by methanol
than [Fe(cmi)3]2+ and CIOA- separately.
A similar pattern of ligand bulk is seen for bsb and tsb (for
bidentate and terdentate ligands). The general trend for preferential
solvation by methanol is noted as ligand bulk and hydrophobicity
increases (with the exception of the special case where the nonpyridine
nitrogen atom bears a proton rather than Ke or Ph (see Chapter 6).
Derived transfer chemical potentials for encapsulated complexes
from Table 4.4, [Fe(cxcage)] (PF6)2 and neutral [Fe(Ke2Bcage»), are
- 50 -
...,'l
TABLE 4.1
Derivation of transfer cheDdcal potentials for the [Fe(bipY)3]2+ cation from solubility of its bis-perchlorate salt in water and methanol/water solvent Ddxtures, at 298.2 K.
Vol~ 0 20 40 60 80 100
~1. ](ethanol 0 16.5 34.5 54.3 76.0 100
m. f. 0 0.10 0.23 0.40 0.64 1.0
~----------------------------------------------------------------------------------------------Dilution 10 30 30
Absorbance 1. 80 1. 01 1.86
Dil. x Absorb. (ASS) <a> 18.00 30.30 55.80
103Solub./mol dDr3 (0) 2.069 3.483 6.414
om~e(Salt)/kJ mol-' -3.8'7 -8.41
om~e{2(ClO.)-}/kJ mol-' <c) +0.20 -0.20
Om~e(Cation)/kJ mol-' -4.0'1 -8.21
(a) ASS = Absorbance of saturated solution (b) Calculated from €S22=8700X-'cDr' for [Fe(bipY)3] (ClO.)2 (c) Calculated from reference 7
50 50 10
1. 87 1. 71 1. 84
93.60 85.50 18.40
10.'760 9.862 2.115
-12.25 -11. 58 -0.16
+0.60 +3.60 +12.60
-12.85 -15.18 -12.76
10
(a) (!)
~ "d
~
ii ....... 5 ~ +' orj
~ orj
,0 ;:J ~
~ (I) 0 rI
0 50 100
Vol~ KeOli -i (ClO .. )-
+5
,. a 100
I [Fe(bipY)al (CIO")2 r-I 0 a -5
~ (b)
~ , -10
~ [ Fe (bi py) 3] 2-+ e Co
FIGURE 4.1
Dependence of solubility (a) and transfer chemical potential (b) of [Fe(bipy)a] (Cl04 )2 on methanol composition for binary aqueous
mixture at 298.2 K.
- 52 -
"-.,'1 W
Table 4.2
Dependence of single ion transfer chemical potentials on methanol caDpOsition and derivation of transfer cheDdcal potential from solubility of salts in Dethanol/water solvent mixtures at 298.2K
](ethanol 1 P~P+=P~I K+ 1 (ClO4)-1 OH- [Fe (gDd)31 (ClO4)2 1 K[Cr(]B3)2(SCI)4] Content 1 1 1 1 1 -------------1--------------
(a) 1 (a) 1 (a) 1 (a) 1 ESS4 = 8600](-' cnr' 1 ES17 = 105.5X- 1 cnr'
~Ol~ . ---------1 ----I ------1 -----1 --------------1 ----------------
'it~ m. f. 1 A or C 1 A C A I ASS S C ASS S C
0 0 0 383 13 10 8.1 0.047 -2.0 1.1 +0 -0.1 416 -0.626 -0.54 14.8 -0.64 -1. 74 20 16.5 0.100 -4.1 2.2 +0.10 -0.2 417 -0.642 -0.71 18.0 -1. 61 -3.81 30 25.3 0.160 -6.2 3.3 +0.05 -0.2 444 -1. 110 -1.26 24.1 -3.07 -6.37 40 34.5 0.229 -9.2 4.5 -0.10 +0.1 465 -1. 450 -1.30 31. 2 -4.34 -8.84 50 44.2 0.308 -12.3 5.3 0.00 +0.5 60 54.3 0.400 -14.9 6.2 +0.30 +1.6 3'16.5 +0.11'1 -0.23 45.6 -6.22 -12.42 '10 64.8 0.509 -17.6 '1.4 +0.90 +3.7 80 76.0 0.640 -19.7 8.6 +1.80 +6.3 138 +'1.580 +4.42 48.9 -6.56 -15.16 90 8'1.'1 0.800 -21. 4 9.3 +3.90 +9.3 48 +15.430 +8.33 48.0 -6.47 -15.7'1
~OO 100 1 -22.6 9.8 +6.30 +12.5
ASS = Absorbance of saturated solution.
Transfer cheBical potentials (8_~e/kJ DOl- 1) of; A = anion, C = cation, S = salt.
(a) transfer cheDical potential calculated from reference 7
'~I
..::--
Table 4.3
Derivation of transfer chemical potentials for Fe{I!) diiDine complex cations from solubility measure.ents of their perchlorate salts in methanol/water solvent Dixtures at 298.2 K.
}(eOR Vol"'
0 10 20 30 40 50 60 70 80 90
100
# # # # [Fe(gnd)3] (CIO~)2 [Fe (mBd)3] (Cl04 )2 [Fe(bDi)3] (CI04)2 [Fe(cDi)3] (CI04)2
--------------------1---------------------1---------------------1-----------------------€ss4=8600I-1CDl1 1 €S61=9560X- 1CDl1 1 €S64=10600I-1CDl1 1 €ss2=12470X- 1 Cm- 1
---------------------1---------------------1----_-----------_____ 1 _____________________ _ ASS S C ASS S C
79.34 91.68 109.30 -1. 31 -1. 23
111.34 -2.52 -2.57 1 137.60 -3.02 -2.96 136.90 -2.98 -3.13
134.50 -3.92 -3.80 I 187.25 -5.30 -5.20
98 -1. 57 -1. 91 240.00 -7.15 -7.50
43 +4.55 +1.39 1 151.00 -3.62 -6.78
I
ASS s c
82.5
187.4 -6.10 -6.04
358.5 -10.92 -16.76
579.0 -14.48 -14.82
459.0 -12.76 -15.88
J J
ASS s C
2.44 4.68 -4.85 -4.77 9.67 -10.24 -10.30
23.37 -16.80 -16.95 56.00 -23.29 -23.13
140.50 -30.13 -30.47
173.00 -31.68 -34.84 131.00 -29.61 -36.72
ASS = absorbance of saturated solution.
Transfer chemical potentials of; S = salt. C = cation.
I Salt equilibrated with solution containing laCl04. 0.4 mol dDr 3•
.10
,. I r1 0 S
40
~ 0
~ "-$:1,
[ Fe(hxsbh)]2+-
,. C ~ [Fe (mmi )8J 2+
[Fe(cml)3]2+
-40 [ Fe ( bs b - Ke 2 ) J 7 +
FIGURE 4.2
Transfer chemical potentials for Fe(II) diimine cODplex cations from water into aqueous methanol at 298.2K
'-' G\
Table 4.4
Derivation of transfer chemical potentials for Fe(II) complex cations from solubility measureDents of their perchlorate salts in Dethanol/water solvent mixtures at 298.2K.
[Fe(mpDd)3] (ClO.):2 [Fe (tsbIe) 2] (ClO.)21 [Fe(tsbPh)21 (ClO.)2I[Fe(cxcage)] (PF6)2 [ Fe (Ie:2Bcage) ] 1--------------------
___________________ 1 ____________________ / ___________________ _
--------------I
Vol~1 €ssa=11500~lcDrl
IeOBI--------------------€S91=13170X- 1 CDL 1 €s9s=6490I- 1 C1Il 1 1 €sS2=50'74X- 1 c1Il 1 €442=16100X- 1 CDl 1
0 10 20 30 40 50 60 70 80 90
100 I
ASS S C ASS s C ASS S C ASS S C
23.25 I 76.8 1. 77 0.235 102.0 -2.10 -2.02
29.20 -1.69 -1.75 1116.0 -3.06 -3.12 4.83 -7.46 -7.52 10.515 -5.83 -6.91 148.5 -4.90 -5.05
42.60 -4.50 -4.66 1190.8 -6.76 -6.60 16.44 -16.57 -16.41 10.950 -10.40 -8.92
56.60 -6.61 -6.95 1274.0 -9.45 -9.79 148.90 -24.67 -25.01 11.590 -14.20 -12.38
40.50 -4.12 -7.28 1213.6 -7.60 -10.76 149.00 -24.68 -27.84 120.0 -3.30 -10.40
I 1 i
ASS = absorbance of saturated solution
Transfer cheDdcal potentials of; S = salt, C = cation
--------------ASS S
0.36
0.89 -2.24
2.43 -4.73
8.37 -7.79
L
'" I r1 0 ~ 0
40 80
~ Vol7. MeUH
~ "-
~ E
'0 [ Fe Ole2Bcage) ]
-10
·~[Fe(CXCage)J2+ ~------------
-20
lFe(cmi)3]2-+-
FIGURR 4.3
Transfer cheDdcal potentials for Fe(II) cage complex cations from water into aqueous methanol at 298.2K
- 57 -
Ul CD
](eOH Vol~
Table 4.5
Derivation of transfer che~cal potentials for Fe(II) and Ii(IV) hexadentate complex cations from solubility measurements of their perchlorate salts in
methanol/water solvent ~xtures at 298.2K
[Fe(bxsble») (CI04)21 [Fe (Ke4L)] (CI04)2 1 [Fe (Je2Pb2L)] (CI04)2 [Ii (Je4 L)] (CI04)2
-------------------1-----€ 605=9280 I-1 C.- 1 I €&1e=10558I-l c.r' €&4S=12460K-' CDl l €&00=6300I-1 cr'
-----------------, ---- ----------------ASS S C , ASS s c ASS s C ASS s c
r----- -------------------,-------------------- --------------------- ---------------------; o 10 20 30 40 50 60 70 80 90
100
48.0
56.0 -1.15 -1.19
73.0 -3.12 -2.96
86.0 -4.33 -4.83
64.0 -2.14 -5.74 37.5 +1.84 -5.66 12.9 +9.78 -2.82
I , i
105.0
120.0 -0.99 -1.05
159.5 --3.10 -2.85
273.0 --7.10 -7.44
221.3 -5.54 --8.70
ASS = absorbance of saturated solution
9.83 11.15 --0.94 --0.85 13.95 -2.30 -2.36 18.75 -4.80 -4.95 25.35 -7.04 --6.88
54.90 -12.78 -13.12
86.80 -16.19 -19.35 79.00 -15.49 -22.60 42.33 -10.85 -23.45
Transfer cheDical potentials of; S = salt, C = cation
13.75
14.84 -0.57 -0.77
19.78 -2.70 -2.50
30.85 -6.00 --6.60
30.57 -5.94 -10.32
.. I r1 o a
O 50 100
~~~------------~----------------~ Vol% MeUIl
~- [ Fe (hxsbh) J 2+
[Fe (hxsbMe) J 2+
-10
-20
FIGURE 4.4
Transfer cheBdcal potentials for Ii{IV) and Fe{II) hexadentate complex cations from water into aqueous methanol at 298.2K
- 59 -
+10
gmt
po
I M 0
100
0 Vol~ KeOR a hxsbh
..., ~ hxsb1le
"-~ =1-E
mpmi
co
-10 tsbKe
bipy
-20
phen
tsbPh
-30
FIGURE 4.5
Trends in the transfer chemical potentials for Fe(II) diimine complexes from water to aqueous methanol, indicating dependence of preferential solvation by methanol on the ligand bulk and hydrophilic/hydrophobic character of the periphery of the complex
TABLE 4.6
Derivation of transfer chemical potential (Om~e/kJ mol- 1) for anions from solubilities, at 298.2 K, of [Fe(diimine)n]2+ salts.
Vol~ 1lethanol 0 20 40 60 60 100
[Fe(bipY)31 lASS 5.01 5.10 10.14 15.60 6.70 8.60 [Fe(CI)6101.4H2O IS -0.09 -3.49 -5.63 -2.73 -2.62
E622=8700J{-l CJr l IA +3.82 +4.76 +7.12 +12.45 +10.14 r------------------I--------------------_______________________________ [ Fe (phen) 3] lASS 0.42 0.68 1. 35 1. 77 1. 71
[Fe(CI)6R01.5H2O IS -3.66 -5.78 -7.13 -6.96 E610=11100J{-l c mr 1 IC (a) -4.5 -11.1 -16.9 -20.7 -21. 50
IA +0.6 +5.3 +9.6 +13.7 ~------------------I---------------------------------------------------
Mean value for A +2.3 +5.0 +6.4 +12.9
lASS 1. 11 1.61 3.74 7.60 9.65 [Fe(bipY)3] (PF6)2 IS -2.77 -9.05 -14.52 -16.10 ES22=8700J{-l CD-l IA +0.56 -0.37 -0.96 -0.70 ------------------1---------------------------------------------------[Fe(tsbh)2] (PF6)2 lASS 19.74 30.24 72.93 127.70 189.90 ES92=14700J(-l c mr' IS -3.27 -9.82 -13.96 -16.93
IA +0.52 -1.11 -0.66 -0.64 ~------------------I---------------------------------------------------
Mean value for A +0.54 -0.74 -0.91 -0.67
lASS 15.73 32.34 45.57 61. 74 57.33 [Fe(tsbh)21 (SCI)2 IS -5.42 -7.90 -10.22 -9.51 ES92=14700J{-l cmL l IA -0.50 -0.23 +1.10 +3.13
I lASS 288 660 1370 1762 1674 507
[Fe(bipY)3] (BF4)2 IS -6.16 -11. 56 -13.46 -13.92 -4.20 ES22=8700J{-l CD-l IA -1.12 -1.65 -0.35 +0.63 +4.26
J
ASS = Absorbance of saturated solution
Transfer chemical potential for; S = salt, A = anion, C = cation
(a) froD reference 5
- 61 -
S0,42- [Fe <elf) sBOJ 2-
+10
t5
po
IJ I.) Vol1. MeOH ~ 0 ~--------~--~~-=~~~~~--------
" ~:t C
Co
-4
FIGURE 4.6
Transfer chemical potentials for some simple and complex anions froa water into aqueous methanol, at 298.2K
- 62 -
compared with previously reported cage complex<S) and their tris-ligand
analogues in Figure 4.3. The larger [Fe(cxcage)]2+ cation is more
preferentially solvated by methanol than its smaller [Fe(Me2cage)]2+
analogue. However both are much less solvated than their corresponding
tris-ligand Fe(II) counterparts, [Fe(bmi)3]2+ and [Fe(cmi)3]2+. This
relatively small preference of the cage cations for methanol is
attributed to the presence of six non-coordinated nitrogen atoms on the
ligand periphery which contribute marked hydrophilic properties to these
otherwise hydrophobic cations. This hypothesis is supported by the
transfer chemical potential for uncharged encapsulated complex
Fe (Ke2Bcage) which on its periphery instead of non-coordinated nitrogens
has oxygen atoms. The trend for the uncharged cage complex comes close
to that of [Fe(Me2cage)] indicating similarity in hydrophobic and
hydrophilic properties.
Solubility data for metal hexadentate bis-perchlorate salts are
tabulated in Table 4.5. The smaller hexadentate complex [Fe(hxsbh)]2+
with its hydrophobic periphery shows only modest preferential solvation
by methanol. As expected from earlier established order, its analogue
[Fe(hxsbm)]2+ shows significantly increasingly favourable stabilisation
on going to methanal rich mixtures, Figure 4.4. The same figure shows
hexadentate oxime ligand which shows preferential salvation pattern as
above. [Fe(II)(Xe4L»)2+ and [Ni(IV) (Me4L») 2+ complexes have similar
preferential salvation by methanol even though the ligand in the latter
complex has two protons less. Solubility analyses were attempted with
other Ri(IV) complexes containing bidentate and terdentate ligands but
these were found to oxidise methanol. Preferential solvation by
methanol of the above mentioned Fe(II) diimine complexes, in terms of
- 63 -
ligand size and hydrophobic/hydrophilic character of the complex, is
summarised in Figure 4.5.
Table 4.6 reports solubilities of the tetrafluoroborate,
hexafluorophosphates, thiocyanate and nitroprussides of iron (II)
diimine cation and om~e(anion) derived from these solubilities and
oM~e{Fe(II) cation) from Tables 4.1-4. The oM~e(anion) for BF.- and
SCN-- and the average of om~e(anion) for PF6- and [Fe(CN)sNO]-, with a
small selection of other anions(4', are shown in Figure 4.6.
Nitroprusside, dinegative anion is destabilised, as expected from its
hydrophilic periphery and its trend is similar to that of 80.2 - and
comes between the medium sized, hydrophilic, thiosulphate and
peroxodisulphate. Thiocyanate ion and BF4- parallel CN- and CIO.- as
expected while PF6- does not show preferential solvation for water or
methanol as its Om~e is along the axis for up to 80% indicating that
there is strong water solvation of this anion.
- 64 -
4.4 RESULTS AND DISCUSSION FOR AQUEOUS i-PrOH AND t-BuOH
The solubility data <mol dm- 1) or (ASS - absorbance units) of
Ph4AsCl04, [Co(en)3] (Cl04)3 and [Co(en)3] (Ph4B)3 and derived transfer
chemical potentials are presented in Table 4.7 and Table 4.10 for
aqueous i-PrOH and t-BuOH solvent mixtures respectively. Although the
dielectric constant of alcohol-rich mixtures is much lower than that for
water-rich media, it is assumed that these salts are completely
dissociated in the whole water-alcohol range. The 8M~e values for the
above salts are calculated using the equation 1.18, where n = 1 for
Ph4AsCI04 and n = 4 for Co(III) salts. The transfer chemical potentials
for tetraphenylarsonium tetraphenylboronate (TATB) values are obtained
using the equation 4.1,
oM~e(Ph4AsPh4B) = 1/3{38M~e(Ph4AsCl04) + 8M~e([Co(en)3] (Ph4B)3) -
- 6'm~e([Co(en)a] (CI04)3)} .. (4.1)
while the split for TATB assumption is obtained using the relationship,
Transfer chemical potential values for Cl04-, [Co(en)a]3+, and
those of other anions and cations were obtained using the above TATE
assumption and 8M~e values of the respective salt derived from
solubility data as reported in Tables 4.7-9 for i-PrOH and Tables 4.10-
11 for t-BuOH. Plots of 6'm~e(ion) vs Vol% of alcohol are shown in
Figure 4.7 and Figure 4.8 for i-PrOH and t-BuOH respectively, which show
the preferential solvation of ions in these solvent mixtures.
- 65 -
c;'I ,:;'\
TABLE 4.7
Derivation of transfer cheDical potentials for simple and complex ions fram solubility measureDents of salts in aqueous i-propanol; at 298.2 K.
i-Propanol Content
Vol1. Yt1. 11. f.
1 Ph.As+/1 PhAAsCIO", 1 PhAB- I-
I E26&=3020X- 1 ClJl 1
A/C AS) S A
1 [Co(en)3] {PhAB)31 [Co{en)3] (CIO.)3 1 KCIO. 1 RbCIO. -------------1--------------1 ---------------1-----------Atomic Absorpionl £465 = 86](-1 CD.- 1 1 Flame Photometry 1 Atomic Absorption ------------1------------1 ------------1--------------
102Sol- S 1 ASS SCI 102801- SCI 102801- S C ~ ________________________________________ --------------_1 1 __________________ ,
32.04 7.160 0 0 0 0.355 0.0318 21.80 10 8.06 0.026 I 16.67 +1.33 5.010 +1.75 20 16.47 0.056 -5.06 1 0.873 -4.75 +0.32 0.1292 -13.9 25.60 +2.22 +1. 28 11. 00 +3.39 +3.081 4.230 +2.61 +2.40 30 25.27 0.092 1 1 9.23 +4.26 I 3.220 +3.96 40 34.37 0.136 1-11.80 I 2.474 -9.62 +2.18 1.440 -37.8 21.15 +4.12 -2.41 7.82 +5.08 +2.911 2.896 +4.49 +2.31 50 44.10 0.191 I I 6.15 +6.27 I 2.370 +5.48 60 54.20 0.262 1-14.77 I 3.268 -11.00 +3.77 4.406 -48.8 16.05 +6.85 -4.46 4.60 +7.71 +3.941 1.667 +7.20 +3.43 70 64.80 0.356 I I 2.86 +10.06 I 0.936 +10.08 80 75.94 0.486 1-16.81 3.135 -11.11 +5.70 6.271 -52.4 6.80 +15.36 -1.74 1.54- +13.16 +7.431 0.610 +12.20 +6.46 90 87.66 0.681 1.934 -8.69 0.795 +16.40 0.117 +20.37
100 100 1 0.632 -3.15 0.018 +35.50 0.006 +35.24
ASS = Absorbance of saturated solution Transfer chemical potential, (o_~e/kJ DOI- 1 ), for; S = salt, C = cation, A = anion a) SOlubility/DOl dDr 3
I
()'\ -..:J
·TABLE 4.8
Derivation of transfer cheDical potentials for Fe(II) diiDdne cations from solubility Deasurements of perchlorate salts in aqueous i-propanol;
at 298.2 K.
I I [Fe(gDd)3] (CIO£)2 1 (Fe (phen) 3] (CI04)21(Fe(bipY)3] {CI04)2 1 [Fe (tsbPh) 21 {CI04)2 I(Fe(tsbKe)2] {CI04)21[Fe(tsbH)21 (CIO£)
. .
1--------------1------------1 ------------1 --------------1-----------1 ----------- ~
i-PrOHIE&&6 = 8600X-'cur 1 Es,o= 11500X-'cur' ES2Z= 8700X- 1 cur' ES9& = 6940X-'c--' 'ES9'= 13170X-'c.r' IEs92 = 14700X-'cur 1
Vol~ 1 ____________ _ --------------------. ASS S C ASS S c ASS S c ASS S c ASS S C ASS S C
0 441 9.2 17.7 1. 77 78 55.00 10 20 435 +0.10 -0.53 24.4 -7.25 -7.881 33.2 -4.68 -5.31 1 7.08 -10.31 -10.93 107 -2.35 -2.98 99.95 -4.44 -5.08 30 40 408 +0.58 -3.77 54.0 -13.15 -17.501 57.2 -8.72 -13.07 124.75 -19.61 -23.96 146 -4.67 -8.92 1 96.77 -4.20 -8.56 50 60 267 +3.73 -3.81 49.0 -12.43 -19.971 47.6 -7.35 -14.89 137.00 -22.60 -30.14 97 -1.62 -9.16 69.60 -1.75 -9.29 70 80 56 +15.31 +3.91 16.2 -4.21 -15.611 14.3 +1.57 -9.83 126.40 -20.09 -31.49 24 +8.83 -2.57 1 24.53 +6.00 -5.40 90
100 1 J ~ __ I l 1
ASS = Absorbance of saturated solution Transfer chemical potential. (8_~e/kJ DOI-'). for; S = salt. C = cation
0' (JJ
TABLE 4.9
Derivation of transfer cheDical potentials for complex ions from solubility measurements of salts in aqueous i-propanol; at 298.2 K.
I I [Fe(.:pam!)3] (ClO~d21 [Fe(Dpmi)3] (Cl04)21 [Co(I'R3)4CCh] (ClO.) K[Cr(IIb)2(SCI).] I [Co(en)3] 13 I [Co(!H3)td (CI04}; 1-------------1 -------------1 --------------1 I-------------I-----------~"
i-PrOHI£s72 = 13000~lcDrll -sse= 11500~lcmr11 £S:2A. = 111~lcmr1 I£S17 = 105.5X- 1 CDr 1 I £.6S = 86X- 1 CDL 1 I £47S = 58X- 1 cmr l
Vol ~ 1--------------1 -----------1--------------1----------1-----------ASS S C ASS S C ASS SCI ~ S A lASS S A lASS S C
0 88.0 1 26.1 5.58 12.0 5.49 3.30 10 3.84 +1.85 7.0 +2.67 20 I 102.6 -1.14 -2.78 31.8 -1.49 -2.10 3.36 +2.52 +2.20 7.4 +2.38 -0.70 3.30 +5.05 +1.26 2.44 +2.99 +2.03 30 2.61 +3.76 10.5 +0.66 40 104.8 -1.30 -5.66 40.4 -3.32 -7.67 2.07 +4.91 +2.74 16.2 -1. 49 -4.39 2.82 +6.60 +3.00 1.89 +5.52 -1. 02 50 20.8 -2.73 60 90.4 -0.20 -7.74 1 27.6 0.42 -7.96 0.96 +8.72 +4.95 24.0 -3.44 -7.38 1.96 +10.21 +4.89 1.17 +10.28 -1.03 70 24.9 -3.62 80 29.8 +8.05 -3.35 7.15 +9.25 -1.78 0.16 +17.60 +11.90 22.7 -3.15 -10.58 0.61 +21.78 +7.84 0.20 +27.60 +10.50 90 22.2 -3.05
100
ASS = Absorbance of saturated solution Transfer che.dcal potential, (o_p8 /kJ DDI- 1
), for; S = salt, C = cation, A = anion
po
I r1 0 a I; ~
" ~:l. ~
<0
OH-
[Co (ltH) ... ab,] +
+10
[ Fe (gm!) 3] 2+
0 [Co(en)3] 3+
[ Fe (tsbMe) 2] 2+
[ Fe (bi py) 3) 2+
-10
[ Cr (Bib) 2 (SCB) ... )-
[Fe (tsbPh) 2] 2+
FIGURE 4.7
Dependence on vol~ of single-ion transfer chemical potentials for complex and simple ions in water/i-propanol mixtures at 298.2 K.
-...J o
t-Butanol Content
Vol~ Yt~ m. f.
Ph.As+1 Ph..s.B-
A/C
TABLE 4.10
Derivation of transfer chemical potentials for simple and complex ions froD solubility measureuents of salts in aqueous t-BuOHj at 298.2 K.
1 Ph..AsCl04 [Co(en)3] (Ph.B)3 1 [Co(en)3] (Cl04)3 1 KCl04 1 RbCI04
--------1 ~ ------------1------------E2&6 = 3020X- 1 CDr 1 AtoDic Absorpt. 1 E4&& = 86]{-1 CD- 1 1 FlaE PhotoB!try 1 Atomic Absorption
---------1 ~ ------------1-------------ASS S A 102 501. S 1 ASS SCI 102 801 SCI 102 S01 S C
r--------------- ------- _________________ ---------------,1 1 I __________________ ~ o 0
10 8.06 20 16.47 30 25.27 40 34.47 50 44.10 60 54.20 70 64.80 80 75.94 90 87.66
100 100
o 0.021 0.046 0.076 0.113 0.161 0.22 0.31 0.43 0.63 1
-2.10 -6.10
-11. 40 -13.89 -15.40# -14.80 -14.10# -13.47 -13.72
0.355
1.056 -5.40 +0.70
3.203 -10.90 +2.99
3.403 -11.20 +3.60
1.820 -8.10 +5.37 0.828 -4.20 +9.52 0.289 +1. 01
i
ASS = Absorbance of saturated solution
0.032 0.064 0.164 0.778 2.565
-6.80 -16.22 -31.62 -43.45
1 1 1 1 1 1
2.829 -44.42 1 1
1. 382 -37. 32 1. 020 -34.30 0.081 -9.21
I
32.50 27.20 21.30 16.68 15.74
21.80 +1.77
1 7.160 1
+4.18 +6.61
+2.08 112.60 +2.70 +2.001 4.245 +2.59
+7.19 -1. 78 9.53 +4.10 +1.101 3.366 +3.74 1
+1.85
+0.75
10.95 +10.78 -0.02 1 5.42 +6.90 +3.301 1.958 +6.43 +2.83 1
4.68 +19.20 +3.09 1 1.32 +13.90 +8.501 0.533 +12.87 +7.50 0.91 +35.41 +6.86 I 0.27 +21.70 +12.201 0.091 +21.65 +12.13
1 0.012 +31.00 i
Transfer che~cal potential. (o_p8 /kJ DDl- 1). for; S = salt. C = cation. A = anion
Sol = Solubility(mol dg-3)
-..,J I -"
1
TABLE 4.11
Derivation of transfer cheDdcal potentials for complex ions from solubility measureDents of salts in aqueous t-BuOHj at 298.2 K.
1 [Fe(gDd)31 (CI04)2 1 [Fe(bipY)3] (ClO4)2 1 Kl Cr (1H3)2 (SCI) 4] 1 [Co(IH3)4CObl (CI04)1 [Co(JH3)sJ (ClO4)3 1-------------1----------1-------------1 I
t-BuOHI ES&6 = 8600X- 1 CDr 1 1 ES22 = 8700J[-l C.-l 1 ES17 = 105.5X- 1 cDr 1 I ES24 = llll1- 1 cDr 1 1 E47S = 5811- 1 c.- 1
Vol~ 1--------------1 I 1 ASS S C 1 ASS S C ASS S A I ASS S C ASS S C
I 0 441. 0 17.70 11. 70 I 5.57 3.40
10 22.56 -1.80 6.70 +2.76 +0.96 I 4.21 +1.39 2.77 +2.04 20 4-42.5 -0.03 -1. 43 32.72 -4.57 -5.9'1 7.13 +2.46 +0.46 I 3.63 +2.12 +1.42 2.18 +4.38 +2.28 30 40.00 -6.06 11. 40 +0.13 -1. 38 I 2.86 +3.30 2.09 +4.83 40 417.0 +0.41 -5.50 45.60 -7.03 -13.01 15.75 -1.47 -2.57 2.39 +4.19 +1.20 1.82 +6.19 -2.78 50 60 232.0 +4.78 -2.42 27.90 -3.38 -10.58 21. 75 -3.07 -6.37 1.62 +6.12 +2.52 1.08 +11.47 +0.67 70 80 36.3 +18.57 +7.83 5.10 +9.25 -1.49 20.40 -2.75 -11. 25 0.14 +17.91 +12.54 0.25 +25.83 +9.72 90 0.92 +23.80 +4.76 13.80 -0.82 -13.02
100
ASS = Absorbance of saturated solution Transfer che~cal potential, (8m~e/kJ DDl- 1
), for; S = salt, C = cation, A = anion
+10
-10
-20
FIGURE 4.8
[Co (I16> 4C(h] +
[Fe (gmi )3] 2+
Cl04 -
[Co (en) 3] 3+
[Fe(bipY)31 2 +
100
t-BuOH
Dependence on vall of single-ion transfer chemical potentials for complex and simple ions in water/t-BuOH mixtures at 298.2 K.
- 72 -
--.l w
I
TABLE 4.12
Derivation of transfer cheDdcal potentials for complex ions from solubility measurements of salts in aqueous ethanol; at 298.2 K.
I P1l4.As+ II K+ (ClOA)- OH- [Fe (gDd)3] (ClOA)2 I [[Cr(IB3)2(SCI)A] I [Co(IB3)4~]Cl04 I[Fe(bipY)3]2+ Ethanol I P1l4B- -----------------1---------------1---------------1 Content I (a) (a) (a) €ssA=8600X- 1 cur 1 I €S17=105.5Jl l cm- 1 I €S24=111I-l cur 1 I (b)
t---+---+-----I--- -----------1 -I 1----------I Vol"' 'it"' m.f. I A/C I C A A
I ASS S C IASJ S A lASS SCI C
0 0 0 408 13.6 5.55 10 8.05 0.072 -2.0 +1. 15 +0.10 11. 4 +0.74 -0.41 I 4.17 +1.42 +1.32 -1.6 20 16.45 0.072 -4.6 +2.00 +0.20 +1.40 408 0 -0.40 12.5 +0.42 -1. 52 I 3.57 +2.19 +2.00 -3.4 30 25.23 0.117 -7.8 +3.00 +0.35 +2.45 14.6 -0.35 -3.35 I 2.64 +3.68 +3.33 -5.6 40 34.43 0.170 -11. 6 +3.45 +0.90 +4.15 408 0 -1.80 20.8 -2.12 -5.57 I 1. 98 +5.12 +4.12 -8.6 50 44.43 0.235 -15.1 +3.30 +2.10 +6.90 I -12.7 60 54.15 0.316 -17.5 +3.90 +2.95 315 +1.92 -3.98 29.4 -3.82 -7.72 I 0.90 +9.02 +6.10 -14.8 70 64.76 0.418 -19.3 +5.65 +3.50 I -14.6 80 75.90 0.552 -20.3 +8.00 +4.30 85 +11.66 +3.06 30.6 -4.02 -12.20 I 0.23 +15.77 +11.52 -14.2 90 87.63 0.735 -20.5 +11. 70 +6.00 19.8 +22.50 +10.50 22.8 -2.56 -14.26
100 100 1. 000 I -20.9 +16.60 +9.90 J ----
ASJ = Absorbance of saturated solution Transfer che~cal potential, (8_~e/kJ mol-I). for; S = salt, C = cation, A = anion (a) from reference 26.27 (b) from reference 29
-..J J>
1
TABLE 4.13
Derivation of transfer cheDical potentials for complex ions from solubility measurements of salts in aqueous acetone; at 298.2 K.
1 1 Acetone 1 P14As+' K+ ClOA.- [Fe(~)3](CIOA)2 [Fe (bipy) 3] (ClOA)2 1 [Fe(phen)3] (ClOA)2IK[Cr(IB3)2(SCI)A.] Content 1 PluB- I 1
1 (a) (b) E&&6 = 8600Jr 1 CDr1 E&22 = 8700Jr1CDr 1 IE&lO= 11500X-1CDr1IEs17= 105.5I-1cDr 1
1 1---------Vol1. 'it1. Lf. I A/C C A " ASS S C ASS S C 1 ASS S C ASS S C
1 I 0 0 0 I, 439 17.7 1 9.2 13.0
10 8.3 0.026 I' 555 -1.74 -2.54 34.1 -4.85 5.651 -11.1 -11.9 10.3 +1.15 +1.85 20 16.6 0.058 I -7.5 -1.5 +0.9 816 -4.61 -6.41 62.6 -9.39 -11. 201 -15.4 -17.2 18.0 -1.60 -0.10 30 25.3 0.095 1 1318 -8.1'7 -11. 37 122.0 -14.35 -17.551 -20.4 -23.6 36.9 -5.17 -2.67 40 34.6 0.144 I -15.8 -3.6 +2.4 1542 -9.34 -14.14 26'7.0 -20.20 -25.001 -25.6 -30.4 51.3 -6.80 -3.20 50 44.2 0.197 I {PJ -29.0 -34.8 60 54.4 0.269 -23.3 -4.5 +3.3 2130 -11.74 -18.34- 588.0 -26.10 -32.701 -32.6 -39.2 71.0 -8.40 -3.90 70 64.7 0.364 1 I -35.3 -42.5 80 76.1 0.495 I -29.2 -3.6 +4.4 2263 -12.19 -21.00 672.0 -2'7.00 -35.801 -35.9 -44.7 77.3 -8.84 -5.24 90 87.'7 0.688 1 691 -3.37 -15.37 401.0 -23.20 -35.201 -33.4 -45.4 97.5 -9.98 -7.58
100 100 1 1 20 +23.10 34.3 -4.92 1 -17.0 1 I 1 I I I
ASS = Absorbance of saturated solution Transfer chewdcal potential, (8_~s'kJ mol- 1
), for; S = salt, C = cation, A = anion (a) and (b) from reference 28,29. (p) calculated from reference 25
Tables 4.12 and 4.13 summarise the solubility results and derived
OM~e for selected complexes in ethanol and acetone water solvent
mixtures. The single ion OM~e for ethanol cosolvent are derived using
literature(2s.27) results based on the TATB assumption, as are those for
acetone water solvent mixtures(2e.29). Further solubility results for
several complexes in binary solvent mixtures of other organic cosolvents
are tabulated in Appendix 2, Table A2(I).
4.5 oM~e(OH-) for Ethanol and Acetone cosolvents.
Transfer chemical potentials of (H+OH-) are calculated from already
published data of pKw, ionization constant for water, in aqueous solvent
mixtures(30) by using the relationship(2e.31), i.e.
oM~e(H+OH-) = -RT{lnKw(mix) - InKw(aq)}
but pKw = -logloKw
Where K is equilibrium constant, therefore
oM~e(H+OH-) = 2.303RT{pKw(mix) - pKw(aq)}
om~e(H+OH-) = -RTln{Kw(mix) / Kw(aq)}
= RT(lnl0)[pKw(mix) - pKw(aq)]
The data for the above calculations are presented in Appendix 2.
The single ion transfer chemical potentials for hydroxide in ethanol
water solvent mixtures are derived from the relationship as shown in
Figure 4.9, [om~e(OH-) = om~e(H+OH-) - om~e(H+)]. Transfer chemical
potentials for OH- ion in aqueous acetone are calculated from
oM~e(H+OH-), om~e(H+Cl-)(32) and om~e(Cl-) as shown in Figure 4.10. For
both of the above calculations conversion from wt% to vol% were carried
- 75 -
+10 OH-
,. I W-OH-r-l -f5
0 ~
I-) ~
~ 0
20 Wt~ EtOH :l. ~
'0 11+
-5
FIGURE 4.9
Dervation of transfer chemical potentials for hydroxide ion from water to aqueous ethanol, at 298.2 K, [for data see Appendix 2, Table A2(II)]
- 76 -
+20
+10
I rl o ~
-10
FIGURE 4.10
OH--
C1 ~
..-----.-- lIe 1
40 'It%. Acetone
H+-
Derivation of transfer chemical potentials for hydroxide ion fro. water to aqueous acetone, at 298.2 K, [for the data and conversion to Vol~ see Appendix 2, Table A2(II)]
- 77 -
out by graphical intrapolations. Table 4.14 summarises Om~e(OH-) in
five binary solvent mixtures up to 60% by volume of the organic
casal vent.
TABLE 4-.14
Transfer chemical potentials of OH- ion in aqueous solvent mixtures of MeOH, EtOH, i-PrOH, t-BuOH and acetone
Vol1. 10 20 30 40 50 60
-----------------------------------------------------------------)(ethanol .. -0.1
Ethanol b +0.'75
i-Propanol c +1.80
t-Butanol - +0.65
Acetone b +3.20
a - from reference 7, c - from reference 28,
4.6 DISCUSSION
-0.2 -0.2 +0.1 +0.5 +1. 6
+1. 40 +2.45 +4.15 +6.90
+3.35 +5.20 +7.20 +9.25 +11. 60°
+2.30 +5.60 +7.60 +9.50
+6.80 +10.85 +15.50 +20.10
b - for calculation see Appendix 2 d - extrapolated, e - reference 29
The interalcohol comparison is carried out using the plots of 6m~e
vs mole fraction of alcohol. Mole fraction is preferred to weight
percentage or volume percentage, the advantage being that the comparison
is directly concerned with the number of moles of hydroxyl groups per
mole of the cosolvent mixture. Figure 4.11 shows the plot of volume %
vs mole fraction for MeOH, EtOH, i-PrOH and t-BuOH. Thus, the
difference in number of moles, of groups for cosolvent interaction, per
mole of cosolvent mixture shows a significant decrease from methanol to
t-butyl alcohol, ie for t-BuOH 80% volume is only 0.43 on our mole
fraction scale while methanol 80% volume is 0.6. In the following
discussion a comparison for several ions will be based on plots of 6~~e
vs mole fraction.
- 78 -
100
80
+J I=l OJ I>
r-1 0 (J) 0 (J
= ](eOH ~ r-1 = EtOH 0 40
I>- = i-PrOH = t-BuOH
20
o ________ ~ ______ ~ ______ ~0~.6_· ______ 0~.8 ______ ~1
m.f. cosolvent
FIGURE 4.11
Volume percent vs DOle fraction for several aqueous cosolvents.
-79-
~
I 0.2 0.4 0.6 0.8 0
~ m. f. casal vent 0 a f-)
~ "-
~ s
co -10
t-BuUH
i-PrOH
-20 ---EtUH
. XeOH
FIGURE 4.12
Dependence of single-ion transfer chemical potentials for Ph4P+/PhAAs+/Ph4 B- ion on composition of aqueous casal vents ,
at 298.2 K
- 80 -
i-PrOU
+10
~
I M
~ ~ ~
" 1\ +5
I ~ 0 v
~ E J{eOH
CO
0 m. f. cosolvent
FIGURE 4.13
Dependence of single-ion transfer chemical potentials for OH- anions on composition of aqueous cosolvents; at 298.2 K
- 81 -
I r-1 0 S
'i ~ "-t ~
" co ~ N
c.o
o = leOH
:: ~~~H / ,.
1;/ I +51 A = t-BuOH
r-1 0
0.5 1 S 0
f casal vent m •. r, :.,.I ~
" (a) ~~
a \<J
FIGURE 4.14
Dependence of single-ion transfer chemical potentials for [Fe(gDd)3]2- (a) and [Fe(bipY)3]2- (b) cations on cODpOsition of aqueous cosolvents; at 298.2 K
+5
I (b)
I 0.2 0.5 1
o l\ I m. f. casal vent
-5
-10
EtOH
-15 ){eOH
(9 = ]leOH • = EtOH ~ = i-PrOH £ = t-BuOH o = Acetone
+5 ~
rl
0 a
c;J I.."J 50
~ 0 ~ "-~ ~
...... ~
~ -5 J \. \.~ " '0
-10
(b)
100
Vol% casal vent
--............. -
(a)
o = JlrOB
I +101 • = EtOB ~ = i-PrOB £ = t-BuOB v = Acetone
~
o +5 a Ij
~
" I~ 50 100
t 0 I ~ ~
'0
-5
~ 101% casal vent
FIGURE 4.15
Dependence of single-ion transfer chemical potentials for [Co(IB3).COb]+ (a) and [Cr(IHa)2(SCI).]- ions on composition of aqueous cosol vents j at 298.2 K
For Ph4B- = Ph4As+,Ph4P+
Plots of Om~e values for (Ph4B-/Ph4As+/Ph4P+) ions vs mole fraction
of the organic cosolvent, Figure 4.12, show that these largely
hydrophobic ions are preferentially solvated by an organic casal vent.
In the low region of organic cosolvents an increase in stabilisation of
OM~e is in the order of t-BuOH > i-PrOH > EtOH > MeOH; whereas in high
region of organic casal vents the order of stabilisation is reversed. As
the Born-type interactions are probably very small for these largely
hydrophobic ions, the observed pronounced stabilisation, and the
relative order in this series of monols can be attributed to the
combined effect of dispersion interactions and the cavity forming
interactions. Since the polarisation of the cosolvents dictating the
dispersion interactions, and the relative content of the H-bonding
association dictating the cavity-forming interactions this series of
monols are in the order of the increasing size of the cosolvents. Thus,
the order of stabilisation of Ph4B- in these co-solvents should be
t-BuOH > i-PrOH > BtOH ) MeOH, which is found to be the case.
For OR- anion
Preferential salvation of hydroxide ion by water is obvious from
Figure 4.13. The OM~e composition profile for this hydrophilic ion
exhibits increasingly positive trends indicating increased
destabilisation, The order of destabilisation in these binary systems
is MeOR < EtOR < i-PrOH < t-BuOH < acetone. Destabilisation and the
relative order suggest that solvation of OH- is dictated by the combined
effect of decreased acidity or anion H-centre type acid base
- 84 -
interactions(33) and the increased Born-type electrostatic contributions
of the respective cosolvent system.
For [Fe(gmi)3]2+ and [Fe(bipY)3]2+ cations
We have seen earlier in this chapter that [Fe(gmi)3]2+ complex
cation, the smallest and therefore most hydrophilic of the Fe(II)
diimine complexes, shows lack of preferential solvation by methanol. On
the other hand the hydrophobic [Fe(bipY)3]2+ complex shows a pronounced
preferential solvation by methanol. A similar pattern is observed for
the other aqueous cosolvent systems as shown in Figures 4.14(a) and
4.14(b) for gmi and bipy complexes respectively. Both complexes are
progressively stabilised in low organic cosolvent region in order
acetone> t-BuOH > i-PrOH > EtOH > MeOH while at higher organic
cosolvent fraction stabilisation is observed for both complex cations in
reverse order, with the inflection occurring at 40-60% by volume.
For [Co(NH3)4C03]+ and [Cr(NH3)2(SCN)4]- ions
The two above complexes, both of which are hydrophilic, show an
unusual behaviour as seen in Figure 4.15. The monopositively charged
[Co(IH3)4C03]+ ion is destabilised with an increase in the organic
cosolvent in the order MeOH > EtOH > i-PrOH > t-BuOH > acetone. An
initial destabilisation in low acetone region is followed by
stabilisation. This sequence of destabilisation is probably due to the
decrease in dielectric constant of these media. Such interactions on
the cation represent the stabilising influence with an increased
basicity of the media, as dictated by the cation - 0 - centers type
interactions(33)
- 85 -
On the other hand the [Cr(NH3)2(SCN)4]- anion is progressively
stabilised, except for an initial hump in t-BuOH and acetone, as shown
in Figure 4.15(b), the order of stablisation being KeOH ) EtOH ) i-PrOH
> t-BuOH ) acetone. The solvation of this complex is probably dictated
by the effect of H-bonding interactions, which increases in the order
acetone < t-BuOH < i-PrOH < EtOH < KeOH, and dispersion interaction
which is dictated by hydrophobicity decrease in the same order.
The initial humps for t-BuOH, i-PrOH and acetone may be partly due
to the structural change involved in these cosolvents. The similar
initial hump, for transfer chemical potential, is observed for
[Co(phen)3]3+<S), while [Co(en)3]2+ and (Co(NH3)S]3+ show a roller
coaster behaviour in aqueous i-PrOH and t-BuOH systems, as seen in
Figures 4.7 and 4.8.
- 86 -
REFERENCES
1. J. Burgess and C. D. Hubbard, J. Chern. Soc. Chem. Comm., (1983)1482
2. J. Burgess and C. D. Hubbard, J. Am. Chern. Soc., 106(1984)1717
3. M. J. Blandamer, J. Burgess and E.-E. A. abu-Garib, Trans. Met. Chem., 9(1984)193
4. J. Burgess and E.-E. A. Abu-Garib, Trans. Met. Chem., 9(1984)234
5. E.-E. A. Abu-Garib, M. J. Blandamer, J. BurgesS,N. Gosal, P. Guardado and F. Sanchez, Trans. Met. Chem., 9(1984)306
6. N. Gosal, Ph.D. Thesis, University of Leicester, 1985. R. Bin-Ali, Ph.D. Thesis, University of Leicester, 1986.
7. M. H. Abraham, T. Hill, H. C. Ling, R. A. Schulz and R. A. C. Watt, 1. Chem. Soc., Faraday Trans. I, 80(1984)489
8. O. Vollarova and J. Benko, J. Chem. Soc. Dalton Trans., (1983)2359
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
I. N. Basu Mullick and K. K. Kundu, Indian J. Chem., 23A(1984)812
F. Franks and D. J. G. Ives, Quarterly Reviews, XX(1966)1.
P. Krumholz, J. Am. Chem. Soc., 75(1953)2163
P. Krumholz, O. A. Serra and M. A. DePaoli, Inorg. Chim. Acta., 15 (1975)25
P. Krumholz, Inorg. Chem., 4(1965)609
P. Krumholz, Inorg. Chem., 4(1965)612
M. 1. Blandamer, J. Burgess, R. I. Hains, F. M. Mekhail and P. Askalani, J.Chem. Soc., Dalton Trans., (1978)1001
J. G. Mohanty, R. P. Singh and A. Chakravorty, Inorg. Chem., 14(975)2178
A. N. Singh, P. P. Singh, J. G. Mohanty and A. Chakravorty, Inorg. Chern., 16(1977)2597.
E. R. Gardener, F. M. Mekhail and J. Burgess, Internat. J. Chem. Kinetics, 6(1974)133
H. Saarinen, J. Korvenranta and E. Nasakkala, Acta Chern. Scand., A34 (1980)443
A. G. Lappin, M. C. M. Laranjeira and L. Yonde-Owen, J. Chem. Soc., Dalton Trans., (1981)721
87
21. S. C. Jackels and N. J. Rose, Inorg. Chern., 12(1073)1232. S. C. Jackels, J. Zektzer and N.J. Rose, Inorg. Synth., 17(1977)139
22. J. Bjerrurn and J. P. McReynolds, Inorganic Synthesis, Vol.2, p.216
23. Y. Shirnura and R. Tschida, Bull. Chern. Soc. Japan, 28(1955)572
24. D. Glick, J. BioI. Chern., 156(1944)650
25. F. M. van Meter and Neumann, J. Am. Chern. Soc., 98(1976)1382
26. O. Popovych, A. Gibofsky, D. M. Berne, Anal. Chern., 44(1972)811
27 D. Elvidge, University of Leicester, Private Communication.
28 B. Clark, Ph.D. Thesis, University of Leicester, 1985.
29. J. Burgess, Unpublished Work.
30. E. M. Woolley, D. G. Hurcot and L. G. Hepler, J. Physic. Chern., 74(1970)3908
31. M. J. Blandamer, Private Communication.
32. C. F. Wells, J. Chern. Soc., Faraday Trans. I, 70(1984)694
33. A. K. Das and K. K. Kundu, J. Soln. Chern., 6(1976)431.
CHAPTER
5
Ambient and High Pressure Kinetics of
Reaction Between Hydroxide and [Fe(gmi)3]2+
Complex in Several Aqueous Cosolvent Systems
5.1 INTRODUCTION
The kinetics of reaction of low spin Fe(II) complexes with the
diimine ligands, phen and bipy, have been the subject of intensive study
for many years Cl-
3 ). The kinetics of reactions of iron (II) complexes
of other diimine ligands have also been investigated(4-S). Despite all
these efforts, there still remain problems relating to the mechanism of
substitution of these complexes, one of which concerns reactions with
hydroxide or cyanide. The dominant term in the rate law, equation 5.1,
for reactions of iron(II) diimine complexes with such strong
nucleophiles is second order, indicating an associative attack(7.s>.
-d[complex]/dt = (k, + k2[Nu])[complex] ..... (5. 1)
Although the k2 term, at moderate concentrations of the nucleophile, can
be assigned with confidence to a bimolecular process, the initial site
of attack whether it is on the iron atom or some position on the ligand
is a question of debate(9).
In this chapter, reaction kinetics at atmospheric and at elevated
pressures are reported for hydroxide attack on the smallest iron(II)
diimine complex, [Fe(gmi)3]2+, in water and in aqueous organic cosolvent
mixtures. The organic cosolvents used are methanol, ethanol,
i-propanol, t-butanol and acetone. Harked rate increases are observed
with increasing content of organic cosolvents at atmospheric pressure at
298.2K. These kinetic data, combined with transfer chemical potentials
for the complex and hydroxide ion in binary mixtures, enables us to
carry out analysis of initial state-transition state solvation in these
binary systems. Volumes of activation, ~V*, which are derived from
effects of pressure on the rate constant in the above mentioned systems
yield initial state-transition state solvation differences. The two
- 89 -
approaches when considered together for this and othe ti ' r reac ons of
Fe(II) diimine complexes, afford a reasonably complete solvation
interpetation(lO>.
5.2 EXPERIMENTAL
Kinetics of reaction of [Fe(gmi)3]2+ and [Fe(mmi)3]2+ complexes
with hydroxide ions were monitored in 0%, 20%, 40%, 60% and 80% by
volume of methanol. For EtOH, i-PrOH, t-BuOH and acetone 60% by volume
of organic cosolvent was the maximum used, as transfer chemical
potential values for the hydroxide anion are not available in these
systems for volumes higher than 60% of the organic cosolvent. The ionic
strength (1 = 0.33M) was maintained by adding appropriate amounts of
NaCl, and constant temperature 298K was maintained throughout. In all
systems the concentrations of the complex were 10-4 mol dm-3 while the
hydroxide ion concentration was much greater than that of the complex.
The reactions were monitored in solution at several different hydroxide
concentrations, where 5 x 10-3 ~ [NaOH] ~ 2 x 10-1 mol dm-3, as shown in
Table 5.1.
Rate constants were calculated from the dependence on time of
absorbance, at Xmax = 554 and 561 nm for the [Fe(gmi)3]2+ and
[Fe(mml)3]2+ complexes respectively(11), characterising the decrease in
concentration of the iron complex with time. No shift in Xmax was
observed when the complex was dissolved in alcohol-water mixtures or
during the reaction period. The dependence of absorbance on time was
followed using an HP8451A diode array spectrophotometer, for 2.5 half
lives for reaction of [Fe(gmi)3]2+ ions in solution. The absorbance
data was stored and analysed for the first-order rate constant. The
- 90 -
second-order rate constants were calculated, by computer fitting, from
the gradient of plots of first-order rate constants as a function of
iaOH concentration for each solvent mixture for each system,
5.3 RESULTS
Table 5.1 contains k(Ob.> values for hydroxide attack on Fe(II)
complexes in water and binary aqueous mixtures. In both, the aqueous
solutions and the aqueous solvent mixtures, the reaction of [Fe(gmi)3]2+
and (Fe(mmi)3]2+ with hydroxide ions were first order in (complex]. In
all cases the absorbance at infinity was close to zero indicating that
the reaction had gone to completion. The rate of reaction between
(Fe(gmi)3)2+ and OH- in aqueous solution is given by equation below,
-d[complex)/dt = (k1 + k2[OH-)[complex)
k, is the first-order rate constant for the aquation reaction, k2 is a
second-order rate constant. For solutions where hydroxide concentration
is much greater than the concentration of the complex, the observed
first-order rate constant is given by k(ob.> = k, + k2[OH-). The
dependence of k(ob.> on OH- concentration, as shown in Figure 5.1, was
linked to the above equation using linear least-square procedures.
Results for k2 terms are reported in Table 5.2 for methanol, Table 5,3
for ethanol and i-propanol and Table 5.4 for t-butanol and acetone, The
kl terms were found to be negligible and therefore not reported.
,5.4 DISCUSSION
A marked increase in the rate constant is observed for hydroxide
attack both on [Fe(gmi)3]2+ and [Fe(mnd)3)2+ complexes with increasing
- 91 -
TABLE 5.1
First-order rate constants, kCob.), for reaction between [Fe(gD1>31 2 +
and OH- ions in methanol-, i-propanol- and acetone-water mixtures; and between [Fe(mDd)31 2
+ and OH- ions in methanol-water mixtures at 298.2K; ionic strength = 0.33 mol dDl3
kCob.)/S-l for [laOH1/mol dur3
[Fe(gDi)3J2+ ------------------------------------------------------
0.02 0.015 0.01 0.0075 0.005
Volt J(eOH
0 1. 14x10-4- 8. 66x10-6 5.51x10- S 4. 65x10-s 2. 63x10-s
20 4. 66x10-4- 3. 58x10-4- 2.07x10-"" 1. 83xlO-"" 9.98x10-s
40 1.'l2x10-3 1. 25x10-3 8. 18x10-"" 6. 26x10-4- 4. 17x 1 0-""
60 5.'l4x10-3 4. 27xl0-3 2. 66x10-3 2.03x10-3 1. 35x10-3
80 1. 56x10-2 1. 13x10-2 '1. 46x10-3 6.03xl0-3 3.97x10-3
Volt i-PrOH
20 6.'lOx10-4- 4. 95x10-4- 3. 18x10-4- 2.4'lx10-"" 1. 59x10-""
40 2. 98x10-3 2. 34x10-3 1. 46x10-3 1. 15x10-3 6. 99x10-4-
60 7.'l5x10-3 5. 52x10-3 3. 31x10-3 2. 77x10-3 1. 60x10-3
Volt Acetone
20 8.40x10-4- 6. 18x10-4- 3. 95x10-"" 2. 97xl0-"" 1. 94x10-""
40 3. 48x10-3 2. 52x10-3 1.81x10-3 1. 32xl0-3 8. 58x10-""
60 1. 04x10-2 8. 28x10-3 5.'lOx10-3 4.40xl0-3 3.03x10-3
[ Fe (DIlIi ) ~1 2+
--------------------------------------------------------------------
Voll J(eOH
0 8. 87x10-6 6. 39x10-6 4. 41x10-6 3. 52xl0-6
20 2. 53x10-6 1. 82x10-s 1. 26x10-6 9. 32xl0-6 7. 83xl0-6
(0 6. 14x10-6 4. 96x10-s 2. 52x10-s 1.80x10-s
60 2. 82x10-4- 1. 92x10-4- 1.48%10-4- 9. 32xl0-s 7.38x10-s
80 8.61x10-3 6.03x10-3 3. 87x10-3 2. 16x10-3 1. 36x10-3
- 92 -
1.0
0.5 ~ , fJ)
...... " , Jl 0 v ~ '4 0 ~
20 (c) 40
10
1. 0 2.0
10"[ "nUII1/mol dm-3
?,oo (e)
100
o "'----__ -----t"---J 1.0 2.0
102 [R1l01lJ/mol dm-:3
FIGURE 5.1
Dependence of k(ob.) on ['aOll] for reaction between OH- and
IPe(gmi)31 2 + ions in water and methanol/water solvent mixtures
at 290.2 K: Voll "ethanol = (a)O, (b)20, (cHO, <d)60, (e)80.
- 93 -
\0 l:'-
TABLE 5.2
Initial state-transition state analysis for iron(II) d1i~ne complexes
hydroxide reaction fro. water to aqueous methanol, at 298.2 K.
[ Fe (g:ud. ) 3] 2+ [ Fe (lIlIIi ) 3] 2+
Vol ~ :le0H 0 20 40 60 80 0 20 40 60
~ 102 k2/I-l s -1 0.570 2.25 8.60 27.5 77.5 0.0444 0.125 0.312 1.350
om~G"/kJ 1101- 1 -3.33 -6.58 -9.40 -11. 92 -2.56 -4.83 -8.46
om~e{Cation}/kJ 1101-1 -0.70 -1.29 -0.22 +4.42 -2.90 -5.20 -7.50
om~e(OH-)/kJ 1101- 1 (a) -0.15 -0.05 +1.35 +5.75 -0.15 -0.05 +1.35
Initial state -0.85 -1.34 +1.13 +10.17 -3.05 -5.25 -6.15
Transition State -4.18 -7.92 -8.27 -1.75 -5.61 -10.08 -14.61
(a) from reference 12
80
3.580
-10.89
-6.80
+5.75
-1. 05
-11. 94
'-0 Vl
I
TABLE 5.3
Initial state-transition state analysis far hydroxide attack on [Fe{gDd>3]2+
reaction fro. water to aqueous ethanol and aqueous i-propanol, at 298.2 K.
Vol ~ EtOB Vol ~ i-PrOB
0 20 30 40 50 0 20 40 60
102 k2/I-1 s -1 0.5'1 2.60 6.38 11.15 15.50 0.5'1 3.30 14-.93 35.60
6 ... 6G*/kJ :.11- 1 -3.'16 -5.98 -'1.3'1 -8.18 -4-.35 -8.09 -10.24-
6 ... ~e{[Fe{gDd)3]2+}/kJ :.11- 1 -0.4-0 -0.95 -1.80 -3.10 -0.52 -3.'1'1 -3.81
6 ... ~e{OH-)/kJ :.11- 1 (a) +1.4-0 +2.45 +4-.15 +6.90 (b) +3.40 +'1.20 +11. 60
Initial State +1.00 +1.50 +2.35 +3.80 +2.88 +3.43 +'1.'19
Transition State -2.'16 -4.4-8 -5.02 -4.38 -1. 4'1 -4.66 -2.45
(a) as calculated in Chapter 4, (b) from reference 13
~ 0\
TABLE 5.4
Initial state-transition state analysis for hydroxide attack on [Fe(gDd)3]2+
reaction frOD water to aqueous t-butanol and aqueous acetone. at 298.2 K.
Vol I t-BuOH Vol "L Acetone
0 10 20 30 40 50 0 20 40 60
102k2/I-l s -1 0.5'1 1.165 2.64 5.10 6.9'1 10.'10 0.5'1 4.20 18.00 26.30
6rnAG* /kJ DOl- 1 -1.'1'1 -3.80 -5.42 -6.20 -'1.25 -4.95 -8.55 -9.49
6rnpe{[Fe(gDd>3]2+}/kJ mol- 1 -0.55 -1.43 -3.30 -5.50 -4.'10 -6.40 -14.14 -18.34
6rnpe(OH-)/kJ DDl- 1 (a) +0.65 +2.40 +5.60 -'1.'10 +9.50 (b) +6.80 +15.50
Initial State +0.10 +1.00 +2.30 +2.20 +4.80 +0.40 +1.36
Transition State -1. 6'1 -2.80 -3.12 -4.00 -2.45 -4.55 -'1.20
(a) from reference 14, (b) as calculated in Chapter 4
~ -..J
FIGURE 5.2
Initial state-transition state analysis of reactivity trends
f or hydroxide attack on [Fe (gmi) 3] 2+ complex from water to
aqueous .ethanol and aqueous iso-propanol, at 298.2 K.
I, +51 :""'I
0 ~ OH-.
1. S. ~ 20 40 [ Fe <gmi ):;oJ 2-
~ 0
"- Vol~ ]leOE
t eo( ... ~ ,
'0 (a)
-5
T.S.
-10 o ... ~G"
OB-
+10
loS.
~
I +5 r-{
0 g ., ~ 0 """I
"-
40
Vol~ i-PrOH
60
~ T.S.
~ ~ .. ~ ,~[Fe(gmi)3J2'"
-5 (b)
-10 o ... ~G"
\.0 co
I -4.66 +11. 6 OH-
-1.92 -8.27 -----T.8. 85.8 kJ DDl- 1 85. 8 k.I lIKJl- 1
1
o
(a)
+1. 35 OH-
-0.15 -0.05./ ~ [Fe <gmi )3] 2. (b)
-0.52
---3.81
-3.77 [Fe(gmi)3]~+
20 40 60 o 20 40 60
Vo1~ ){eOR Vol1.. i-PrOH
FIGURE 5.3
I Step-diagraJI' showing in! tial state-transition state analysis of reacti vi ty
trends for hydroxide attack on [Fe(gm)3]2. cODplex froD water to aqueous
.ethanol and aqueous iso-propanol; at 298.2 K, kJ IKll- ' .
ou·-
·HO
,. I r1 0 l.s. a
0 Vol?; Acetone
~ 40 60
~ "-
~ E
<0
T. S.
-10 ~ 6' ... 6Gft
-20
FIGURB 5.4
Initial state-transition state analysis of reactivity trends for hydro~tde
attack on lFe(gml)31 2 + complex frOD water to aqueous acetone, at 290.2 R.
- 99 -
-4.55
85.8 kJ 1ID1- 1
1
+0.40
o 20
-7.20 --- 'f.B.
-t15.50 011-
-t1.36 ~--- I. S.
40
FIGURH 5.5
-18.34 [ Fe <gmt) ",I ;0 •
60
Vol~ Acetone
'Step-diagram' showing initial state-transition state analysis of
reactivity trends for hydroxide attack on [Fe(gmi)3]2+ complex
from wnter to aqueous ncetone; at 298.2 I, kJ DOl-'.
- 100 -
85.8 kJ R11- 1
o ,
- - -
o = MeOn • = Blon ~ = i-PrOB A = t-BuoOB
- --0- ---0.1
Transition State
Initiai State
-- --0- .....
0.2 0.3 0.4
•. f. cosolvent
FIGURE 5.6
'Step-diagra.' showing initial state-transition state analysis uf
reactivity trends for hydroxide attack on [Pe(gmi)31 2• complex
from water to aqueous cosolvents; at 298.2 K, kJ JllJl- 1•
- 101 -
content of methanol. The same increase in the rate constant is also
observed for hydroxide attack on [Fe(gmi)3]2+ complex with the increase
in content of other alcohols. The details of initial state-transition
state analysis for hydroxide attack on these complexes is set out in
Tables 5.2-4 for various cosolvent systems. As initial state-transition
state analysis for the [Fe(mmi)3]2+ complex is very similar to that of
the [Fe(gmi)3]2+ complex, the latter is being discussed first while
hydroxide attack on [Fe(mnd)3]2+ complex will be discussed later. The
results for initial state-transition state analysis are shown in Figures
5.2-3(a) and 5.2-3(b) for methanol and i-propanol cosolvents
respectively. Figures 5.2-3(a), for methanol/water solvent mixtures,
show [Fe(gmi)3]2+ cation to be only slightly preferentially solvated by
methanol in the water-rich region. On the other hand the hydroxide ion
shows no preferential solvation by methanol as its transfer chemical
potential is zero up ~o 40% methanol(12), destabilisation of this anion
only occurs in the high methanol region (which is due to strong
hydrophilic hydration, ~.water molecules around the hydroxide ion are
negligibly displaced up to 40% methanol/water mixture). The initial
state shows only small stabilisation while the transition state is
stabilised much more than the corresponding initial states by methanol.
As mentioned in the previous chapter, [Fe(gmi)3]2+ cation is
progresively stabilised in order MeOH < EtOH < i-PrOH < t-BuOH up to at
least 40% of the organic cosolvent. The hydrophilic hydroxide ion is
destabilised in the same order for the same cosolvents. Figures 5.2-
3(b) show initial state-transition state analysis for the i-propanol
cosolvent. The much higher destabilisation of the hydroxide ion(13)
than the complex cation leads to destabilisation of the initial state.
- 102 -
Once again the transition state is stabilised on transfer from water to
i-PrOH. Unlike the methanol system, where reactivity is dominated
entirely by stabilisation of the transition state, in the i-PrOH system
the reactivity trend may be attributed to destabilisation of initial
state and stabilisation of transition state.
Equal and opposite preferential solvation of the hydroxide ion and
[Fe(gmi)3]2+ complex by acetone (the [Fe(gmi)3]2+ cation is greatly
stabilised while the hydroxide ion is destabilised on transfer from
water to acetone) leaves the initial state relatively constant on
transfer from water to acetone. However, the corresponding transition
state is stabilised indicating that contribution to the reactivity is
dominated by stabilisation of transition state, as shown in Figures 5.4-5.
It is of particular interest to note the trends in initial state
transition state analy~is for four alcohol water systems as shown in
Figure 5.6. Although the transition state is stabilised for all
systems, the initial state is increasingly destabilised in the order
XeOH < EtOH < i-PrOH < t-BuOH, which is the same as the trend for
hydroxide ion destabilisation on transfer in these systems. From this
it may be realised that hydroxide preferential solvation in these
casal vent systems plays an important part in determining the
reactivity.
The higher stabilisation of the transition state than the
corresponding initial state may be attributed to two main factors, one
would have a destabilising effect while the other a stabilising effect
on transition state on transfer from water into organic cosolvent: (a)
charge reduction from 2+ plus 1- (initial state) to 1+ (transition
- 103 -
state) which has a stabilising effect on transition t t s a e as we go from
water into the organic cosolvent mixture (increase in solvation by
methanol). This charge effect becomes less influential as the size of
the complex cation increases due to the decrease in charge density.
Evidence for the preferential solvation of complexes with a lower charge
is provided by ternary complexes (see Chapter 7), [Fe(bipY)2(CN)21 and
[Fe(bipY)2(CN)2]+, where the former is more preferentially solvated by
methanal; (b) the second factor is the decrease in hydrophobicity of the
cation as hydroxide ian is incorporated into the complex periphery.
This may have a destabilising influence an transition state on going
from water into the organic cosolvent (decrease in solvation). The
decrease in hydrophobicity increases with increase in the ligand size of
the cation. Evidence that the introduction of a hydrophilic ligand into
the complex periphery causes destabilisation is provided by already
mentioned examples in Chapter 4. Further evidence for this is also
provided by examples in Chapter 7 where, on transfer from water into
methanal, the three analogues are stabilised in order [Fe(bipY)31 2+ )
The above mentioned factors may be of relevance when considering
the differences in the increase of the second-order rate constant, for
hydroxide attack on various Fe(II) diimine complexes, with an increase
in organic casal vent. If so, these factors would have direct relevance
and hence offer an explanation for the difference in stabilisation of
transition state to corresponding initial state for various Fe(II)
diimine cation complexes. The increase in k2 an transfer from water to
methanal is reflected in negative values of OM~G·. Taking several
- 104 -
examples in methanol/water mixtures, i e tris-gmi mIni h (16) d . . " p en an
bsbMe2(6) Fe(II) complexes, which are in increasing order of ligand
size, their OM~G* values are found to decrease in order of ligand size
as shown in Figure 5.7.
On going from initial state to transition state incorporation of
the hydrophilic hydroxide ion into the predominantly hydrophilic
periphery of (Fe(gmi)3]2+ complex would produce hardly any reduction in
hydrophobicity (the [Fe(gmi)3]2+ complex also shows preferential
solvation by water), thus there is no destabilising influence on
transition state on going into methanol. However as the (Fe(gmi)3]2+
complex is a small cation, where there is a high charge density, charge
reduction would have stabilising effect on transition state as we go
into the less polar medium (i.e. into methanol solvent). In the case of
the (Fe(mmi)3]2+ complex, which is more stabilised by methanol than its
gmi analogue, stabilisation would be increased further in transition
state as a result of charge reduction (charge density smaller on
[Fe(mmi)3J2+ complex). However this increased stabilisation of
transition state may be opposed by decrease in hydrophobicity of the
complex periphery due to incorporation of hydroxide ion. As a result
this would cause methanol desolvation. The same would hold true for
hydroxide attack on Fe(II) complexes of phen and bsbMe2 which are much
more preferentially solvated by methanol than the previous two
complexes. The larger size indicates smaller charge density, therefore
the charge reduction would have an even smaller effect than in the
previous two cases on the transition state stabilisation but greater and
opposing effect due to reduction in hydrophobicity.
- 105 -
-t4
,... .. , r-f
i 20 40 60 80 100
~ 0 v Vol' J(eOIl ..., " t .., ~ " " i ..,
-4 ~ ~ ..... ~ I phen
" • 1 to
-8
-12 gill
FIGURB 5.7 ,
VariatIon of the ratio of the second-order rate constants for the reactlon
of hydroxide attack on [Fe(LL)31 2 + complexes in aqueous methanol, at 290.2 K
- 106 -
The above reasoning and the observed pattern in om6G*, for these
complexes, suggest that the reactivity of Fe(II) diimine complexes with
the hydroxide ion is strongly influenced by the nature of the ligand
periphery. From the high pressure study in the next section, the
activation volumes for hydroxide attack on [Fe(gmi)3]2+ are analysed for
the differences in solvation between initial and transition state in
several binary systems. The pattern of 6V* in different cosolvents for
this reaction, and other reactions in the same cosolvent will be
discussed to confirm the above results, Le. that the reactivity of Fe(II)
diimine complexes with the hydroxide ion is strongly influenced by the
nature of both the ligand periphery and the cosolvent.
- 107 -
5.5 HIGH PRESSURE KINETICS
The effect of pressure on a chemical reaction in solution is
attributed to the volume change which occurs in the activation step of
that reaction. That is, if there is a decrease in the partial molar
volumes on going from reactants to an activated complex, pressure will
drive the activation equilibrium in favour of products of smaller volume
and the reaction will be accelerated. Thus if a reaction is accelerated
by an increase in pressure this indicates a negative volume of
activation, ~V· ,{p~V* = -RTln(kp/ko) equation 5.2 }(16'; i.e. decrease
in partial molar volumes from the initial state to the transition state.
The converse of this situation is also true.
Pressures that provide useful changes in rate constants, hence ~V*,
are measured in kilobars. Typical volumes of activation lie in the
range(17) +35 to -60 cm3 mol- 1 and are associated with definite
structural changes, the most important of which are:-
(a) Association of two reactant molecules to form a product. This is
accompanied by reduction in volume, a negative contribution to ~V*.
Conversly, dissociation of a molecule into fragments brings about an
increase in volume change, +~V·.
(b) An increase in solvation, i.e. association with the solvent,
causes a reduction in volume which is particularly large when an ionic
charge is created. Ionic reactions or those reactions with strongly
dipolar transition states are accompanied by very negative ~V* which are
strongly solvent-dependent and are more negative the less polar the
solvent. This is because the influence of the electric field, which
brings about the solvent structuring in the vicinity of the ion, extends
to a greater distance from the charge centre the lower the dielectric
- 108 -
constant of the medium. Therefore, this 'electrostrictive effect' for
which ~V·_ is the volume change associated with electrostriction and
~V·i, the intrinsic volume of solute, are related by, ~V· = ~V.i + ~V._.
Thus the increase in charge in transition state predicts negative ~V •.
The reverse of this, the relaxation of solvation that accompanies charge
neutralisation brings about an increase in activation volume, +~V •.
Therefore the inorganic reactions that give a transition state less
dipolar than the initial state are retarded by pressure, and the volume
of activation is positive.
In this section, reaction kinetics at elevated pressures are
followed for hydroxide attack on [Fe(gmi)3]2+ complex in aqueous
solution and in binary mixtures for several cosolvents. Rate constants,
for atmospheric and elevated pressures, give positive volumes of
activation for the whole range of all cosolvent systems indicating a
high desolvation of the transition state.
5.6 EXPERIMENTAL
Kinetics of dissociation at elevated pressure were studied by using
the high pressure apparatus described in Chapter 2. A typical run
involved making up a solution (approximately 150 cm3 ) of the appropriate
composition, except for the dissociating agent, with cosolvent added by
suitable volume for the final desired value. Upon (reaction) initiation
by adding aqueous hydroxide solution, the reaction solution was divided,
with the major part (~ 120 cm3 ) used in the high pressure apparatus
while some of the remaining solution was used in a cell for the
atmospheric pressure comparison run. Both solutions were maintained at
298.2K; aliquots were withdrawn from the high pressure system and
- 109 -
monitored in the SP 8-100 spectrophotometer. The concentration of the
complex was in the range of 9 x 10-& and 3 x 10-4 mol dm- 3 and hydroxide
ion concentration was varied as necessary, with the ionic strength made
up to 0.33 mol dm-3 with sodium chloride as appropriate.
The solvent composition ranges and hydroxide ion concentrations
used in high pressure runs reflected the practical limitations. In
methanol solutions approaching 100% by volume methanol the dissociation
of the gmi complex is sufficiently rapid that the time resolution acts
as the limit. This time limit is also approached as the EtOH, i-PrOH
and t-BuOH cosolvents reach volumes of approximately 60%.
5.7 RESULTS
The kinetic results reported in Table 5.6 show the variation in
first-order rate constants with hydroxide concentration and solvent
composition. The val~es of rate constants at atmospheric and at
elevated pressures, ko and kp respectively, were obtained from slopes
for plots of In<abs) vs time. The reactions were initially followed for
three half-lives, Table 5.5(a). However such plots showed linearity
only for the first half-life as shown in Figure 5.8(a) for water.
Therefore most of the reactions were followed for the first half-life
only as shown in Table 5.5(b) and Figure 5.8(b) for 20% i-PrOH/water
solvent mixture at atmospheric and 344 bars pressure.
Volumes of activation were calculated using the slopes from plots
of In(kplko) vs pressure using equation 5.2. Plots of In(kplko) vs
pressure showed no systematic deviation from linearity as seen in Figure
5.9 for water and up to 80% methanol/water solvent mixtures. Volume of
- 110 -
TABLE 5.5
Decrease in absorbance vs time for hydroxide attack on lFe(gmi)31 2 +
complex at atmospheric and elevated pressure in water(a) and 20~ i-propanol-water solvent Dixture (b), at 298.2 K; I = 0.33 mol dDr 3 •
(a)
[Fe(gmi)31 2+ + hydroxide(0.02X) in water
p = atDKJspheric p = 689 bars
Abs In(Abs) Aba In(Abs) tille/min ----------------------------------------------------------------------
1. 016 0.910 0.819 0.745 0.689 0.638 0.581 0.512 0.447 0.375 0.345 0.285 0.253
(b)
0.016 -0.094 -0.200 -0.295 -0.372 -0.450 -0.543 -0.670 -0.805 -0.982 -1.064 -1. 255 -1. 374
1. 017 0.940 0.882 0.828 0.787 0.754 0.708 0.655 0.600 0.537 0.505 0.430 0.364
0.017 -0.062 -0.126 -0.189 -0.240 -0.283 -0.346 -0.423 -0.510 -0.621 -0.684 -0.844 -1.010
o 20 40 60 74 90
109 135 162 200 223 287 355
lFe(gDi)312+ + hydroxide(O.OlX) in 20~ by volume of iso-Prop~~~:_ -----------------------------------------------------------------
p = atDllspheric P = 344 bars -----------------------------------
Abs In(Abs) Aba In(Abs) ti~/min
----------------------------------------------------------------------
0.823 -0.192 0.816 -0.203 2
0.783 -0.245 0.782 -0.246 5 7.5
0.757 -0.278 10 0.727 -0.322 0.735 -0.308
13 0.694 -0.365
-0.390 16 0.658 -0.418 0.677
20 0.622 -0.475
-0.494 24 -0.536 0.610 0.585 28
0.551 -0.596 -0.598 32
0.519 -0.656 0.550 37
0.484 -0.726 -0.715 42
0.451 -0.796 0.489 -0.794 49
0.410 -0.892 0.452 -0.870 56
0.376 -0.978 0.419 0.402 -0.911 60
- 111 -
o ~~ ________ ~10~0 __________ ~2~0~0 __________ ~3~O~0~ ____ __
time/min
-0.8
•
• (a)
-1.2
o
o
40 20 60 o ti~/min
-0.2
-0.6
-1. 0
FIGURE 5.8
rIot of In(Abs) vs time for hydroxide attack on (Fe(gml)3]2+ cO~llex at atmospheric (0) and elevated pressure (.) in water (a) aod 20t i-propanol/water mixture (b). For details see Table 5.5.
112 -
P/bar 689 1030\
o
-0.2
" 0 ~
" a -0." ~ v ~ M -0.6
401.
20~
-0.0 O~
FIGURE 5.9
Plot of logarithm of ratio of (kp/ko) vs pressure for hydroxide attack on [Fe(gai)31 2 + couplex, at 298.2 K, at various compositions of aqueous methanol.
- 113 -
,TABLE 5.6 First-order rate constants, k<Ob.>, for hydroxide attack on lFe(gmi)3]2- complex cation, at atmospheric and at elevated pressures, and ratios In(kplko) in water and in methanolethanol-, i-propanol-, t-butanol- and acetone-water solve~t mixtures; and derived volumes of activation, (~V.), at 298.21
lethanoll 1 I 1 Voll IlOH-]1 pI 104 ko/s- 1 1 104kp/s-l I In(kp/ko) 1 ~V.I m. f. 1 DOldI-3 1 bars 1 1 1 1 cr.,l-l
---------1-------1------1------------1------------1-----____ 1 _______ _ 01 0.04 344 1.820 1.402 -0.265
0.00 0.04 689 1.812 1.140 -0.460 +16.7 0.02 689 0.834 0.526 -0.468 0.04 1034 1.785 0.862 -0.725
---------1-------1------1------------1------------1---------1--------20~ 0.02 344 2.849 2.233 -0.242
0.100 0.02 689 2.850 1.848 -0.433 +16.2 0.02 1034 2.845 1.354 -0.740
---------1-------1------1------------1------------1---------1--------0.01 344 6.700 5.420 -0.212
40~ 0.01 689 6.710 4.350 -0.433 0.229 0.01 1034 6.602 3.350 -0.678 +15.5
0.01 1206 6.810 2.960 -0.833 0.01 1378 6.540 2.550 -0.942
---------1-------1------1------------1------------1---------1--------50~ 0.005 689 5.680 4.050 -0.338
0.308 0.005 1034 5.620 3.320 -0.526 +12.6 ---------1-------1------1------------1------------1---------1--------
60~ 0.002 344 4.690 4.260 -0.096 0.400 0.002 689 4.710 3.760 -0.225 +8.1
0.002 1034 4.640 3.240 -0.230 ---------1-------1------1------------1------------1---------1--------
80~ 0.002 344 19.80 18.80 -0.052 0.640 0.002 689 21.66 18.90 -0.136 +4.8
0.002 1034 19.20 15.25 -0.230 ------_1------ ________________________ _
EtOH 1 Voll 1 11. f. 1 1
---------1-------1------1------------1------------1--------- --------30~ 0.005 344 2.520 2.080 -0.192
0.117 0.005 689 2.520 1.750 -0.365 +14.4 0.005 1034 2.550 1.350 -0.636
---------1-------1------1------------1------------1---------1--------40~ 0.005 344 5.980 5.230 -0.134
0.170 0.0025 689 2.300 1.780 -0.256 +10.4 0.0025 1034 2.300 1.480 -0.441 _______ _
---------1-------1------1------------1------------1---------1 4 780 -0.015 50~ 0 0025 344 4.820 .
. 1 680 -0.085 0.235 0.0015 689 1.830 . 0.0015 1034 1.830 1.490 -0.210
+4.3
________ ----_1----- _________ _
TABLE 5.6 (Continued) ..... .
i PrOH 1 1 1 1 Vol~ I[OH-] I pI 1 10 .... ko/s- 1 10 .... kp/s-l Iln(kp/ko>l flV"1 1I.f. IlIOldr<31 bars 1 1 1 Ic.- 3 DlJl- 1
---------1-------1------1------------1------------1----_____ 1 ________ _ 20~ 0.010 344 2.480 2.040 -0.195
0.056 0.010 689 2.510 1.650 -0.419 +15.3 0.010 1034 2.530 1.300 -0.666
---------1-------1------1------------1------------1---------1 ________ _ 30~ 0.010 344 5.910 5.080 -0.152
0.092 0.005 689 2.520 1.730 -0.376 +13.7 0.010 1034 6.020 3.280 -0.605
---------1-------1------1------------1------------1---------1---------40~ 0.005 344 5.310 4.770 -0.110
0.136 0.005 689 5.320 4.070 -0.268 +9.7 0.005 1034 5.320 3.520 -0.410
---------1-------1------1------------1------------1---------1---------50~ 0.003 689 2.260 1.850 -0.203 +7.6
0.191 0.003 1034 2.230 1.680 -0.326 ---------1-------1------1------------1------------1---------1---------
60~ 0.003 689 3.940 3.400 -0.147 +5.8 0.262 0.002 1034 2.580 1.990 -0.260
t-BuOH Vol~
11. f. 1 ---------1-------1------1------------1------------1---------1---------
10~ 0.020 344 1.872 1.525 -0.205 0.021 0.020 689 1.838 1.215 -0.414 +14.9
0.020 1034 1.852 1.050 -0.567 ---------1-------1------1------------1------------1---------1---------
17~ 0.020 344 3.630 3.140 -0.145 0.040 0.020 689 3.580 2.530 -0.344 +12.2
0.010 1034 1.735 1.035 -0.519 ---------1-------1------1------------1------------1---------1---------
30~ 0.010 344 3.950 3.640 -0.082 0.076 0.010 689 4.030 3.286 -0.204 +7.4
0.010 1034 4.100 2.930 -0.336 ---------1-------1------1------------1------------1---------1---------
40~ 0.010 689 5.500 5.370 -0.024 0.113
---------1-------1------1------------1------------1---------1---------501 0.005 689 4.125 3.670 -0.117 +3.9
0.161 0.005 1034 4.250 3.180 -0.290
Acetone
20~
0.058
0.010 0.010 0.010
344 689
1034
2.647 2.680 2.631
- 115 -
2.286 1.969 1.666
0.146 -0.308 -0.457
+11. 1
activation values for water and aqueous cosolvent systems are reported,
together with other data, in Table 5.6.
5.8 DISCUSSION
A significant decrease in the rate constant is observed for
hydroxide attack on [Fe<gmi)3]2+ complex with an increase in pressure,
resulting in positive 6V· values. In aqueous solution the volume of
activation for this reaction has a value of +16.7 cm3 mol- 1 , as shown in
Table 5.6 and Figure 5.10 <uncertainties for 6V· values, in Table 5.6,
are between +0.6 and +1.5 cm3 mol- 1). Nucleophilic attack by hydroxide
in bimolecular mode is a reasonable conclusion from the reaction
kinetics. The expected value of the volume of activation for
bimolecular reaction based on the intrinsic volume change is about -10
cr03 mol- 1 (18). The difference in observed and expected 6V· values, of
the order of +25 to +35 cm3 mol- 1 , implies a very large solvation
5.8.1 Reaction in water
The large positive 6V· value for hydroxide attack on gmi complex
may be attributed to decrease in solvation of the species involved on
going from initial state to transition state. Initially both the
complex cation and the hydroxide ion are specifically hydrated but to a
different extent because of the difference in charge densities and
nature of periphery. On reaction with the iron complex whether its
point of attack is the iron nitrogen bond or initially at some other
position on the ligand, the negatively charged hydroxide ion will lose
its specific hydration water as it becomes engulfed in the body of a
- 116 -
complex ion. This loss of specific hydration water by hydroxide ion
will be accompanied by increase in 6V*. Further increase in 6V. will be
brought about by relaxation of the electrostricted solvent which
accompanies a decrease in charge, as the net positive charge of the
resultant species is one (1). The higher 6V· value for hydroxide attack
in aqueous solution on [Fe(gmi)3]2+ to that on [Fe(hsbh)]2+ complexeS)
may be explained in terms of difference in electrostriction. The latter
complex is slightly larger and more hydrophobic therefore of a lower
charge density, hence less solvent structuring which accompanies an
increase in 6V*. This is further confirmed by a much larger and totally
hydrophobic [Fe(bsb-Me2)3]2+ complexeS) whose 6V· is in Table 5.7.
However, this trend is not followed by 6V· values for hydroxide attack
on phen and bipy Fe(II) complexes (+19.7 and +21.6y cr03 mol- 1
respectively(20», which are much higher than that for the [Fe(gmi)3]2+
complex.
Complex [Fe(hxsbh)]2+
16.7 13.4 11.1
TABLE 5.7 6V. for Fe(II) diimine complexes + OH- reactions in water
5.8.2 Reactions in Binary Systems
On going from water into methanol cosolvent the volume of
activation, for hydroxide attack on [Fe(gmi)2]2+ reaction, remains
relatively constant ( ~ +16 cr03 mol- 1 ) up to 40% by volume of methanol.
- 117 -
I"
r-I 0 S
~i S U "-.... ~ co ~ <J
+20
+16
+12~ t \\ "- o = 1IeOB
" • = EtOB ~ = i-PrOB
\ \ • " • = t-BuOB
+8
+4
m.f. Cosolvent O~I--------------__ ------------__ --------------~--------------~--
0.2 0.4 0.6 0.8
FIGURE 5.10
Plot of 6V· VB BOle fraction of aqueous cosolvents for hydroxide attack on [Fe(s-d>3)2- co.plex. at 298.2 K.
+30
+20
I'"
I ri
~ ~~ flO
0 "-# :> <J
~
d 0
= lFe(bsb-Xe2)31 2+ <.>
= [Fe(hxsbh)12+ c.>
= [Fe(gmi)3]2+
Vol~ MeOH o r---------~~--------~----------_r----------~------
20 40 60 80
FIGURE 5.11
Plot of 6V· vs Vol~ of aqueous methanol for hydroxide attack on Fe(II) diiDine complexes, at 298.2 K. (a) Data froD reference 6.
- 119 -
However, the decrease in 6V* is rapid in higher methanol region, ie +8
and +5 cm3 mol- 1 in 60 and 80% by volume (or 0.40 and 0.64 mole
fraction) of methanol respectively. This decrease in volume of
activation is also observed in other alcohols and takes place at
progressively lower mole fraction in order MeOH ) EtOH ) i-PrOH > t-BuOH
as shown in Figure 5.10. This trend in 6V* for our reaction in these
cosolvent systems may be explained by considering the following. As
[Fe(gmi)3]2~ cation exhibits some preferential solvation by methanol,
the magnitude and sign of 6V* will be determined by: (a) decrease in
charge resulting in higher solvation by organic cosolvent in transition
state which will give rise to electrostriction, therefore a decrease in
6V*i and (b) incorporation of hydrophilic hydroxide ion into the
periphery of the complex leading to decrease in hydrophobic character
which will in effect cause desolvation of organic cosolvent in
transition state - hence an increase in 6V*. The influence of the above
two factors on the magnitude of 6V*, although opposite, will vary in
dominance with the ligand size of the complex cation as mentioned
earlier in, this chapter. The former predominates in complexes
containing small ligands while the opposite is true for the latter.
Further, it is assumed that an increase in 6V*, due to release of
hydroxide hydration water, is constant throughout the binary
composition.
In a small and mainly hydrophilic [Fe(gmi)3]2~ the positive 6V* is
mainly due to the desolvation of the hydroxide ion on going from water
into the methanol cosolv~nt The decrease in the AV* value in high
h high i nfluence that decrease in the charge methanol region reflects t e
in transit ion state has on methanol solvent density on cation
- 120 -
structuring, <electrostriction of methanol). Further, it is of
particular interest to note the decrease in 6V· appears in the mole
fraction region where water structure-breaking by methanol is at the
maximum. The same is true for other water-alcohol binary mixtures where
the decrease in 6V· is at progressively lower mole fraction from
methanol to t-butanol.
In the above discussion we have seen that very small, if any,
influence in 6V· may be attributed to a change in hydrophobicity in the
transition state due to hydroxide incorporation in [Fe(gmi)a]2+ complex
periphery. However for a large and mostly hydrophobic [Fe(bsb-He2)a]2+
complex cation, which is preferentially solvated by methanol, the charge
density is minimal while the decrease in hydrophobicity of the complex.
due to hydroxide incorporation into the complex periphery, is
pronounced. This is reflected in the increase in 6V· on going from
water into the methanol region, Figure 5.11, which is a result of
methanol desolvation in transition state(6). It would be of interest to
determine 6V· values for hydroxide attack on [Fe(phen)a]2+ and
£Fe(bipy)a]2+ complexes in methanol/water cosolvent mixtures. From the
above reasoning it would be expected that 6V· would remain constant
throughout the whole composition range or may show a small increase in
6V. in the higher methanol region. These expectations may be justified
as in these two complexes desolvation of methanol may be equally opposed
by the charge density decrease in transition state.
- 121 -
REFERENCES
1. J. Burgess, Inorg. React. Kech., 7(1981>163j 7(1981>232.
2. M. J. Blandamer, J. Burgess, T. Digman, P. P. Duce, J. P. MaCann R. H. Reynolds and D. K. Sweeney, Transition Met. Chem., 8(1983)148
3. F. M. Mikhail, P. Askalani, J. Burgess and R. Sherry, Transition Met. Chern., 6(1981)51.
4. M. J. Blandamer, J. Burgess, R. I. Hains, F. M. Mekhail and P. Askalani, 1. Chern. Soc., Dalton Trans., (978)1001
5. J. Burgess and C. D. Hubbard, J. Chern. Soc., Chern. Commun., (1983) 1482.
6. J. Burgess and C. D. Hubbard, J. Am. Chern. Soc., 106(1984)1717.
7. D. W. Margerum and L. P. Morgenthaler, J. Am. Chern. Soc., 84(1962)706
8.
9.
10.
11.
12.
13.
14.
15.
16.
J. Burgess, Inorg. Chim. Acta, 5(1971)133
R. D. Gillard, D. W. Knight and P. A. Williams, Transition Met. Chern., 5(1980)321.
M. J. Blandamer and J. Burgess, Pure and Appl. Chern., 55(1983)55.
P. Krumholz, O. A. Serra and M. A. De Paoli, Inorg. Chim. Acta, 15(1975)25.
M. H. Abraham, T. Hill, H. C. Ling, R. A. Schulz and R. A. C. Watt, J. Chern. Soc., Faraday Trans., 80(1984)489
B. Clark, Ph.D. Thesis, University of Leicester, 1985.
J. Burgess, Unpublished Work.
M. J. Blandamer, J. Burgess and D. L. Roberts, J. Chern. Soc., Dalton Trans., <1978>1086.
D. R. Stranks, Pure Appl. Chern., 38(1974)303. 1981 N. S. Isaacs, "Liquid Phase High Pressure Chemistry", Wiley, .
17 D P 1 d H K 1 Coord. Chern. Rev., 36(1981)89. . . a mer an . e m,
18. T. Asano and W. L. Noble, J. Chern. Rev. 87(1978)407.
19. and R. Sherry, J. Chem. Soc., Chem. Com., J. Burgess, A. J. Duffild (1980)350
20. A L D R Stranks and S. Suvachittanont, lnorg. Chem., G. . awrance, . . 18(1979)82
-122-
CHAPTER
6
Solvation and Reactivity of
Semiaromatic Fe(II) Diimine
Complexes
6.1 INTRODUCTION
The semiaromatic bidentate diimine ligand, with the general formula
(I), provides the connection between the aromatic(bipy) and
aliphatic(gmi or bmi) diimine ligands. A whole series of iron(II)
complexes with semiaromatic diimine ligands has been prepared starting
from 2-pyridine aldehyde, 2-acetylpyridine and 2-benzoylpyridine with
ammonia, methylamine and phenylamine(l,2). An unsuccessful attempt by
Krumholz(3) to prepare an iron(II) complex of 2-pyridinalimine, pami,
from 2-pyridine aldehyde and ammonia resulted in preparation of a
terdentate. Further, the iron(II) complexes with the terdentate
semiaromatic ligand, of general formula (II), derived from condensation
of 2,6-diacetylpyridine with an amine(4-6),are also known.
R R' R'
~ H H pami H tsbh H 1{e hpmi .Me tsbJle I
I ]{e H mpami -R Ph tsbPh J-R ]{e ]{e mpmi Ph H ppami Ph ](e ppmi
(I) (11)
These semiaromatic diimines are of interest because not only can
the groups bound to the exocyclic C atom be altered with ease, as is the
case for the aliphatic diimine ligands, but also the groups bound to N
atom within the imine molecule can be altered with equal ease. The
latter makes such complexes ideal for studying factors involved in
determining stabilisation and destabilisation properties of Fe(II)
complexes containing these ligands. Of particular interest are those
complexes with the proton on the imine nitrogen. In this chapter,
- 123 -
preparation of the [Fe(pami)3]2+ complex is reported it as s perchlorate
salt. Solubility measurements and derived transfer parameters, in
methanol-water solvent mixtures, for the above and related bidentate and
terdentate diimine complex cations are reported and compared with those
of their analogues, R = Me. Although complexes with bidentate ligands
may be composed of a mixture of two isomers(7), we have not been able to
detect this from the kinetic results of hydroxide attack on the
complexes (ie we have seen no two stage reactions). The solvent effect
on the rate constant for hydroxide attack is the opposite for these
iron(II) complexes containing semiaromatic ligands when compared with
those containing symmetric ligands such as phen or gmi in the same
media. The solvation contribution to the initial state-transition state
is analysed for hydroxide attack on several of these complexes in
methanol/water solvent mixtures.
6.2 EXPERIMENTAL
Preparation of terdentate,[Fe(tsb)2]2+, complexes was carried out
by following published methods by condensation of 2.6-diacetylpyridine
with an appropriate amine in aqueous ethanol followed by addition of 0.5
mole equivalent of FeCl 2 in solvent(6). The same procedures were used
for preparing Fe(II) bidentate Schiff bases from 2-pyridine aldehyde,
2-acetylpyridine and 2-benzoylpyridine with ammonia and methylamine.
which were precipitated as perchlorate salts for solubility purposes.
Some of these complexes were also precipitated as hexafluorophosphate
salts for 'H-nmr spectra analysis.
The iron(II) complex containing the smallest ligand (paml) of this
bidentate semiaromatic series was prepared in nitrogen atmosphere where
- 124 -
to one equivalent of 2-pyridine aldehyde in ethanol solution (at 30-
35·C) 1.1 equivalent of 0.880 ammonia solution was added dropwise. The
solution was kept at the same temperature for 30 minutes with occasional
stirring. An aqueous solution containing 0.3 equivalent of FeC1 2 was
added gradually, accompanied by stirring. After a period of one hour
nitrogen was bubbled through the solution and the complex precipitated
as the perchlorate salt, and separated by filtration. Recrystallisation
was from the minimum amount of 50% ethanol/water solution. Each complex
was characterised from its ~max and extinction coefficient, whereas for
the newly prepared complex,[Fe(pami)3J2+, microanalysis was carried out
for C, H, and H, the results of which are reported below.
C H H
Calculated 37.63 14.63 3.14
Actual 39.93 15.37 4.02
Solubilities of the perchlorate salts of all the above mentioned
complexes were determined as described previously in Chapters 2 and 4.
Kinetics of hydroxide attack on these complexes were followed using the
HP8451A diode array spectrophotometer, as explained in Chapter 5, while
the kinetics of reactions for the smallest complex, [Fe(pami)3J 2+, were
followed by the "SFA-11 Fast Kinetics Accessory" unit which was fitted
on the SP8-l00 spectrophotometer (see Chapter 2).
6.3 RESULTS AND DISCUSSION
The absorption spectrum for the [Fe(pami)3J2+ complex cation is
shown in Figure 6.1 together with the absorption spectrum of the complex
d (3) That the 1 f i il r preparatory proce ures . obtained by Krumho z rom sma
- 125 -
:'-)
C'\
~ -II
400
--/ , / \
/ \ / \
/ \ // \
/ \ I
/ --------.... \ / ~ ~ \ ",-; / /, /
/, / I , / , /
" " ...... _-
------
500 )./D.D.
FIGURE 6.1
Spectra of the [Fe(paIi.)3]2+ ccmplex( ) £574 = 6800 ll-lCm- 1t and that
obtained by Krumholz from. reference 3, (- ---) £579 = 10500 )(-lClI.- 1 •
\ \ \ \ \ \ \
600
products are two different complexes is evident from diff t eren X~ax and
extinction coefficient values obtained. The complex is found to be
labile and undergoes substitution in the presence of phen or bipy. The
substitution reaction,
[Fe(pami)31 2+ + 3(phen/bipy) = [Fe(phen/bipY)31 2+ + 3pami
is multistage, resulting in formation of Fe(II) tris-phen or tris-bipy,
the first stage of which follows first-order kinetics with either phen
or bipy.
6.3.1 Transfer chemical potentials
Solubility measurements, and derived transfer chemical potentials
for salts and cations, are presented in Tables 6.1 and 6.2 for iron(II)
terdentate and bidentate complexes respectively. The plot of Om~e
versus methanol composition, Figure 6.2 for terdentate Fe(II) complexes,
reveals that the [Fe(tsbH)2]2+ complex is stabilised to a larger extent
on transfer from water to methanol medium than its analogue
[Fe(tsbMe)2]2+, (R' = Me). This trend in stabilisation is in reverse
order to that already established. The expected order of stabilisation
according to ligand bulk and hydrophobicity, as has already been
discussed in Chapter 4, is Ph ) Me ) H, whereas the actual order for
these terdentate complexes is Ph ) H ) Me, as seen in Figure 6.2.
Figure 6.3 (a,b,c) show plots of Om~e vs volume percentage of
methanol for three, proton on imine nitrogen, bidentate complexes
([Fe(pami)3]2+, [Fe(mpami)3]2+ and [Fe(ppami)31 2+) where they are
compared with their analogue complexes (i.e. ligands containing methyl
h arder of stabilisation group on imine nitrogen). Once again t e reverse
is observed for [Fe(pami)3]2+ and [Fe(hpmi)3]2+ as seen in Figure 6.3(a)
- 127 -
N CD
TABLE 6.1
Derivation of transfer cheMical potential for [Fe(L-L-L)2]2+ complex ions from solubility Deasurements of their salts in water and methanol-water solvent mixtures, at 298.2K.
1 [Fe(tsbh)21 (ClO4)2 1 [Fe (tsbh)21 (PFG )2 1 [Fe (tsbJe) 21 (Cl04)2 1 [Fe(tsbPh)21 (ClO4)2 1 ----------------1 ------------. ----I ----------------1 -----------------
lIeOB 1 €s9~14700 1(-1 CIJ('"""l 1 €s92=14700 X- 1 CIJ('"""1 1 €s91=13170 1(-lCIJ('"""l 1 €s9s=6490 X- 1 CIJ('"""1
.----------------1-----------------1----------------1--------------------Vol~ ASS S C 1 ASS S A 1 ASS S C 1 ASS S C
1 1 0 54.4 1 19.7 1 76.8 1. 77
10 80.0 -2.95 -2.84 23.9 -1.54 +0.63 I 102.0 -2.10 -2.02 20 95.5 -4.24 -4.30 30.2 -3.27 +0.52 116.0 -3.06 -3.12 4.83 -7.46 -7.52 30 120.0 -5.94 -6.08 43.4 -5.96 +0.07 148.5 -4.90 -5.05 40 153.5 -7.76 -7.60 51. 0 -9.82 -1.11 190.8 -6.76 -6.60 16.44 -16.57 -16.41 50 60 272.0 -12.02 -12.36 76.0 -13.98 -0.86 274.0 -9.45 -9.79 48.90 -24.67 -25.01 70 80 290.0 -12.49 -15.65 101. 0 -16.93 -0.64 213.6 -7.60 -10.76 49.00 -24.68 -27.84 90 202.0 -9.81 -16.91 120.0 -3.30 -10.40
100 ----- -- - ._--- _1_ _-.l
ASS = Absorbance of saturated solution Transfer chemical potential (om~e/kJ mol- 1
) for: S = salt, A = anion and C = cation o~~e for perchlorate ian calculated from reference 8
LLL =
I-R'
80 0
Vol% KeOH
'"
rI 0 S
f; ~ "-~~ -10
R'::: Me , ~
, R ::: II
-20
R'= Ph
-30
FIGURE 6.2
Transfer chemical potential for Fe(II) terdentate complexes, lFe(LLL)21 2+, from water to aqueous methanol, at 298.2 K.
- 129 -
'..;.J o
TABLE 6.2 Derivation of transfer cheDdcal potential for [Fe(L-L)31 2 + complexes cations from solubility
measurements of their salts in water and methanol-water solvent mixtures, at 298.2 K.
Complex € 1)(-1 c:xr 1
~ ... _x/llll. Volt JeOH ~
�----------------------------------------------------------• 1 0 20 40 60 80 100
~--------------- ---______ ----1---------------------------------------------------4 [Fe(paDd)3] (Cl04)2
6800 574
ASS S C
13.0 21.8 -3.84 -3.90
47.0 -9.55 -9.39
91.2 -14.48 -14.82
128.0 -16.99 -20.10
32.0 -6.70
-19.30 -------------------1-----------1-----1----------------------------------------------------------
13000 ASS 88.5 126.0 217.0 325.0 334.0 [Fe(~)3] (Cl04)2 572 S -2.63 -6.67 -9.67 -9.87
C -2.69 -6.51 -10.01 -12.93 ~-------------------I-----------I-----I----------------------------------------------------------
17000 ASS 15.5 18.2 34.7 85.9 141.0 [Fe(ppaDd)3] (Cl04)2 590 S -1.37 -6.17 -12.90 -16.58
C -1.43 -6.01 -13.24 -19.74 ~--------------------I-----------I-----I----------------------------------------------------------
11000 ASS 237.0 327.0 427.5 532.5 372.0 42.0 [Fe(hpmi)31 (Cl04)2 551 S -2.39 -4.38 -6.02 -3.35 +12.86
C -2.45 -4.23 -6.36 -6.51 +0.26 r-------------------I-----------I-----I----------------------------------------------------------
11500 ASS 23.2 29.2 42.6 56.6 40.5 [Fe(mpD1)3] (Cl04)2 558 S -1.69 -4.50 -6.61 -4.12
C -1.75 -4.66 -6.95 -7.28 -------------------1-----------1-----1----------------------------------------------------------
14300 ASS 14.3 20.7 42.1 78.0 127.0 94.0 [Fe(ppDd)3] (Cl04)2 565 S -2.76 -8.04 -12.61 -16.24 -14.00
C -2.82 -7.88 -12.95 -19.40 -26.60 I 1 I
ASS = Absorbance of saturated solution Transfer chemical potential (om~e/kJ mol- 1
) for; S = salt, C = cation
'" I 0
r-l 0 a
--' Ij w ...... ~
" ~ -10
:l. C
'<l
-20 a
LL =
'-P--{ 40 80
R-R' I 40 80 0
Vol~ ]leOH
= H, R' = ]Ie ~. I ~R = Ph, R' = H D - Tl, = Jle
R = Ph, R' = R = R' = H R = Xe, R' = H
b c
FIGURE 6.3
Transfer chemical potential for Fe(II) semiaramatic bidentate complexes, [Fe(LL)3]2+. from water to aqueous methanol (same scale for a. b and c). at 298.2 K.
where the farmer is stabilised to a larger extent an tr f f ans er rom water
to methanal (i.e. higher preferential salvation by methanol). The same
pattern is observed for twa complexes derived from 2-acetylpyridine,
which is shawn in Figure 6.3(b), but the difference in stabilisation is
much smaller. In the case of the twa analogues derived from
2-benzoylpyridine the difference in stabilisation is less pronounced but
still obvious. Although bath complexes shaw moderate stabilisation
throughout the whale range, the complex with a proton an the imine
nitrogen shows somewhat higher stabilisation in the methanal region,
while the reverse is true for the water region.
Unlike the aliphatic series of diimine ligands (Chapter 4) where an
increase in the ligand size (different organic group an the ligand's
carbon atom) increases the stabilisation, substitution of a group an the
imine nitrogen has a different effect. Evidently, when R' = H an imine
nitrogens, stabilisation of the complex is higher than when R' = Me.
This can be interpreted in terms of inductive effect within the imine
moiety and hence the acidity of N-H band. The basicity of aqueous
methanol solvent mixtures increases as the methanal fraction increases,
it would be expected that [Fe(pami)3]2+ which contains the most acidic
proton on N would be much more stabilised than its analogue derived from
2-acetylpyridine. In turn both complexes would be more stabilised than
their Me on imine nitrogen analogues. On the other hand, the two
phenyl derivatives, Figure 6.3(c), are well stabilised because of the
hydrophobicity of the ligand, but the N-H analogue would be more so in
the high methanol fraction (Ph is a poor electron donor). This inductive
effect is in the same order as will be seen in Chapter 8.
- 132 -
The acidity of the imine proton is also seen from the lH-nmr
spectra for the [Fe(tsbH)2]2+ complex when the H peak at 8.5 ppm
disappears after introduction of D20 as seen in Figure A3(I) and A3(II)
in Appendix 3. The same was attempted for bsb Fe(II) complexes but
complicated spectra were obtained due to cis and trans isomers<7>. The
presence of two isomers make final positioning of the proton on the
imine nitrogen difficult as shown, for [Fe(mpami)3]2+, in Figure A3(III)
in Appendix 3.
6.3.2 RESULTS AND DISCUSSION OF THE INITIAL STATE - TRANSITION STATE
Unlike the aliphatic and aromatic ligand Fe(II) diimine complexes,
where the rate constant for dissociation with hydroxide increases
markedly with an increase in the content of the organic cosolvent<9-10),
the rate constant for dissociation with hydroxide of semiaromatic ligand
Fe(II) complexes is relatively stable or decreases with an increase in
the content of the organic cosolvent. The k<ob.> values of hydroxide
attack on the Fe(II) diimine complexes are reported in Tables 6.3 and
6.4 for methanol/water solvent mixtures. The second order rate
constant, k2, was calculated, as explained earlier in Chapter 5, from
variation of kabs with hydroxide concentration. The rate constant for
aquation is found to be small, in comparison to k2, and therefore not
reported. The second-order rate constants are reported in Tables 6.5-7,
which show the calculation of solvation contribution on the initial
state-transition state analysis for hydroxide attack on an appropriate
complex.
Of interest is the variation of k2 within a particular ligand
series for a given solvent mixture. In water, the second-order rate
- 133 -
(a)
TABLE 6.3
First-order rate constant. k<obs>. for hydroxide attack on
(a) [Fe(tsbh)21 2+ and (b) [Fe(hpDf)3]2+ complex ions in
water and methanol-water solvent mixtures at 298.21; ionic
strength = 0.33 mol dDr3
104 k(obs>/s-1 for [laDH1/mol dDr3
L-L-L Vol~ KeDR ---------------------------------_________ _
0.18 0.15 0.12 0.10 0.05
0 12.64- 8.40 6.61 4.77 1. 59
20 7.20 5.56 3.86 2.89 1. 10
tsbh 40 5.41 3.99 3.00 2.23 0.94
60 4.41 3.24 2.63 1. 85 0.82
80 2.81 2.33 1. 66 1. 55 0.67
(b)
10sk(obs>/s-1 for [laDHl/mol dDr3
L-L Vol~ KeDR ----------------------------------------0.02 0.015 0.010 0.0050 0.001
0 2.93 2.48 2.01 1. 02 0.50
20 3.08 2.50 1. 77 0.95 0.41
hpui 40 4.01 3.37 2.38 1. 66 0.50
60 6.80 5.67 3.91 2.26 0.94
80 12.40 9.92 7.83 5.04 2.60
- 134 -
TABLE 6.4
First-order rate constant, k<Ob.>, for hydroxide attack on [Fe(LL)31 2 +
complex ions in water and methanol-water solvent mixtures at 298.2(j
ionic strength = 0.33 mol dm-3
L-L Vol~ ReOR --------------------______________________ _
0.02 0.015 0.010 0.005 0.001
0 56.1 40.2 26.0 14.4 3.58
20 42.6 33.1 22.9 11. 2 1.82
mpami 40 38.3 27.9 19.2 9.7 1. 84
60 40.9 32.2 21. 9 10.7 2.74
80 47.9 36.9 25.6 12.6 2.89
10"k(ob.>/S-1 for [Iaml] /mol dDl 3
L-L Vol~ XeOH ---------------------------------------------0.066 0.045 0.025 0.010
0 10.92 7.32 4.08 1. 83
20 14.06 9.82 5.60 2.28
ppami 40 18.12 12.20 7.48 2.74
60 19.08 12.95 7.30 3.23
k /S-l for [laOH1/mol dur 3 (ob.>
L-L Vol~ XeOH ---------------------------------------------
0.04 0.025 0.01 0.005
0 0.059 0.036 0.016 0.0081
0.019 0.0081 0.0053 pami 20 0.028
40 0.020 0.013 0.0049 ------
1]'-) -
W U'\
TABLE 6.5
Initial state-transition state analysis for hydroxide attack on
[Fe(tsbh)2]2+ complex froD water to aqueous methanol at 298.2 K.
Vol ~ ](eGH o 20 40 60 80
--------------------------------------------------------------------------1()3k2/1I- 1 S-l 8.20 4.74 3.41 2.72 1.64
~G·/kJ DlJl- 1 +1.36 +2.18 +2.73 +3.99
m~e{[Fe(tsbh)2]2+}/kJ mcl- 1 -4.30 -7.60 -12.36 -15.65
m~e(OH-)/kJ mcl- 1 -0.15 -0.05 +1.35 +5.75
Initial State -4.45 -7.65 -11.01 -9.90
il'ransition State -3.09 -5.47 -8.27 -5.90
84. 8 kJ D:Jl- 1
1
-0.15
o 40
-8.27
-5.90 T8
"5.75 UH-
80
Vol" MeUB
FIGURE 6.4
Initial state-transition state analysis of reactivity for hydroxide attack on [Fe(tsbh)2)]2+ complex in wateqmethanol mixtures, at 298.2 K.
- 137 -
-> w co
TABLE 6,6
Initial state-transition state analysis for hydroxide attack an [Fe(LL)3]2+
complexes from water to aqueous methanol at 298.2K.
LL = pami LL = ppaDi
Vol 1. )(eOR o 20 40 o 20 40 60
-----------------------------------------------------------------------------------------------k21Jl- 1 S-l 1. 48 0,72 0.513 0.0165 0,0214 0.0276 0.0289
c5 ... ~G·/kJ DKJl- 1 +1. 78 +2.62 -0.64 -1.27 -1.39
c5m~e{[Fe(LL)3]2+}/kJ mol- 1 -3.90 -9.40 -1.43 -6.01 -13.24
c5 ... ~e(OB-)/kJ DKJl- 1 -0.15 -0.05 -0.15 -0.05 +1.35
Initial State -4.05 -9.45 -1.58 -6.06 -11.90
Transition State -2.26 -6,82 -2.22 -7.33 -13.29
...... W \.D
Vol"' ](eon
TABLE 6,7
Initial state-transition state analysis for hydroxide attack an [Fe{LL)3]2+
complexes from water to aqueous methanal at 298.2K,
LL = mpami LL = hpmi
o 20 40 60 80 o 20 40 60 80
--------------------------------------------------------------------------------------------------------------------102 h/X- 1 S-l 2.73 2.15 1. 90 2,04 2.38 0,13 0,14 0.18 0,31 0.52
o .. f1G*/kJ mol- 1 +0.59 +0.90 +0,73 +0.33 -0,22 -0.80 -2,17 -3.44
omp6{[Fe{LL)3]2·}/kJ mol- 1 -2.70 -6.50 -10,00 -13.00 -2.45 -4,23 -6,36 -6.51
o .. p6(On-)/kJ mol- 1 -0,15 -0,05 +1,35 +5.75 -0.15 -0.05 +1.35 +5.75
Initial State -2,85 -6.55 -8,65 -7.25 -2,60 -4.28 -5,01 -0.76
Transition State -2.22 -5.63 -7.84 -6.90 -2.82 -5.08 -7.18 -4.20
--" ;;:-o
-2.22
70.0 kJ DOl- 1 81.9 kJ 1JIJl- 1 83. 1 kJ lKll- 1
1
o
(a) \ -9.40 IS
[Fe (pami) 3) ] 2+
1 -7.84
(b)
TS 1
TS
(c)
-13.00 -13.24
40 o
[Fe (mpaDi) 3)] 2+
40 80 o 40 - [Fe (ppllDi) 3) « •
Vol~ )(eGB Vol~ 1IeOB
FIGURE 6.5 Initial state-transition state analysis of reactivity for hydroxide attack on (a) [Fe(paDd)3)]2+, (b) [Fe(Dpami)3)]2- and (c) [Fe(ppami)3)]2- complex in water/methanol Ddxtures, at 298.2 K.
Vol~ JleGB
constants for [Fe(pami)a]2+, [Fe(mpami)a]2+ and [Fe(ppami)a]2+ are 1.84,
0.027 and 0.0165 M-1 S-l respectively. This decrease in rate constant
for dissociation probably suggests that the steric effects(ll) are
important factors influencing the rate of dissociation of the above
complexes. Further evidence, that this may be the case, is provided by
rate constant (0.0013 M-1 S -l) for [Fe(hpmi)a)2+ when compared with its
analogue [Fe(mpami)3]2+ above. This order of decrease in the rate
constant is the same as that already reported in Chapter 3 for iron(II)
diimine complexes containing aliphatic ligands (gmi, mmi, bmi and cmi).
Figure 6.4 shows the solvation contribution on the initial state-
transition state analysis for hydroxide attack on the FeCII) terdentate
complex. This is the first case for an FeCII) complex where the solvent
effects on the rate constant are dominated by the initial state. In the
case of hydroxide attack on phen(9) and other complexes(10), as well as
that of gmi in the previous chapter, the solvation contribution is
dominated by the transition state which is preferentially stabilised
relative to the corresponding initial state.
As for the terdentate complex above, the same initial state-
transition state analysis is observed in Figure 6.5, for the
[Fe(pami)a]2+ complex, where the initial state is found to dominate. On
the other hand in the case of [Fe(mpami)3]2+ complex bath initial states
and transition states are stabilised to the same extent (solvation
contribution is the same in both cases). Figure 6.5 shows the split
diagram for the [Fe(ppami)a]2+ complex, which contains the bulkiest
ligand of the three within this series. In this latter case, the
solvation contribution to the initial state-transition state is
- 141 -
dominated by the transition state effects, as the transition state is
stabilised to a larger extent than the initial state.
Variation in contribution to the transition state in these three
complexes may be explained in the following manner. The hydroxide
attack is most probably at some position on the ligand(12), in each case
there would be charge reduction on going from initial to transition
state from 2+ to 1+. This charge reduction should lead to preferential
solvation of the transition state. However, in the three ligands
although similar ( Ph-C(R)=N-H ) the R groups vary from hydrophilic
proton in the pami ligand to hydrophobic phenyl in the ppami ligand.
Therefore in the case of the least hydrophobic complex, [Fe(pami)a1 2+,
although there is a charge reduction from 2+ to 1+, on going from
initial state to transition state, this should in effect make the
transition state preferentially solvated by methanol. However, the
incorporation of hydrophilic OH- ion within the bulk of the complex
would make the species in the transition state less hydrophobic than
that in the initial state. Thus the contribution of the incoming
hydrophilic OH- ion may have a higher effect than the reduction of
charge which would lead to reduction in stabilisation on going from
initial to transition state.
In the second complex, [Fe(mpami)a]2+, reduction in charge and
reduction in hydrophobicity of the exterior make opposite contributions
but may be of equal magnitude on going from initial to transition state
thus cancelling each other out - leading to no change in stabilisation
of initial and transition state. On the other hand a very small
reduction in the hydrophobic exterior would be experienced on engulfing
the hydroxide ion into the already hydrophobic lFe(ppami)a]2+ complex,
- 142 -
thus the reduction in charge 2+(initial state) to l+(transition state)
would have a much greater effect leading to the transition state being
preferentially solvated by methanol.
1~3 -
REFERENCES
1. D. H. Bhsch and J. C. Bailar, J. Am. Chern. Soc., 78(1956)6016
2. P. Krumholz, Inorg. Chern. 4(1965)609.
3. P. Krumholz, Inorg. Chem. 4(1965)757.
4. P. F. Figgins and D. H. Busch, 1. Am. Chern. Soc. , 82(1960)820.
5. L. J. Wilson and 1. Bertino, J. Coord. Chem. 1(1971>237
6. M. J. Blandamer, J. Burgess, R. 1. Hains, F. M. Mekhail and P. Askalani, J. Chern. Sac. , Dalton Trans. , (1978) 1001
7. L. 1. Wilson and 1. Bertini, J. Coord. Chem., (1971>237
8. M. H. Abraham, T. Hill, H. C. Ling, R. A. Schulz and R. A. C. Watt J. Chem. Soc. Dalton Trans. I, 80(1984)489
9. M. J. Blandamer, J. Burgess and D. L. Roberts, J. Chern. Soc., Dalton Trans., (1978)1086
10. J. Burgess and C. D. Hubbard, J. Am. Chern. Soc., 106(1984)1717. B. Clark, Ph.D. Thesis, University of Leicester, 1985. N. Gosal, Ph.D. Thesis, University of Leicester, 1986.
11. P. Krumholz, J. Phys. Chern., 60(1956)87. F. Basolo and R. G. Pearson, "Mechanisms of Inorganic Reactions",
John Wiley, (958)152. R. D. Gillard, D. W. Knight and P. A. Williams, Transition Met.
Chem. 5(1980)321 12. R. D. Gillard, Coord. Chern. Rev., 16(1975)67
R. D. Gillard, R. J. Lancashire and P. A. Williams, Transition Met. Chern., 4(1979)115
G. A. Lowrance, D. R. Stranks and S. Suvachittanont, Inorg. Chem., 18(1979)82.
CHAP~'ER
7
Solvatochromism and Solvation of
Fe(II) and Fe(III) Ternary Complexes
7.1 INTRODUCTION
Maximum absorption frequencies of the charge transfer band for
[Fe(bipY)31 2+ and related ligands are insensitive to the solvent.
However, the dependence of maximum absorption frequencies of the charge
transfer band on the nature of the solvent was first reported for its
mixed ligand analogue [Fe(bipY)2(CN)21(1). Further investigation showed
that such solvatochromic behaviour is exhibited by a range of mixed
ligand complexes of transition metals with t2g6 electronic
configuration(2.3). This range included complexes of Fe(I!) bis
diimine, bis-cyanide, [FeCLL)2(CN)21 where LL = bipy or phen(4.S),
substituted phen(S) and bidentate Schiff bases(s>; FeCI!) diimine
tetracyano complexes of the type [Fe(LL) (CN)41 2-, where LL = phen and
bipy(4.S); and MCD) diimine tetracarbonyl complexes of the type
X(CD4)(LL) where K = Cr, Ko and W, and LL = phen, bipy(7.s> or
diazabutadiene(9). Although the above complexes all contain bidentate
ligands and have in common the cis-octahedral geometry(10.11>,
solvatochromic behaviour is also exhibited by complexes containing
monodentate and polydentate ligands(12) as well as complexes with square
planar geometries c13 >.
In this chapter, the range of Fe(II) and FeCI!I) ternary complexes
which exhibit solvatochromism has been extended. Several of the
complexes, of the type [Fe(LL)2CCN)2]O or 1+ and [Fe(LL) (CN)41 2- orl ,
have been prepared and their solvatochromic properties studied in a
restricted number of solvents. Solvent dependence of the lowest energy
charge-transfer bands, obtained in the form of slopes from the
correlation of frequencies of maximum absorption against those for
14~ -
[FeCbipY)2(CN)2] in respective solvents, begin to show a systematic
pattern emerging which is governed by the nature of the ligand and the
oxidation state of the metal. For Fe(II) ternary complexes the charge
transfer band is that of metal to ligand (MLCT) where the frequency of
maximum absorption increases as the polarity of solvent increases (as is
shown by Photograph 7.1) and is dependent on the ligand's rr-acceptor
abilities. On the other hand in Fe(III) complexes which also show
solvatochromic behaviour, the frequency of maximum absorption decreases
as the polarity of the solvent increases indicating that the charge
transfer is that of ligand to metal (LMCT).
Use has been made of Fe(II) and Fe(III) mixed ligand complexes for
their ability to have different charges. Solubility measurements on a
number of these complexes have been carried out and hence transfer
chemical potentials derived and compared, in methanol/water mixtures and
ather aqueous cosolvents.
7.2 EXPERIMENTAL
Compounds of the type [Fe(LL)2(CN)21, where LL = phen or bipy, were
prepared by procedures published by Schilt(4). The Fe(III) analogues of
these two compounds were prepared as nitrate salts; for solubility
purposes an unsuccessful attempt was made to prepare perchlorate salts.
Tetracyano-complexes were also prepared according to Schilt but with
minor modifications whereas their FeCI II) analogues were prepared as
hydrogen salts. [Fe(bipY)3][Fe(bipy) (CN)4] was prepared for solubility
purposes by precipitation of the perchlorate and potassium salts
respectively in 60% aqueous ethanol. The solution was cooled for half
an hour, then the precipitated KCI04 was removed by filtration. The
- 146 -
volume of the solution was reduced by evaporation to a half, which
resulted in crystallisation of the salt. This was collected by
filtration and was washed with the minimum amount of ethanol.
R R'
1 H Me = hpmi
2 H Ph-pMe = hppi-Me
3 Me H = mpami
4 Me Me = mpmi J-R'
5 Me Ph = mppi
6 Ph Ph = bppi
The ternary complexes of the type [Fe(LL)2(CN)2] , where LL is a
semiaromatic(14) ligand 1-6 above, were prepared from their tris-ligand
analogues and three equivalents of KCN in 50% aqueous methanol. The
solutions were left at room temperature for 24 hours and then evaporated
to dryness and complexes were extracted in 60% ethanol solution at O·C.
Oxidation of these complexes to their Fe(III) analogues was
unsuccessful, resulting in the oxidation of the ligand. Microanalyses
were carried out for C, H, and N in characterising (Fe(mpmi)2(CN)2] and
[Fe(mpami)2(CN)2] complexes, both of which were found to contain water
of crystallisation. The results are set out in the Table below.
wt% C H N
( Fe (mpm!) 2 (CN) 2] calculated 57.47 5.36 22.34
[Fe(mpmi)2(CN)2] found 54.88 5.65 21.30
[Fe(mpmi)2(CN)2] . H20 calculated 54.84 5.62 21.32
- 147 -
[Fe(mpami)2(CN)2J calculated
[Fe(mpami)2(CN)2J found
[Fe(mpami)2(CN)2J.3.5H20 calculated
C
55.17
46.79
46.71
H
4.59
4.99
5.50
N
24.14
20. 18
20.38
[Fe(bmi)2(CN)21 and [Fe(cmi)2(CN)2] complexes were prepared from
their tris-ligand analogues by the method described above. Both
complexes were found to be highly hygroscopic - also hygroscopic was
K2[Fe(bmi) (CN)4J. This was prepared from its bis-ligand analogue
<above) and three equivalents of KCN in 50% methanol solution warmed at
60·C for eight hours. The resulting solution was evaporated to dryness,
followed by extraction of the complex in a minimum amount of water.
The electronic spectra of all the above compounds were measured
using SP 800 and SP 8-100 Pye-Unicam Spectrophotometers, using 10mm path
length silica cells. Solubilities for several of the mixed ligand
complexes, and hence transfer chemical potentials, were determined in
methanol/water solvent mixtures at 298.2K. In the case of
[Fe(bipY>2(CN)2] these were also carried out in other alcohols and
acetone/water mixtures. Solubilities were determined by the use of a
spectrophotometer and/or atomic absorption spectrophotometry. In most
cases water dilutions were used, therefore the Amax followed was that
for water since these complexes are solvatochromic. For consistency,
solubility measurements of [Fe(bipY)2(CN)2] were determined by using
both instruments mentioned above.
- 148 -
-> \ .,;)
H;20
MeOH
(Fe (L Lh(( N)2] EtOH
i - PrOH
LL = CH::::3C N
I-J{e
DMSO
CH::::3 NO;2
A CETONE
PHOTOGRAPH 7. 1
Solvatochromic behaviour exhibited by [Fe(hpDi)2(C1J)2] ternary complex
7.3 RESULTS
All the Fe(II) complexes were stable in most f th o e solvents for at
least a few weeks, some as long as several years. On the other hand
Fe(III) complexes were less stable and were found to oxidise alcohols
and other solvents resulting in Fe(II) analogues. Due to their lack of
solubility the electronic spectra of these complexes were measured in
only a limited number of solvents.
Spectroscopic results are reported in Tables 7.1 and 7.2 in the
form of v (frequencies of maximum absorption> together with solvent ET
values(lS). From previous work on ternary complexes correlation of the
lowest energy charge transfer band against solvent ET values has shown
linearity(16). However, it is important to realise that charge transfer
frequencies in hydroxylic and non-hydroxylic solvents give two separate
correlation lines when plotted against respective ET values, which is a
reflection of differing relative importance of hydrogen-bonding on the
charge-transfer energies of the system studied. For iron(II) ternary
complexes the two correlation lines for hydroxylic and non-hydroxylic
solvents are of the same slope(6), as shown in Figure 7.1. As
solubilities of most of these Fe(II) and Fe(III) complexes are
restricted, spectroscopic measurements were obtained in a small number
of solvents (as few as three in some cases), an attempt to eliminate
these two correlation lines was made. This is possible by correlating
frequencies of maximum absorption of two closely related complexes, L~
bmi vs cmi or phen vs bipy in which case the slopes are always
approximately 1. Better correlation however is achieved by taking an
[ F (bi ) (eN)] Therefore, slopes in arbitary standard, in this case e py 2 2·
Tables 7.1 and 7.2 are determined graphically from plots of Vmax, the
- 150 -
~I
60
.. J 50
~4e
.-1
i ~
.-1
",'''':''''v
a:3
JJ ~O""
..... /""'>~
t4i 0'9-~
40-
15 16 17 v/103 CDr 1
FIGURE 7.1
The dependence of the frequency of maximum absorption (v) for the lowest energy charge-transfer band of [Fe(LL)2<CK)21 on the solvent parameter HT ; LL = I-<pyridylmethylene)-3,4-di.ethylaniline). Data fro. Ref. 6.
--' Vl rv
TABLE 7.1
Frequencies of JmxilllJm absorption for the lowest energy charge-transfer bands of [Fe (LL)2 (CI)21 cOlllplexes. Solvent ET values (kcal .,1-1 ), and the solvent sensitivity in the form. of a slope
derived from plots as explained in the text.
v/103 c.- 1
Solvent ET LL
hpmi hppi-Ke mpam mplli mppi bppi bipy
H~ 63.1 17.95 17.86 18.41 19.42 ]leOH 55.5 16.89 16.18 16.95 17.30 16.78 16.13 18.02 EtOH 51. 9 16.58 15.97 16.75 16.98 16.50 15.77 17.61 n-PrOH 50.7 16.53 15.82 16.64 16.89 16.34 15.72 17.48 I
i-PrOH 48.6 16.34 15.72 16.56 16.72 16.18 15.67 17.30 CIbC. 46.0 15.75 15.12 15.87 16.15 15.70 15.15 16.56 DJISO 45.0 15.53 15.01 15.63 15.87 15.43 14.93 16.23 DKF 43.8 15.72 14.96 15.53 15.82 15.43 15.04 16.16 Cl6ICk 43.3 15.97 15.29 16.03 16.29 15.82 15.43 16.56 Cfi2C12 41.1 15.11 16.18 15.53 15.06 16.18 CHC13 39.1 15.45 15.15 15.92 16.13 15.67 15.11 16.34
Slope 0.76 0.65 0.67 0.79 0.66 0.58 1. 00
~
-" \..,'1 W
I I
TABLE 7.2
Frequencies of DaXiDUm absorption (v) for the lowest energy charge-transfer bands of Fe{II) and Fe{II!) dicyano and tetracyano ternary co~lexes. Solvent ET values (kcal mol-'). and solvent
sensitivity in the form of slope derived from plot as explained in the text.
Solvent E'T I
IbO 63.1 18.45 IeOH 55.5 17.79 EtOH 51.9 17.57 n-PrOB 50.7 17.30 i-PrOB 48.6 17.39 ClbCI 46.0 17.09 DJISJ 45.0 16.89 D:IF 43.8 C1bI£k 43.3 17.06 C~12 41.1 17.06 CHC13 39.1 17.00
Slope 0.52
I = [Fe(bDd)2(CI)2] IV = [Fe(bipY)2(CI)2]
VIII= [Fe(phen)2(CI)2]
v/l0:3c.-'
II III IV
19.84 17.98 19.42 18.66 17.52 18.02 18.38 17.36 17.61
17.24 17.48 18.02 17.15 17.30 17.42 16.89 16.56 17.30 16.84 16.23
16.78 16.16 16.50 16.56
17.48 16.92 16.18 17.64 16.95 16.34
0.80 0.44 1. 00
II = [Fe(b~){CI)4]2-V = [Fe(bipY)2(CI)2]+
II = [Fe(phen)2(CI)2]+
V VI VII
18.40 20.49 23.92 19.23 18.38 24.04 19.30
19.42 25.51 25.64
15.75
19.84 26.88
-0.50 1.49 -1.03
III = [Fe(cD!)2(CI)2 VI = [Fe (bipy) (CI)41 2-
I = [Fe(phen) (CI) 4] 2-
VIII IX X
19.30 18.94 21.50 18.18 19.23 19.49 17.76 19.42 17.67 19.49 17.51
16.52 16.61 16.45 16.88
16.78
0.95 -0.26 1.55
VII = [Fe (bipy) (CI)4]-
20.00
18.00
16.00
16.00
/
• = [Fe(bmi) (CN)41 2 -
8 = [Fe(bmi)2(CB)21 o - £Fe(cmi)2(CB)21
18.00 20.00
FIGURE 7.2
Correlation between frequencies of maximum absorption (v) for the lowest energy charge-transfer band, of several Fe(II) ternary complexes, and those for [Fe(bipY)2(CB)2] in the corresponding solvents.
.. I~ ()
(') o M ....... t>
20.00
18.00
16.00
16.00
6 = lFe(phen) (CB)41 2-0= [Fe(bipy) (CB)41 2-T= lFe(phen)2(CB)21 A = lFe(phen)2(CB)21+ • = [Fe(bipY)2(CB)2]+
18.00
FIGURE 7.3
-,
20.00
Correlation between frequencies of maximum absorption (v) for the lowest energy charge-transfer band, of several Fe(II) and Fe(III) ternary cODplexes, and those for lFe(bipY)2(CI)2] in the corresponding solvents.
- 155 -
lowest energy charge transfer bands i t , aga ns those of [Fe(bipY)2(CN)2J
for corresponding solvents as shown in Figures 7.2 and 7.3.
7.4 DISCUSSION
Solvatochromic behaviour was investigated for ternary Fe(II) and
Fe(III) dicyano and tetracyano complexes with bidentate Schiff base
ligands from three series. From the aliphatic series were the bmi and
cmi ligands; mpmi and related ligands from the semiaromatic series,
while from the aromatic series the ligands employed are phen and bipy.
The lowest energy band in the visible region for these complexes, for
which results are reported in Tables 7.1 and 7.2, is attributed to the
charge transfer from iron to Schiff base ligand (t2Q ~ rr*). The solvent
sensitivity of this charge transfer band is reported in Tables 7.1 and
7.2 in the form of slopes. An increase in solvent sensitivity, for
Fe(II) ternary complexes corresponds to increasing value of such slopes,
i.e. positive solvatochromism. The negative solvatochromism is a
feature characteristic of Fe(III) ternary complexes. This sign reversal
may be attributed to the reversal of the direction of rr-charge-transfer,
which is in the direction metal ~ ligand for iron(II) compounds, but for
iron(III) compounds is ligand ~ metal.
The above trend in solvent sensitivity of the charge transfer band
suggests that the cyanide ligand is a specific site of solvation which
affects such trends. For example, solvent effects on the charge
transfer band of lFe(bipY)3]2+ are negligibly small (slope ~ 0), while
for [Fe(bipY)2(CN)2] (slope = 1) and [Fe(bipy) (CN)4]2- (slope = 1.49)
are relatively large. This evidence suggests the direct solvent
ligand is unli kely to be the cause of solvent interaction with bipy
- 156 -
effect on the frequency of maximum absorption. Therefore the source of
solvent effect on the metal to ligand h c arge transfer absorption must be
indirectly through the cyanide ligand in these systems. Consider for
example [Fe(bipY)2(CN)2J complex where the charge transfer band is metal
to ligand and operates through the rr-orbital system, this complex has
maximum absorption frequency value of 19420 cm- 1 in water, a change to a
less polar solvent decreases solvation of the cyanide ligand, this will
effectively decrease the Fe ~ CN rr-back-bonding (decrease in rr-acceptor
ability of the cyanide), thus raising the ground state energy. As the
energy of the excited state will remain approximately constant(17),
there will be a net change in the energy of the absorption as the nature
of the solvent and hence the solvation of the cyanide ligand is changed.
A phenomenon which has been explained by a similar argument is the
protonation of [Fe(bipY)2(CN)2J complex(lS), which is believed to occur
at the cyanide ligand. Here, protonation of the cyanide ligand produces
an increase in the rr-acceptor properties of the cyanide group
(now C=N-H) thus lowering the ground state energy and producing a shift
to higher wavelength. Therefore the solvent dependence of charge
transfer band is ascribed to the solvation variation at the cyanide
ligand on the appropriate iron and Schiff base ligand energy.
7.4.1 Fe(II) dicyano and tetracyano complexes
The lowest energy charge transfer band for Fe(II) ternary complexes
i h with an increase in proton donating power of the shifts to h g er energy
solvent, as illustrated for [Fe(hpmi)2(CN)2J in Photograph 7.1. This
solvent sensitivity is reflected in slopes derived for each complex,
- 157 -
Tables 7.1 and 7.2, the larger the slope the greater the solvent
sensitivity.
Within the group of closely related compounds, such as those
containing semiaromatic ligands, the solvent sensitivity of the charge
transfer band, although it does not vary to a large extent, is
sufficient to reflect the presence of different substituents on the
ligand's ethene carbon and imine nitrogen positions. The presence of
electron releasing methyl groups is reflected in higher solvent
sensitivity particularly when the methyl group is on the imine nitrogen
atom. Thus for the complexes containing hpmi and mpmi ligands, where
both have methyl groups on imine nitrogens and the former has hydrogen
instead of a methyl group on ethene carbon, the slopes are 0.76 and 0.79
respectively. The lower slope of 0.67 is observed for mpami which has
methyl on carbon and proton on imine nitrogen. A lower slope of 0.58 is
observed for the bppi ligand where a phenyl is a substituent in both
positions.
The larger changes in solvent sensitivity are found between the
ligand series. Thus for the entire [Fe(LL)2(CN)2J series the slope is
least <solvent sensitivity least) for the complexes containing aliphatic
ligands, i.e. bmi and cmi, whereas the solvent sensitivity is greatest
for phen and bipy complexes. The order of slopes for the respective
ligand series is as follows; aliphatic < semiaromatic < aromatic, e.g.
(0.52) bmi < (0.79) mpmi < (0.95) phen. This solvent sensitivity order
t bilit (19,20) corresponds to the reverse order of ligand rr-accep or a y .
Therefore the entire [Fe(LL)2(CN)2J series from Tables 7.1 and 7.2 may
be placed in order of rr-acceptor ability according to their slope of
solvent sensitivity.
- 158 -
Tetracyano complexes [Fe(LL) (CN) ]2-4 , are found to have charge
transfer bands which are more sensitive to solvent variation than their
dicyano analogues, results which are summarised in a table below. This
large increase in solvent sensitivity for tetracyano Fe(II) complexes is
in accordance with the view f th i if o e s gn icance attributed to solvation
effects at cyanide ligands as explained earlier.
11gand(LL) [Fe(LL)2(CN)2J [Fe(LL) (CN)4]2
bmi 0.52 0.80
phen 0.95 1.49
bipy 1.00 1.55
7.4.2 Fe(III) dicyano and tetracyano complexes
As with tetracyano Fe(II) complexes, Fe(III) ternary complexes are
found to be insoluble in most of the non aqueous solvents. A further
complication was that they were found to be susceptible to reduction
when dissolved in most binary solvent mixtures. However it was possible
to follow solvent effects on lowest energy charge transfer bands in a
limited number of solvents, results of which are in Table 7.2.
It is of interest to notice that the frequencies of maximum
absorption decrease as the polarity of the solvent increases, resulting
in a negative slope which is opposite to their Fe(II) derivatives, as
shown in Figure 7.3. The slopes of solvent sensitivity for
[Fe(bipY)2(CN)2]+ and [Fe(phen)2(CN)2]+ are -0.26 and -0.50
respectively. A much higher negative slope value (-1.03) is obtained
for the tetracyano complex, (Fe(bipy) (CN)4]-, which once more confirms
the significance attributed to solvation effects at cyanide ligands.
- 159 -
7.5 TRANSFER CHEMICAL POTENTIAL OF TERNARY COMPLEXES
Solubility results and derived transfer chemical potentials of
ternary complexes are summarised in Tables 7.3-5. In Chapter 4 we have
seen the Fe(II) tris-ligand complexes of phen and bipy are
preferentially solvated by methanol, i.e. both complexes are stabilised
on transfer from water to methanol, the phen complex to a larger extent.
The Fe(II) ternary complexes of phen and bipy are also found to be
stabilised on transfer from water to methanol, however stabilisation is
much less than that of their tris analogues. The same is found for
other ternary complexes, i.e. those containing ligands mpmi and mpami,
when compared with their tris analogues as shown in Figure 7.4. Neutral
complexes are expected to be stabilised by methanol, however the
presence of cyanides in ternary complexes increases the hydrophilic
character in the periphery of the complex which consequently leads to
lower stabilisation, i.e. the cyanides are preferentially solvated by
water.
Evidence that the neutral complex is more stabilised than the
charged complex is provided by an example of Fe(ll) and Fe(III) bis
cyanide bis-bipy complexes as shown in Figure 7.5. Replacement of
another bipy ligand by two cyanides leads to (Fe(bipy) (CN)41 2-, a
dinegatively charged tetracyano complex which has a predominately
hydrophilic periphery and would be expected to be preferentially
solvated by water. Figure 7.5 shows this complex to be destabilised on
transfer from water to methanol almost to the same extent as the
dinegatively charged pentacyano complex nitroprusside.
In order to investigate preferential solvation of ternary complexes
9 f (F (bi ) (CN)21 were obtained in a series of organic cosolvents Om~ 0 e py 2
- 160 -
from solubilities of this complex in several binary aqueous mixtures.
Figure 7.6 shows a plot of transfer chemical potentials of this complex
vs mole fraction of four alcohols and acetone. The complex is found to
be initially preferentially solvated by organic cosolvents in the order
KeOH < EtOH < i-PrOH < t-BuOH < Acetone. This trend in preferential
solvation is only true for the lower region of organic cosolvent. At
higher mole fraction of the organic cosolvent an inflection of "roller
coaster type" takes place and the trend in preferential solvation is
reversed, i.e. MeOH ) EtOH ) i-PrOH ) t-BuOH ) Acetone. This trend in
stabilisation may be interpreted in terms of decrease in acidity of the
cosolvent media which decreases from methanol to t-butyl alcohol to
acetone.
- 161 -
~
7' ~J
TABLE 7.3
Solubility and derived transfer chemical potentialCkJ mol- 1 ) for ternary FeCII) dicyano complexes in methanol/water solvent mixtures,at 298.2 K.
Vol~ ](eaR 0 20 40 60 80 100
sol (a) 0.585 0.877 5.230 18.00 26.50 OmJi8 -1. 01 -5.43 -8.49 -9.45
[Fe (bipy) 2 (CI) 2] ASS 2.76 5.25 34.76 102.60 207.00 160.00 OmJi8 -1.59 -6.30 -8.96 -10.70 -10.06 ---------------------------------------------------
avo OmJi8 -1.30 -5.86 -8.72 -10.70 -9.75
ASS 0.220 0.648 3.360 11. 70 27.60 25.00 [ Fe (phen) 2 (CI) 2] OmJi8 -2.67 -6.75 -9.84 -11. 9? -11. 73
sol (a) 0.210 0.243 0.365 0.716 1.206 2.98 OmJi8 -0.38 -1.36 -3.03 -4.33 -6.57
[Fe(mpaD!)2(CI)2 ------------------------------------------------------sol (a) 0.21 0.361 0.728 1. 210 2.98 OmJi8 -1.34 -3.08 -4.34- -6.57 ---------------------------------------------------
avo OmJl8 -0.86 -2.22 -3.68 -4.33 -6.57
[Fe(mpDd)2(CI)2] ASS 83 176 322 881 626 OmJl8 -1. 87 -3.36 -5.85 -5.01
- - ------------
(a) = Solubility(mol dBr3) fro. atomic absorption spectroscopy
ASS = Absorbance of saturated solution
--' (J"\ w
TABLE 7.4
Solubility and derived transfer cheDdcal potential for ternary Fe{!I) cOBplexes
in uethanol/water solvent mixtures,at 298.2 K.
Vol"' Jlethanol
[Fe{phen)2{CI)2]I~
~
S
A
C
o
3.40
20
8.30
-4.40
-0.20
-4.20
40
13.30
-6.76
+0.80
-7.56
60
14.90
-7.32
+2.50
-9.82
80
17.20
-8.03
100
~----------------------------------------------------------------------------------------------
[Fe(bipy)2{CI)2]I~
ASS
S
C
5.70 7.50
-1.35
-1.15
12.00
-3.66
-4.55
14.10
-4.48
-6.99
~----------------------------------------------------------------------------------------------
[Fe (bipY)3]
[Fe (bipy) (CILd
ASS
S
A
28.2
ASS = absorbance of saturated solution
39.0
-1. 61
+2.30
6.p8/kJ mol- 1 for salt(S), anion(!) and cation(C)
6mp8([ID3)-) from reference 21
71. 0
-4.57
+3.68
124.0
-7.34
+5.41
132.0
-7.65
+7.53
24.66
+0.66
+12.10
-> ()\ .:::-
TABLE 7.5
Solubility and derived transfer chemical potential for [Fe(bipY)2(CI)2] complex
in EtOH/, i-PrOH/, t-BuOHI and acetone/water solvent Ddxtures, at 298.2 K.
Vol1.
EtOH ASS
C
o
1. 20
10 20 30 40 60 80 90 100
2.10 4.36 10.22 20.68 39.48 40.30 38.44 40.92
-1.38 -3.19 -5.31 -7.05 -8.66 -8.71 -8.59 -8.74
~--------I-----------------------------------------------------
i-PrOH ASS
C
1.19 2.37 4.98 10.01 14.85 14.08 6.40 3.84 17.70
-1.70 -3.55 -5.28 -6.25 -6.12 -4.17 -2.90 -6.69
~---------I------------------------------------------------------------
t-BuOH ASS
C
1.10 2.61 5.64 8.20 14.50 10.20 3.40 1. 60
-2.14 -4.05 -4.98 -6.39 -5.25 -2.79 -0.93
-------1----------------------------------------------------------Acetone 1.10 3.06 6.39 12.10 18.40 19.32 7.37 1. 98 0.20 ASS
C -2.53 -4.36 -5.94 -6.96 -7.10 -4.71 -1.46 +3.99
~
ASS = absorbance of saturated solution
C = o~~e/kJ .01- 1
0 50 100
Vol"' Mp.UII
....... po ....... I ~" r1 " " 0 ......... @ ~
....... .......
"-....... ~
I-) ....... -.......
,!st ~
"-~~ E
<0
-10
® = [ Fe (mpn m 1 ) :2 (C R ) :2 1 X = [ Fe (mpllml )3] 2+
• _. [Fe (mpml):2 (CR):21 0 = (Fe <mpml )31:2+
• = ( Fe(blpY)2(CR):21 -20 0 = [ Fe(hlpY)3]:2+
A = ( Fe{phen):2(CR):;d 6. := [ Fe (phen) 31:2+
FIGURB 7.4
Transfer chemical potentials for some Fe(II) ternary complexes and their tris-ligand analogues from water to aqueous methanol, at 298.2 K.
- 165 -
+10
v = £Fe(CB)&BO]2-~ = £Fe(bipy) {CR)4]2-6. = CI-A = o =
£Fe(bipY)2(CH)2]+ £Fe(bipY)2(CH)2]
• = [Fe(bipY)3]2+
I /
/ I:::.
I:::.
/
50 / 0~~~ ____________ ~ ______ ~/~ ______ ~1~OO
/ 1:::./
Vol~ J(eOH
,..".,
""- -/:).- -
-10
FIGURE 7.5
Influence of hydrophobic/hydrophilic character and the charge on transfer chemical potentials for iron complexes fro. water to aqueous methanol, at 298.2K
- 1()6 -
+10
'" I r-l 0 I a 0
~ ~
" t
1_J \\ ...... 0\ --.:J
-20
0.2 0.4 0.6 /
lI.f.
\_/ o = XeOH • = EtOH • = i-PrOH • = t-BuOH ~ = Acetone
FIGURE 7.6
Transfer chemical potentials for [Fe(bipY)2(CK)2] comple~ from water to aqueous cosolvents, at 298.2 K.
1.0
REFERENCES
1. J. Bjerrum, A. W. Adamson and O. Bostrup, Acta Chern. S cando , 10 <1956 )329.
2. H. Back and H. tam Dieck, Angew. Chern. Internat. Ed., 5(1966)520.
3. H. Back and H. tam Dieck, Chern. Rev., 100(1967)228.
4. A. A. Schilt, J. Am. Chern. Sac., 82(1960)3000.
5. J. Burgess, Spectrochim. Acta, 26A(1970) 1369.
6. J. Burgess, Spectrochim. Acta, 26A (1970 >1957.
7. J. Burgess, J. Organometal. Chern. , 19 (1969) 218.
8. J. Burgess, J. G. Chambers and I. R. Haines, Transition Met. Chern. 6(1981)145.
9. H. tam Dieck and I. W. Renk, Angew. Chern. Internat. Ed., 9(1970)793
10. E. D. McKenzie, Coord. Chern. Rev., 6(1971)187.
11. M. Davidson, T. W. Faulkner, M. A. Green, and E. D. McKenzie, Inorg. Chim. Acta, 9(1974)231.
12 J. Burgess and M. W. Twigg, J. Chern. Soc. Dalton Trans., (1974)2032
13. P. M. Gidney, R. D. Gillard and B. T. Heaton, J. Chern. Soc. Dalton Trans., (1973)132
I. G. Dance and T. R. Miller, Chern. Comm., (1973)433.
14. P. Krumholz, Inorg. Chern. 4(1965)609.
15. C. Reichardt, Angew. Chern. Internat. Ed., 4(1965)29.
16. R. I. Hains, PhD Thesis, University of Leicester, 1977.
17. H. Kobayashi, B. V. Agarwala and Y. Kaizu, Bull. Chern. Sac. Japan, 48(1975)465
18. N. K. Hamer and L. E. Orgel, Nature, 190(1961)439.
19. L. H. Staal, D. J. Stufkens and A. Oskam, Inorg, Chim. Acta, 26(1978)255
J. Reinhold, R. Benedix, P. Birner and H. Hennig, Inorg. Chim. Acta 33(1979)209
20. D. Walther and E. Uhlig, Coord. Chem. Rev., 33(1980)3
21. J. Burgess and E-E. A. Abu-Garib, Transition Met. Chern., 9(1984)234
- 168 -
CHAP'rER
a
Redox Reactions of Fe(II) and
Fe(III) Ternary Complexes
8.1 INTRODUCTION
The dissection of solvent effects on reaction rate constants into
initial state and transition state components in the previous chapters
was carried out for a number of reactions of inorganic complexes. All
of those systems involved SUbstitution processes. The work in this
chapter deals with oxidation-reduction or redox reactions. Of the two
mechanisms recognised in redox reactions involving transition metal
complexes<1.2), outer- and inner-sphere, the outer-sphere reactions are
simpler, since electron transfer leaves the coordination sphere of each
metal ion unaltered.
Solvation effects on reactivities for outer-sphere redox reactions
have recently been investigated between pairs of transition metal
complexes<3-S) and simple ions<6>, whose rate constant data have been
rationalised in terms of the Marcus theory<7>. Most of the reactions
studied were for the high charge reactants. However, even for outer-
sphere redox reactions, the solvent may affect one or both of the two
components, the pre-association of the reaction:
Ox + Red OX,Red
and subsequent electron transfer:
OX,Red ==~ Products
To minimise the importance of the association step uncharged
organic reductants are preferred such as catechols (where catechol =
1,2-dihydroxybenzene L). In this chapter the preliminary study of the
solvent effect on reactivity of a well characterised outer-sphere
electron transfer reaction was investigated, for oxidation of catechols
with [IrCl6]2- and with [Fe(bipy) (CN)4]-' The rate constants are
- 169 -
reported for [Fe(bipy)(CN)4J- oxidation of catechol and of
t-butylcatechol and for [IrCls]2- oxidation of t-butylcatechol in
methanol/water solvent mixtures. Transfer chemical potentials for
several catechols and quinols <quinol = 1,4-dihydroxybenzene a) are
OH IJ0-O--0H 2.
derived from their solubilities in binary solvent mixtures. With the aid
of the necessary kinetic data and the thermodynamic results the
dissection of solvent effects on reaction rates constant into initial
state and transition state components has been carried out for
hexachloroiridate<IV) oxidation of t-butylcatechol and peroxodisulphate
oxidation (4 ) of Fe(II) dicyanide and tetracyanide ternary complexes.
Oxidation of Fe(II) complexes of hexadentate ligands, a, by
peroxodisulphate and by Ce 4+ gives an indication of ligand oxidation
resulting in formation of the third imine moiety. This belief prompted
the preparation of the analogue Fe(II) hexadentate complex containing
three imine moieties.
8.2 EXPERIMENTAL
The Fe(II) and Fe(III) dicyano and tetracyano complexes of bipy and
phen were prepared by published methods CS', as explained in Chapter 7.
The other
purified.
reagents were Analar grade; the organic solvents used were
The mixed solvents used were prepared where composition by
- 170 -
volume is before mixing. The solubility determinations were carried out
at 298.2 K as explained previously. In the case of catechols and
quinols, solubility measurements were carried out by i hi we g ng the dry
compound after evaporation of the solvent from a known volume of
saturated solution. This method for solubility measurements was
preferred to spectrophotometric methods because the saturated solutions
were very strong, therefore the enormous dilutions required could have
introduced some air oxidation to the solute. Kinetic runs were carried
out on the 'Hi-Tech Scientific SF-3L' stopped-flow spectrophotometer;
reagent concentrations and conditions are specified in the following
section together with results.
8.3 RESULTS AND DISCUSSION
Prior to redox reactions, oxidising species were analysed for their
reactivity towards solvents employed. The Fe(III) complexes of bipy,
analogues, t~ ~ 40 hours, in methanol and methanol/water solvent
mixtures. However, marginally shorter half lives were observed for
oxidation of ethanol by these two complexes. The latter complex was
found to undergo disproportionation reaction in water, resulting in
tris-bipy Fe(II) complex. It is known that [IrClG]2- undergoes
aquation(9) much more slowly than it oxidises reactant substances. An
attempt was made to study systems involving Ni(IV) dioximes as oxidants.
Unfortunately it was found that these oxidised organic casal vents far
too quickly(lO). Oxidation of methanol and ethanol in particular was
found to be far too fast even to obtain reasonable solubility
measurements of these complexes in such cosolvent systems.
- 171 -
8.3.1 Oxidation of catechols by [Fe(bipy) (CN)4]- and [IrC16
]2-
Oxidation of catechols by several complexes has been analysed
recently in terms of different pathways(11,12>. The kinetic
measurements, for oxidation of catechol by [Fe(bipy)(CN)4)-, were made
in a solution where pH was varied from 2.6 to 5.6 (using HCl04 and
phthalate buffer) and an ionic strength of 0.1 mol dm-3 maintained by
NaCl04. The experiments were run under pseudo first order conditions
with an excess of catechol (5.26 - 79.0 x 10-3 mol dm-3 ) while the
initial concentration of the [Fe(bipy) (CN)4)- was maintained at
approximately 5.2 x 10-4 mol dm-3 in all runs. The rates of oxidation
were monitored at the absorption maximum of the [Fe(bipy) (CN)4)2-
complex at 484 nm using a stopped-flow spectrophotometer thermostated at
298.2K. The k<ob.> values obtained at various pHs and different
concentrations of catechols are summarised as an average of a number of
runs in Table 8.1.
A plot of k(ob.> against concentration of catechol, at constant pH,
gives a straight line as shown in Figure 8.1(a); indicating that the
oxidation process shows a first order dependence on catechol
concentration. The k2 values are derived from variation of K<ob.)
values with catechol concentration. A plot of k2 values vs pH, in
Figure 8.1(b), shows an increase in rate constant with a decrease in
acid concentration. Dependence of rate constant on acid concentration
is also observed in oxidation of t-butylcatechol by [Fe(bipy) (CN)4]-'
This pH dependence may indicate that the principal reductant may be HC
where C = i, rather than the protonated form af catechol, 'H2 C'.
However this pH dependence may also be due to the nature of the oxidant
which is believed to be a cyanide ligand protonated species below, ~.
- 172 -
TABLE 8.1
Mean k<ob.> and derived k2 values for catechol and t-butylcatechol
oxidation by [Fe(bipy) (CI)A1- in aqueous solution at 298.2K.
pH 10. runs [catecho11/mol dmr3 k(ob.>/S-l b/){-l S-l
2.6 9 0.0124 0.88 71. 0
3.0 12 0.0124 0.94 76.0
3.6 11 0.0124 1.10 88.7
3.6 8 0.0180 1.98 110.0
3.6 9 0.0248 2.54 102.4
4.0 10 0.0124 1.63 131.5
4.6 9 0.00562 1.58 281.0
4.6 12 0.01065 3.45 324.0
4.6 15 0.0255 7.84 307.0
4.6 17 0.0500 16.20 324.0
4.6 12 0.0790 24.4 309.0
5.0 8 0.0133 4.8 361.0
5.6 9 0.0120 7.4 617.0
pH 10. runs [4-t-butylcatecholJ k<ob.)/S-l b/){-1 6 - 1
4.0 6 0.0113 1. 45 128.3
4.45 7 0.0102 2.73 267.6
5.0 6 0.0133 4.56 342.8
5.0 5 0.0364 13.44 369.2
- 173 -
20
~
I U'l
....... 1"1
• .a --.l 0 ;;::- " ~ 110
800
~
/ <a.)
I
t600 to ~
I <b)
:.: ....... ~ ~
/ J. 400
0,
i 200
0.02 0.04 0.06 0.08
[catechol]/mol dm- 3 , (pH = 4.6) 2.0 4.0
pH
FIGURE 8.1
Dependence of kcoo-) on concentration of catechol (a), and dependence of k2 on pH (b) for reduction of [Fe(bipy) (CI)A]- cOBplex.
6.0
The latter is of particular relevance t h o t e solvatochromic property
displayed by this Fe(III) complex.
CI
1;: ",I /" CI-,. 'r-! Fe ": H .c""'/I 'CI/-
CI 5.
Catechol oxidation by [IrCl 6]2-, which was carried out in Torino,
is found to follow a second-order rate law(13). The results in Table
8.2 summarise k2 terms for t-butylcatechol oxidation by [IrCl6]2- in
water and aqueous methanol, ethanol and DMSO together with the derived
om6G* terms. Transfer chemical potentials of several catechols and
quinols have been derived from their solubilities in methanol/water
solvent mixtures which are shown in Table 8.3 (further solubilities and
transfer chemical potentials for hydroxyquinone in several water/alcohol
solvent mixtures are summarised in Appendix 4). A plot of transfer
chemical potential vs volume percentage of methanol, Figure 8.2, shows
preferential solvation of catechols and quinols by methanol, which
increases in the order of the hydrophobic character, that is
2,4-bi-t-butylcatechol and t-butylhydroquinone are most preferentially
solvated by methanol. It is evident from Figure 8.2 that quinols are
generally more preferentially solvated than the corresponding catechols.
Due to the difficulty experienced in determining solubilities for
t-butylcatechol, in methanol/water solvent mixtures, its transfer
chemical potentials were derived from extrapolation of relative values
of analogous catechols and quinols as shown in Figure 8.4.
The analysis of the reactivity trend into initial state-transition
state contributions, for t-butylcatechol oxidation by [lrC16]2- in
- 175 -
TABLE 8.2
Second order rate constants, k2/I-1 S-l, and derived o-6G*/kJ mol-3 for
oxidation of t-butylcatechol by (IrC16]2- in water and binary aqueous
solvent Ddxtures at pH = 2(HCIO.) and at 298.2K
--' Vol~ 0 10 20 30 40 50 60
--J (J'\
1---------------------------------------------------------------------------------------------)'ethanol 103 k2 20.2 11. 0 6.36 3.00 1. 50 0.657 0.27
o .. ~G* +1.53 +2.80 +4.70 +6.44- +8.46 +10.69
Ethanol 103 k2 20.2 11. 0 6.40 3.25 1. 36 0.605 0.31
o_~G* +1.53 +2.85 +4.53 +6.68 +8.69 +10.35
DJISO 103 k2 20.2 6.17 1. 90 0.53 0.09 0.016
o .. ~G* +2.94 +5.86 +9.02 +13.41 +17.7
--.J --.J
TABLE 8.3
Solubility and transfer chemical potentials for catechols and quinols from water to aqueous methanol, at 298.2 K.
Xethanol 2,4-bi-t-butyl- Xethyl- t-Butyl-Catechol -Catechol Hydroquinonel Hydroquinonel Hydroquinone
Voll Vt X. F. ------------- --------------- -------------1------------- ----------------~ I ~ A B
I 0 0 0 3.65
10 8.1 0.047 4.558 -0.55 20 16.5 0.100 5.002 -0.78 30 25.3 0.160 5.121 -0.84 40 34.5 0.229 5.300 -0.92 50 44.2 0.308 5.580 -1. 05 60 54.3 0.400 5.420 -0.98 70 64.8 0.509 80 76.0 0.640 5.510 -1. 05 90 87.7 0.800 5.470 -1. 00
100 100 1.000 5.300 -0.93 ------- ~--~----~~
A = Solubility/mol dDr3
B = OM~e/kJ mol- 1
A B A B I A B A B I
0.0036 0.650 0.517 0.0232 0.0041 -0.29 0.841 -0.64 0.0063 -1.39 0.988 -1. 04 1. 449 -2.28 0.0894 -3.34 0.0099 -2.52 1. 244 -1. 61 0.0171 -3.86 1.641 -2.31 2.720 -3.84 0.8551 -8.94
2.012 -2.80 2.4770 -16.19 2.380 -3.22 I 3.717 -4.65 2.9081 -11. 98
2.604 -3.44 3.1060 -16.75 2.935 -3.74 4.766 -5.23 3.8900 -12.69
2.948 -3.75 I 4.268 -4.96 4.0480 -12.19 3.3910 -16.91 3.034 -3.82 I 4.084 -4.85 3.5545 -12.47
20 O~~~ ___________ 4~O ______ ~6~0 ______ ~80~ ____ ~1~00 Vol'h KeOll
-10
-15
"" -
• • catechol
hydroquinone
Methyl-hydroquinone
"'0- __ ~ "'-- _ - ~ \!.r- - -0- - - ~
t-butylcatechol
t-butylhydroqulnone
2,4-t-butylcatechol
FIGURB 8.2
Transfer cheBdcal potentials for catechols and qUinols fro. water to aqueous methanol, at 298.2 K
- 178 -
-> --J \.0
TABLE 8.4
Initial state and transition state solvation contributions to reactivity for the
[IrC16 ]2- oxidation of t-butylcatechol in aqueous methanol, at 298.2K.
Volt. ](eOR o 10 20 30 40 50 60
I 103 k2/X- 1 S-l 20.2 11. 0 6.36 3.00 1. 50 0.65'1 0.2'1
o .. 6.G*/kJ Ill- 1 +1.51 +2.80 +4.'10 +6.44 +8.46 +10.69
o .. pe(t-butylcatecol)/kJ mol- 1 -0.90 -1.90 -3.30 -4.40 -5.00 -5.50
o .. pe([IrC16]2-)/kJ DDl- 1 - +2.10 +3.90 +5.40 +'1.20 +9.30
Initial State +1.20 +2.00 +2.10 +2.80 +4.30
Transition State +2.'11 +4.80 +6.80 +9.24 +12.76
a from reference 6
T 4U.2 kJ mol-I
+3.90
f2.UO
-4.40
o 20 tiO
FIGURE 8.3
T.S.
[lrCl,d2-
1. S.
t-butylcatechul -~.~O
60
Vol7. KeOH
'Step-diagram' showing initial state-transition state analysis of reactivity trend for the lIrCl61 2 - oxidation of t-butylcatechol from water to aqueous uetbanol, at 298.2 K.
- 180 -
methanol/water solvent mixtures, is set in Table 8.4 and depicted in
Figure 8.3. Both initial state and transition state are destabilised on
transfer from water to methanol transition state bei d t bi ' ng es a lised to
a much larger extent. The marked decrease in rate constant is therefore
largely due to destabilisation of the transition state as the proportion
of methanol increases. The same is probably true for aqueous ethanol,
but due to lack of transfer parameters in ethanol- and DKSO-water binary
systems it is not possible to carry out dissection of solvent effects on
reactivity into initial state-transition state components for these
cosolvents.
8.3.2 Peroxodisulphate oxidation of Fe(II) ternary complexes
It has already been suggested(3) that transition state effects are
very much more important than initial state effects for peroxodisulphate
oxidation of [Fe(bipY)2(CN)21 and [Fe(bipy) (CN)41 2- complexes in aqueous
methanol. The rate constants for both systems(4) are found to decrease
on transfer from water to methanol. Transfer chemical potentials for
the two Fe(II) complexes, as reported in Chapter 7, show that the
neutral Fe(II) dicyano complex is stabilised on transfer from water to
methanol while the dinegatively charged Fe(II) tetracyano complex is
destabilised.
The analysis of reactivity trends into initial state-transition
state contributions for peroxodisulphate oxidation of the two Fe(II)
complexes in aqueous methanol are shown in Table 8.5(a)(b) and depicted
in Figure 8.4. It is of interest to note that initial state and
transition state effects for the Fe(II) dicyano complex are opposite,
the initial state being stabilised while the transition state is
- 181 -
Table 8.5
Initial state and transition state contributions to reactivity
in the peroxodisulphate oxidation of [Fe(bipY)2(CI)21 (a) and
lFe(bipy)(CI)4]2- (b) in aqueous methanol. at 298.2 I.
(a)
Vol~ XeGH o 20 40
~---------------------------------------------------------------------
~/J(-1 S-l - 0.600 0.056 0.015 6'",6G"/kJ 1101- 1 +5.90 +9.00
6''''~e{[Fe(bipy)2(CI)21}/kJ mol- 1 b -1.30 -5.86 6'm~e{(S2De)2-}/kJ mol- 1 c +1.00 +2.70
Init.State -0.30 -3.16
Tran.State +5.60 +5.84
(b)
Vol~ XeGH o 20 40
~---------------------------------------------------------------------
d 0.195
6'm6G"/kJ 1101- 1
6'm~e{[Fe(bipy)(CI)4]2-}/kJ mol- 1 -
6'm~e{(~De)2-}/kJ mo1- 1
Init.State
Trns.State
a from reference 4 b from Table 7.3 c from reference 13(a) d from reference 4 e from Table 7.4
c
- 182 -
0.123 0.170
+1.14 +0.34
+2.30 +3.68
+1.00 +2.70
+3.30 +6.38
+4.44 +6.72
:a '....J
+6.72 +5.84
+5.60 - IS ~IS
+4.44
/
(a)
I <b) +6.38
74.2 kJ Kll- 1
77.0 kJ Kll- 1
IS
o
+3.68 +2.70
+3.30 ___ [Fe (bipy) (C.NLd 2-
+1.00
-0.30
-1.30
[S2 Ga ] 2-
IS
-5.86 [Fe (bipY)2 <c:rn2]
[~Oa]2-
+1.00
Vol't XeOH
20 40 o 20 40
FIGURE 8.4
'Step-diagram' showing initial state-transition state analysis of reactivity trend for peroxodisulphate oxidation of [Pe(bipY)2(CI>2] (a) and [Fe(bipy) (CI>4]2- (b) from water to aqueous methanol, at 298.2 K.
destabilised on transfer from water into methanol, Figure 8.4(a). On
the other hand, for the Fe(II) tetracyano complex, both initial state
and transition state are destabilised by almost the same amount on such
transfer,as shown in Figure 8.4(b).
8.3.3 Ligand oxidation in Fe(II) hexadentate complexes
Complexes dealt with here are Schiff bases of Fe(Il) where the
ligands hexadentate dioxime and hxsbPh were derived by condensation of
2,3-butadione monoxime and 2-benzoyl pyridine with trien
respectively(14,lS>, as described in Chapter 4. The common element in
these complexes is that two nitrogens from the ligand do not form an
imine moiety, 3..
The aliphatic (gmi and bmi) and semiaromatic (hpmi and mpmi)
ligands make the resulting Fe(II) complexes display acidity-dependent
oxidation properties. If the acid concentration is large enough only
the central metal ion is oxidised to Fe(lII) state(16). At lower
acidities ligand oxidation takes place resulting in Q, L and a species
for gmi, hpmi and mpmi Fe(II) complexes respectively(17,lB'.
H
H OR OB I C=O " / C-c ! ,
Ie-I I-Me (Q'> <
I_lie (Q'> <
I-J(e
This behaviour is also displayed by Fe(II) hexadentate complexes but of
interest is [Fe(hxsbPh)]2+ complex where oxidation of the ligand takes
place resulting in formation of the third diimine moiety ~.
oxidation in this complex is analogousto oxidative dehydrogenation of
diamine ligands reported by Goedken(19)( S h see c eme 1 below>.
lib IH / ---Clb (h / ~CH
(Cff) 4Fe I ------. (CN) 4Fe I '" _______ Clb '" ~CH IH2 IH
Sche1De 1
Oxidation of the (Fe(hxsbPh)]2- complex with peroxodisulphate as an
oxidant, was followed spectophotometrically. Th ti t e reac on ra e was
found to be fast initially and decreasing gradually to a stop, at which
point the absorbance of the resulting species in solution was 2/3 that
of the original complex. There was no observed change in ~m_x,
(606nm)(20), but the reaction was found to go to completion within one
day. The nature of this oxidised species at equilibrium was
investigated using cerium(IV) as an oxidant. Cerium(IV) oxidation of
the (Fe(hxsbPh)]2+ complex, with a large excess of oxidant, decolourised
the purple-blue solution instantaneously. However, slow return of the
purple-blue colour was observed. From stoichiometric titrations,
results show that 3.9 to 5.0 equvivalents of cerium(IV) are consumed per
mole of (Fe(hxsbPh)]2- in order to reduce the absorbance of the solution
by one third. Further addition of oxidant did not result in any
significant decrease in absorbance. The same was attempted for the
Fe <I 1) oxime complex, 7-9 equivalents of Ce 4 + were consumed in order to
reduce the absorbance of the solution by one third. This decrease in
absorbance is accompanied by a change in ~m.x from 518 to 579 nm.
- 185 -
The approximately four equivalents of Ce d + involved in oxidation of
the [Fe(hxsbPh)]2+ complex give further indication that the four
electron oxidation at only one possible position in the ligand takes
place, that being the formation of the third diimine moiety.
Preparation of an Fe(II) complex with h d ttl a exa en a e igand containing
three diimine moieties, was attempted using several different procedures
and starting materials. Only one of these gave the desired product, as
represented below.
2(en) + glyoxal (in 501, aqueous ethanol) ~ ~ ;----\
to Ibl I I 102
1 I 2 (2-8ce~yl pyridine)
The complex obtained as perchlorate salt from the above procedure has
the same ~M.~(605 nm) as its parent complex, (Fe(hxsbKe»)2+, prepared by
condensation of 2-acetyl pyridine with trien (see Chapter 4). Further,
its extinction coefficient (5840 M-'cm-') is approximately one third
lower than that for the parent complex (9280 M-l cm-'). The complex was
found to be unstable in water over a long period of time (~ 24 - 30
hours). The unexpectedly lower value for the extinction coefficient of
this complex is probably due to the strain introduced, and therefore
decrease in MLCT/rr-bonding, in the complex by formation of the third
imine moiety, which was revealed on building a model of such a complex.
This preliminary redox study is indicative that Fe([I) compounds
are good substrates to be used for initial state-transition state
- 186 -
analyses. The solubility data and derived transfer parameters for phe~
and bipy dicyano Fe(III) complexes in methanol/water mixtures, from
Chapter 7, might be of interest to study further the oxidation of
catechols and quinols in this binary system in the future. We have seen
that Fe(III) bipy tetracyano complex is a suitable oxidant for
catechols. However it might prove difficult to study in binary systems
since it has not so far proved possible to find a suitable salt for
solubility measurements. This redox reaction might be of interest for
fast high pressure kinetics study. [IrCl 6 ]2- as an oxidant of catechols
and quinols in organic cosolvents may prove a suitable system to follow.
The complete solvation effect on the reactivity of these redox reactions
will become apparent as thermodynamic data for reactant species becomes
available in these aqueous cosolvent mixtures. In the case of ligand
oxidation it might be of interest to prepare a crystal of the complex
discussed above and hence carry out crystal structure determination
which may prove the presence of the third imine moiety within the
hexadentate ligand. Further such determination may be followed by nmr
study.
- 187 -
REFERENCES
1.
2.
3.
4.
5.
6.
M. L. Tobe, "Inorganic Reaction Mechanisms", Nelson, 1972.
K. F. Purcelland J. C. Katz, "Inorganic Chemistry",
Holt-Saunders, 1977
M. J. Blandamer, J. Burgess, P. P. Duce and R. I. Haines, J. Chern. Soc., Dalton Trans., (1980) 2443
M. J. Blandamer, J. Burgess, N. V. Reed and P. Wellings, J. Inorg. Hucl. Chem., 43(1981)2345
W. F. Prou, S. K. Garmestani and R. D. Farina, Inorg. Chem., 20(1981)1297
P. Braun and R. van Eldik, J. Chem. Soc., Chem. Comm., (1985)1349
M. J. Blandamer, J. Burgess, S. J. Hampshere, C. White, R. I. Haines and A. McAuley, Can. J. Chern., (1983)1361
7. R. A. Marcus, J. Phys. Chern., 67(1963)853; 67(1963)2889
8. A. A. Schilt, J. Am. Chern. Soc., 82(1960)3000.
9. L. Maggi, G. Varani, M. F. Manfrin and V. Balzani, Inorg. Chim. Acta, 4(1970)335.
10. M. Schilling, SS(III) Inorg. Project, University of Leicester, 1985
11. E. Mentarsti, E. Pelizzetti and C. Baiocchi, J. Chem. Soc., Dalton Trans., (1977)132
D. F. C. Morris and T. J. Ritter, Intern. J. Chem. Kinet., 11(1979)1081
12. S. F. Kunn, A. M. Lannon, K. C. N. Laranjeira and A. G. Lappin, J. Chern. Soc., Dalton Trans., <1984>1371
13. E. Pelizzetti, University of Torino, Italy 13(a) J. Burgess and E-E. A. Abu-Gharib, Trans. Met. Chern., 9(1984)234
14. J. G. Mohanty, R. P. Singh and A. Charavorty, Inorg. Chem., 14 (1975)2178
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A. G. Lappin, M. C. M. Laranjeira and L. Yonde-Owen, J. Chem. Soc., Dalton Trans., (1981>721
E. R. Gardener, F. M. Mekhail and J. Burgess, Internat. J. Chern. Kinetics, 6(1974)133
P. Kromholz, H. L. Chum, M. A. De Paoli and T. Rabockai, J. Electroanal. Chern .. 51(1974)465
H. L. Chum and P. Krumholz, Inorg. Chem., 3(1974)514
- 188 -
18. D. Soria, M. L. De Castro and H. L. Chum, Inorg. Chern. Acta, 42(1980)121
19. V. L. Goedken, J. Chem. Sac., Chem. Comm., (1972)207
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APPENDICES
APPEIDlX 1
TABLE AI(I) I-ray diffraction data for lFe(cxcage)] (PF&)2
Bond distances (A) for [Fe(cxcage)] (PF&)2
N(l)-Fe N(5)-Fe N(9)-Fe H(11)-C(1) N(2)-C(1) H(21)-C(2) N(4)-C(2) H(31 )-C(3) N(2)-C(3) H(41)-C(4) N(8)-C(4) H(51 )-C(5) N(10)-C(5) H ( 61 ) -c ( 6 ) N(8}-C(6) C(12)-C(11) N(1)-C(11) N(7}-C(12) H(132)-C(13} C(15)A-C(14) C(15)B-C(15)A C(16}-C(15}B C(26)-C(21) C(23)-C(22) H(231)-C(23) C(24)-C(23) H(242)-C(24) H(251)-C(25) C(26)-C(25) H(262)-C(26) C(36)-C(31) C(33)-C(32) H(331)-C(33} C(34}-C(33) H(342)-C(34) H(351)-C(35) C(36)-C(35} H(362)-C(36) N(4)-N(3) N(8)-N(7) N(12)-N(11) F(2)-P(1) F(4)-P(1) F(6)-P(1) F(S)-P(2) F( 10)-P(2) F(12)-P(2)
1.914(6) 1.897(6) 1.930(5) 1.080(11) 1.498(8) 1.080(9) 1.458(10) 1 .080 ( 9 ) 1. 445 (10) 1.080(10) 1. 498 (10) 1.0S0( 12) 1.458(11) 1.080(12) 1 . 476.( 1 1 ) 1.442(9) 1.262(10) 1.265(10) 1.080(11) 1.56(4) 1.18(5) 1.66(3) 1.514(9) 1.511(10) 1.080( 13) 1.477(15) 1 .080 ( 14) 1 .080 ( 19) 1.516( 12) 1 .080 ( 12) 1.509(12) 1.459(10) 1.080(10) 1.488(13) 1.080(12) 1 .080 ( 14) 1.492(13) 1 .080 ( 12) 1.437(7) 1.431(8) 1.433(8) 1.S2fi(8) 1.567(9) 1.559(9) 1.529(9) 1.524(9) 1.533(8)
- 189 -
N(3)-Fe N(7)-Fe N(11)-Fe H( 12)-C( 1) N(4)-C(1) H(22)-C(2) N(6)-C(2) H(32)-C(3) N(6)-C(3) H(42)-C(4) N(10)-C(4) H(52)-C(5) N(12)-C(5) H(62)-C(6) N(12)-C(6) C(16)-C(11) C(13)-C(12) H(131)-C(13) C(14)-C(13) C(15)B-C(14) C(16)-C(15)A C(22)-C(21) N(3)-C(21) N(9)-C(22) H(232)-C(23) H(241 )-C(24) C(25)-C(24) H(252)-C(25) H(261 )-C(26) C(32)-C(31) N ( 11 ) -C ( 3 1 ) N(5)-C(32) H(332)-C(33) H(341)-C(34) C(35)-C(34) H(352)-C(35) H(361)-C(36) N(2)-N(1) N(6)-N(5) N(10)-N(9) F(l)-P(l) F(3)-P(1) F(5)-P(1) F(7)-P(2) F(9)-P(2) F( 11 )-P(2)
1.921(5) 1.89fl(6) 1.910(6) 1.080(9) 1.458(10) 1.080(11) 1.475(8) 1.080(9) 1.470(9) 1 .080 ( 12) 1.405(11) 1.080(10) 1.475(10) 1.080(10) 1.457(12) 1.511(11) 1.479(11) 1.080( 12) 1.467(16) 1.29(3) 1.52(4) 1.425(9) 1.298(8) 1.291(9) 1.080(13) 1.080(20) 1.448(16) 1.080(15) 1.080(12) 1.450(9) 1.273(10) 1.298(9) 1 .oao( 10) 1.080( 13) 1.519(13) 1.080( 13) 1 .080 ( 11 ) 1.434(7) 1.428(7) 1.457(8) 1.528(7) 1.545(8) 1.537(9) 1.559(5) 1.556(6) 1.558(7)
TABLE Al (1) (continued) . .. Bond angles for [Fe (cxcage)] (PFS)2
N(3)-Fe-N(l) 86.7(2) N(S)-Fe-N(l) 86.5(2) N(S)-Fp--N(3) 86.4. (2) N(7)-Fe-N(1) 78.3(2) H(262)-C(26)-C(21) 109.2(8) H(262)-C(26)-C(25) 109.0 N(7)-Fe-N(1) 117.1(3) N(7)-Fe-N(S) 150.7(2) H(262)-C(26)-H(261) 109.5(8) C(36)-C(31)-C(32) 121. S N(q)-Fe-N(l) 149.6(3) N(9)-Fe-N(3) 77.2(2) N(11)-C(31)-C(32) 112.8(6) N(11)-C(31)-C(36) 125.6 N(9)-Fe-N(S) 117.6(3) N(9)-Fe-N(7) 86.2(3) C(33)-C(32)-C(31) 122.3(6) N(5)-C(32)-C(31) 111. 4 N(ll)-Fe-N(l) 117.S(2) N(11)-Fe-N(3) 1S0.5(3) N(5)-C(32)-C(33) 126.3(6) H(331)-C(33)-C(32) 109.0 N(11)-Fe-N(5) 79.1 (2) N(11)-Fe-N(7) 85.9(3) H(332)-C(33)-C(32) 108.9(7) H(332)-C(33)-H(331) 109.5 H(11)-Fe-N(9) 86.9(3) H ( 12) -C ( 1 ) -H ( 11) 1013.5(9) C(34)-C(33)-C(32) 111.5(6) C(34)-C(33)-H(331) 108.9 N(2)-C(1)-H(11) 108.3(6) N(2)-C(1)-H(12) 108.6(6) C(34)-C(33)-H(332) 109.0(8) H(341)-C(34)-C(33) 108.7 Jl(II)-C(1)-H(11) 10R.3(7) N(4)-C( 1 )-H( 17.) 108.5(7) H(342)-C(34)-C(33) 108.5(9) H(342)-C(34)-H(341) 109.5 • N(4)-C(1)-N(2) 113.6(6) H(22)-C(2)-H(21) 109.5(8) C(35)-C(34)-C(33) 113.1(8) . C(35)-C(34)-H(341) 108.5 N(4)-C(2)-H(21) 108.5(7) N(4)-C(2)-H(22) 108.S(7) C(3S)-C(34)-H(342) 108.6(9) H(351)-C(3S)-C(34) 108.7 N(6)-C(2)-H(21) 108.4(7) N(6)-C(2)-H(22) 108.7(7) H(352)-C(3S)-C(34) 108.7(9) H(352)-C(35)-H(351) 109.S ·N(6)-C(2)-N(4) 111.2(5) H(32)-C(3)-H(31) 109.5(7) C(36)-C(35)-C(34) 112.7(7) C(36)-C(35)-H(351) 108.S N(2)-C(3)-H(31) 108.4(7) N(2)-C(3)-H(32) 108.6(8) C(36)-C(35)-H(352) 1Q.8.8(10) C(35)-C(36)-C(31) 110.9 N(6)-C(3)-H(31) 108.6(7) N(6)-C(3)-H(32) 108.4(7) H(361)-C(36)-C(31) 109.0(9) H(361)-C(36)-C(3S) 109.1 .... N(6)-C(3)-N(2) 113.3(S) H(42)-C(4)-H(41) 109.5(9) H(362)-C(36)-C(31) 109.2(8) H(362)-C(36)-C(35) 109.3 N(8)-C(4)-H(41) 108.4(7) N(8)-C(4)-H(42) 108.6(8) H(362)-C(36)-H(361) 109.5 ( 10) C(11)-N(1)-Fe 118.2 N(10)-C(4)-H(41) 108.6(8) N(10)-C(4)-H(42) 108.S(8) N (2) -N ( 1) -Fe 121.5(4) N(2)-N( 1 )-C( 11) 116.0 I: ( 1 0 ) - C ( 4 ) - N ( 8 ) 113.2(6) H(52)-C(S)-H(S1) 109.S(10) C(3)-N(2)-C(1) 109.5(5) N(1)-N(2)-C(1) 106.4 N(10)-C(S)-H(51) 108.2(7) N(10)-C(S)-H(S2) 108.4(8) N( 1)-N(2)-C(3) 112.0(S) C(21 )-N(3)-Fe 119. S \.D N(12)-C(S)-H(51) 108.2(7) N(12)-C(S)-H(S2) 108.3(7) N(4)-N(3)-Fe 123.0(4) N(4)-N(3)-C(21) 1 15.4 0 tl ( 12) - C ( 5 ) - N ( 10) 114.2(7) H(62)-C(6)-H(61) 109.S(8) C(2)-N(II)-C( 1) 109.6(S) N(3)-N(4)-C(1) 111 .0 II (8) -C (6) -H (61 ) 108.8(9) N(8)-C(6)-H(62) 108.9(8) N(3)-N(4)-C(2) 108.1 (S) C(32)-N(S)-Fe 118. 1 1"( 12)-C(6)-H(61) 108.8(9) N(12)-C(6)-H(62) 108.8(9) N(6)-N(5)-Fe 124.7(S) N(6)-N(5}-C(32) 1 15. 1 N( 12}-C(6)-N(8) 112.1(6) C(16)-C(11)-C(12) 122.6(7) C(3)-N(6)-C(2} 110.5(6) N(S)-N(6}-C(2) 110.2 N(1)-C(11)-C(12) 111.8(6) N(1)-C(11)-C(16) 12S.S(6) N(S)-N(6)-C(3) 107.2(5) C(12}-N(7)-Fe 118.3 C(13)-C(12)-C(11) 122.3(7) I" (7) -C ( 12) -C ( 11) 112.6(7) N(8)-N(7}-Fe 124.3(S} N(8}-N(7}-C(12} 115.2 N(7}-C(12)-C(13) 12S.1(6) H(131)-C(13)-C(12) 108.4(8) C(6}-N(8)-C(4) 110.1(6} N(7}-N(8}-C(4) 1 10.6 H(132}-C(13)-C(12) 108.S(9) H(132)-C(13)-H(131) 109.S(9) N(7)-N(8)-C(6) 107.9 ( 6 ) C(22}-N(9)-Fe 118.4 C(14)-C(13)-C(12) 113.1(7) C(14)-C(13)-H(131) 108.9(10) I" (10) -N (9) -Fe 122.6(4) N(10)-N(9}-C(22) 1 lS . 9 C(14)-C(13)-H(132) 108.4(9) C(1S)A-C(14)-C(13) 118.8(17) C(S)-N( 10)-C(4) 110.5(6) N(9)-N(10)-C(4) 108.5 C(15)B-C(14)-C(13) 119.3(18) C(1S)B-C(14)-C(1S)A 47.8(21) N(9)-N( 10)-C(S) 110.3(6) C(31)-N(11)-Fe 117 . 8 C(15)B-C(1S)A-C(14) 54.4(22) C(16)-C(lS)A-C(14) 111.0(26) N ( 12) -N ( 1 1 ) -Fe 124.S(S) N(12)-N(11)-C(31) 114. 8 C( 16)-C( 15)1>.-C( 1:,)B 74.6(27) C(1~)~-C(15)B-C(14) 77.8(2S) C(6)-N(12)-C(S) . 109.7(7) N ( 1 1 ) - N ( 12) -C ( 5 ) 107.4 C(16}-C(15)B-C(14) 117.7(21) C(16)-C(15}B-C(lS)A 62.0(24} N(11)-N(12)-C(6) 110.6(6} F(2)-P(1)-F(1) 89.8 (, C ( 1:» ;"-C ( 16) -C ( 11 ) 111.2(15) C(15)B-C(16)-C(11) 102.1(13) F(3)-P(1)-F(1) 179.1(5) F(3}-P(1}-F(2) 91. 1 (! ~(1':.)f',-C( If,)-C( 1S}A 43.4(113) C(26)-C(21)-C(22) 122.2(6) F(4)-P( 1)-F(1) 88.S(5) F(4)-P(1}-F(2) 175.71 1;(3}-C(21)-C(22) 111.3(6} N(3)-C(21}-C(26) 126.4(6} F(4}-P(1}-F(3) 90.6(S) F(S)-P(1}-F(1) 92.4(; C(21)-C(22)-C(21} 121.7(G) H(9)-C(22)-C(21) 112.8(6) F(5)-P(1)-F(?) 97.2(5) F(5)-P(1)-F(3) 87.7(1 tl(9)-C(22)-C(23) 125.5(6) H(231)-C(23)-C(22) 109.0(8) F(5)-P( 1)-F(4) 86.8(5) F(6}-P( 1 )-F( 1} 89.4 (~ H(?12)-C(23)-C(22) 108.8(8} H(232}-C(23)-H(231} 109.5(9) F(6)-P(1}-F(2} 85.4(5} F(6}-P(1}-F(3) 90. 4( ~ C(24)-C(23)-C(22) 111.7(7) C(24)-C(23)-H(231} 109.1(9) F(6)-P(1}-F(4) 90.7(S) F(6}-P(1)-F(5) 176.9 ( C(24)-C(23)-H(232) 108.8(10) H(241)-C(24)-C(23) 107.4(12) F ( 8 ) - P ( 2 ) - F"(7 ) 91.2(4) F(9)-P(2)-F(7) 178.5 ( H(242)-C(24)-C(23) 107.6(11} H(242)-C(24}-H(241) 109.5(13} F(9}-P{2)-F(8) 88.1(5) F(10)-P(2}-F{7} 89.0 (~ C(25)-C(24)-C(23} 117.5(9) C(25}-C(24)-H(241) 107 . 1 ( 12 ) F( 10)-P(2)-F(8) 179.0(5) F(10)-P(2}-F(9) 91.6(~ C(25)-C(24)-H(242) 107.7(13} H(251)-C(25)-C(24} 107.3(11) F{")-P(2}-F(7) 90.0(3) F(11)-P(2)-F(O) B5.7(~ H(252)-C(25)-C(24) 10B.0(11) H(252)-C(25}-H(251) 109.5(13) F(11)-P(2)-F(9) 88.6(4) F(11)-P(2)-F(10) 93.3« C(26)-C(25)-C(24) 116.9(9) C(26)-C(2S)-H(251) 107.4(10) F( 12}-P(2)-F(7) 89.6(4) F(12)-P(2)-F(R) 95.0 (~ C(~L)-C(25}-H(252) 107.7(10) C(25)-C(26)-C(21) 110.5(6) F(12)-P(2)-F(9) 91.8(4) F(12)-P(2)-F(10) B6.0(5 H(2~1)-C(26)-C(21) 109.3(7) H(261}-C(26}-C(25} 109.5(9) F(12)-P(2)-F(11) 179.2(5) --- ------ ---
TABLE Ai 0) (continued> ... Fractional atODic co-ordinates for [Fe (cxcage>] (PF6)2
Atom x y z
Fe 0.25700(5) 0.10254(7) 0.60726(7) C (3 1 ) 0.1647(4) 0.0504(6) 0.7324(5) C ( 1 ) 0.4171(4) -0.0371(6) 0.5419(5) C(32) 0.2390(4) -0.0146(5) 0.7460(5) H ( 1 1 ) 0.4111(4) 0.0044(6) 0.4853(5) C(33) 0.2514(4) -0.0924(6) 0.8126(5) H ( 12) 0.4683(4) -0.0951(6) 0.5430(5) H(331) 0.2385(4) -0.1793(6) 0.7934(5) C(2) 0.4337(4) -0.0009(6) 0.6791(5) H (332) 0.3138(4) -0.0866(6) 0.8342(5) H (21 ) 0.4853(4) -0.0580(6) 0.6850(5) C(34) 0.1973(6) -0.0629(7) 0.8772(6) H(22) 0.4400(4) 0.06fi7(6) 0.7225(5) H(341) 0.1988(6) -0.1326(7) 0.9190(6) C (3) 0.3448(4) -0. 1538 ( 5 ) 0.6318(5) H(342) 0.2199(6) 0.0146(7) 0.9054(6) H(31) 0.2875(4) -0.1955(5) 0.6412(5) C(35) 0.1097(5) -0.0431(8) 0.8488(6) H(32) 0.3931(4) -0.2167(5) 0.6359(5) H(351) 0.0741(5) -0.0251(8) 0.8990(6) C(4) 0.1549(5) 0.3341(6) 0.5416(5) H(352) 0.0873(5) -0.1203(8) 0.8201(6) H (41) 0.1058(5) 0.3956(6) 0.5359(5) C(36) 0.1006(5) 0.0551(7) 0.7928(6) H(42) 0.2016 (5) 0.3569(6) 0.5020(5) H(361) 0.0414(5) 0.0508(7) 0.7632(6) C(5) 0.1294(4) 0.2984(6) 0.6740(5) H(362) 0.1064(5) 0.1350(7) 0.8248(6) H (51 ) 0.1590(4) 0.2960(6) 0.7318(5) N ( 1 ) 0.2748(3) -0.0276(4) 0.5416(4)
~ H(52) 0.0796 (4 ) 0.3587(6) 0.6736(5) N(2) 0.3416(3) -0.1055(5) 0.5537(4) \.0 C(6) 0.Ofi27(4) 0.1800(7) 0.5757(fi) N(3) 0.3700 (3) 0.1375(4) 0.5977(4) ...."
H (61 ) 0.0444(4) 0.0930(7) 0.5617(6) N(4) 0.4325(3) 0.0512(5) 0.6014(4) H(fi2) 0.0104(4) 0.2359(7) 0.5711(fi) N(5) 0.2908(3) 0.0068(4) 0.6927(4) C ( 1 1 ) 0.2393(4) -0.0300(5) 0.4745(5) N (6) 0.3598(3) -0.0675(5) 0.6937(4) C(12) 0.1811(4) 0.0620(6) 0.4642(5) N(7) 0.1862(3) 0.1348(5) 0.5193(4) C ( 13) 0.1225(5) O. 061'"3 ( 7 ) 0.3961(5) N(8) 0.1222(3) 0.2179(5) 0.5190(4) H ( 131 ) 0.1273(5) 0.1509(7) 0.3690(5) N(9) 0.2607(3) 0.2667(5) 0.6236(4) H ( 132) 0.0621 (5) 0.0563(7) 0.4168(5) N ( 10) 0.1882(4) 0.3385(5) 0.6188(5) C ( 14) 0.1364(7) -0.0215(12) 0.3371(fl) N ( 1 1 ) 0.1616(3) 0.1024(5) 0.6667(4) C(15)A 0.2236(23) -0.071 (3) 0.3309(24) N ( 12) 0.0962(3) 0.1 fl30 (6) 0.6562(4) C(15)R 0.1688(20) -0.1185(24) 0.3592(21) P ( 1 ) 0.89718(13) 0.29080(21) 0.39823(18) C ( 16 ) 0.2568(6) -0.1142(6) 0.4099(6) P(2) 0.37102(12) 0.25178(17) 0.36300(15) C (21 ) 0.3950(4) 0.242fi(6) 0.6069(5) F( 1) 0.9776(3) 0.2592(6) 0.3626(6) C(22) 0.3288(4) 0.3195(6) 0.6157(5) F(2) 0.8970(5) 0.4051 (6) 0.354](5) C(23) 0.33<)7(5) 0.44fl4(6) 0.6127(6) F(3) 0.8162(4) 0.3212(8) 0.4353(6) H(231) 0.3020(5) 0.4878(6) 0.6553(6) F(4) 0.9031(5) 0.1767(7) 0.44fi9(6) H(232) 0.3212(5) 0.4783(6) 0.5547(6) F(5) 0.8476(4) 0.2219(8) 0.3359(5) C(24) 0.4251(7) 0.4820(8) 0.6298(8) F(6) 0.9452(5) 0.3568(8) 0.4650(fi) H (241 ) 0.4352(7) 0.4819(8) 0.6928(A) F(7) 0.3748(3) 0.3847(4) 0.3518(4) H(242) 0.4327(7) 0.5683(8) 0.6077(8) F{A) 0.4240(5) 0.2326(7) 0.2929(5) C(25) 0.4875(6) 0.4108(8) 0.5978(8) F(9) 0.3698(4) 0.1191(5) 0.3747(5) H(251) 0.4844(6) 0.4230(8) 0.5350(8) F (10) 0.3194(5) 0.2705(8) 0.4338(5) H(252) 0.5456(6) 0.4405(8) 0.6216(8) F ( 1 1 ) 0.4524(3) 0.2557(6) 0.4131(5) C(26) o 4825(4) 0.2826(6) 0.6130(5) Fe 12) 0.2905(4) 0.2489(6) 0.3145(6) H(2F;1) 0.5168(4) 0.2370(6) 0.5705(5) H(262) 0.5076(4) 0.2648(6) 0.6713(5)
TABLE Al (1) (continued) ... Atollic thermal panmeers (x10-4) for [Fe (cxcage)] (PF6)2
Atom U or U 1 1 U22 U33 U23 U13 U12
Fe 328(5) 275(5) 834(9) 20(5) 66(4) 16 ( 4 ) C (31 ) 406 (36) 419(38) 847(65) -40(37) 172(35) -3 (31 C ( 1 ) 396(36) 503(43) 929(68) 26(39) 155(36) 130(31) C(32) 465(38) 268(32) 798(60) -3(33) 81(35) -20 (21 H ( 11 ) 500(0) C(33) 622(45) 367(38) 892(69) 19(39) 142(41) -30(3-H ( 12) 500(0) H(331) 500(0) C(2) 318(34) 467(39) 980(69) 76(39) 7(35) 78(30) H (332) 500(0) H (21 ) 500(0) C(34) 986(69) 611(52) 952(87) 136(48) 184(56) 97 (4' H(22) 500(0) H(341) 500(0) c. ( 3 ) 467(39) 323(34) 945(64) 89(36) 97(37) 83(30) H(342) "500(0) H ( 31 ) 500(0) C(35 ) 831(66) 871(66) 995(87) 101(57) 401(55) -85 (5: H(:12) 500(0) H(351) 500(0) C(II) 646(48) 547(4<l) 940(71) 92(42) 11 (43) 228(38) H(352) 500(0) H ( 41 ) 500(0) C(36) 672(52) 712(55) 1048(83) 46(52) 305(48) 62 (4: H(42) 500(0) H061 ) 500(0) C( 5) 635(4fi) 535(45) 967(72) -63(43) 52(43) 289(38) H(362) 500(0) H ( 5 1 ) 500(0) N ( 1) 391(30) 327(30) 806(50) 26(28) 21(28) -10(24
'.D H(5?) 500(0) N(2) 425(30) 338(29) 950(55) 33(32) 63(30) 85(26 N C(fi) 403(39) 723(56) 1250(83) 74(50) 95(42) 163(38) N(3) 375(29) 332(29) 793(51) -1(28) 1( 28) -36(23
H ( 6 1 ) 500(0) N(4) 329(29) 442(33) 908(54) 59(31) 86(29) 14(25 H(62) 500(0) N(5) 423(30) 327(29) 772(49) 27(28) 139(28) 48(24 C ( 11 ) 475(39) 354(36) 835(63) 24(34) 8(36) -109(30) N(6) 452 (32) 341 ( 30) 787(52) 65(28) 94(29) 140(25 C ( 12) 406(37) 375(36) 1000(70) 3308) 29(39) -79(30) N(7) 404 ( 3 1 ) 398(32) 764(50) 72 (28) 2(29) 62(24 C ( 13) 586(47) 636(50) 896(76) 18(45) -81(45) -48(38) N(8) 452(34) 582(41) 1013(61) 77(36) 9 (33) 157(30 H ( 13 1 ) 500(0) N(9) 570(36) 297(30) 987(60) -45(31) 77(34) 91 (27 H(132) 500(0) N ( 10) 700(42) 352(33) 1161 (67) 27(35) -12 (39) 173(30 C ( 14) 1212(98) 1437(115) 1208(115) -387(87) -228(75) 368(84) N ( 11 ) 329(28) 488(34) 866(54) 103(33) 119(29) 43(27 C(15)'; 60(J(363) 660(260) 395(299) 16(214) 636(254) -343(223) N ( 12) 459(34) 700(45) 968(61) 78(37) 97(34) 219(32 C ( 15) B 952(255) 567(163) 1034(240) -265(139) -226(203) 109(152) P ( 1 ) 496(12) 692(15) 1148(25) 97 ( 15) -35(13) -62(11 C ( 16) 1139(69) 422(44) 826(79) -135(44) -93(57) 67(43) P(2) 523(12) 431(11) 901(21) 124(11) 75 ( 12) 21 (9 ) C (21 ) 429(36) 423(38) 762(62) 24(35) -40(35) -99(31) F (1) 782(39) 1274(58) 2844(113) 427(62) 627(51) 159(38 C(22) 672(46) 361(38) 731(63) 34(35) -82(40) -128(34) F(2) 1843(69) 1036(51) 1687(7fI) 517(46) -14(54) 459('17 C(23) '312(fi2) 340(40) 1112(82) -14(42) -118(52) - 1 60 (41 ) F (3) 985(48) 2509(101) 1910(93) -533(70) 544(49) 58(57 H(231) 500(0) F(4) 1911(77) 1530(72) 2189(103) 1082(68) -229(66) -513(59 H(232) 500(0) F(5) 1180(51) 2075(83) 1770(87) -829(63) 177(50) -547(54 C':(24) 1544(109) 469(57) 1790( 132) -8(66) -166(91) -372(67) F(6) 1662(71) 1757(77) 1991(97) -184(63) -580(65) -510(60 H(241) 500(0) F(7) 1042(42) 4!i6(30) 2165(76) 167(35) -89(42) 21 (28 H(242) 500(0) F (8) 2185(84) 1743(79) 1513(87) 189(57) 881(64) 866(65 C(25) 927(72) 730(66) 1598(117) 242(66) -35(71) -510(59) F(9) 1627(64) 525(35) 2.4 8 9 (93 ) 4fil(43) -315(59) -213(36 H(251) 500(0) F ( 10) 1528(62) 2352(98) 1446(83) 368(63) 660(54) 705(63 H(252) 500(0) F ( 1 1 ) 794 (38) 1015(47) 2470(92) -29(48) -551(45) 156(34 C(2li) 501 ( 42) 561(48) 987(73) 54(43) 52(41) -216(35) F ( 12) 1354(56) 1076(50) 2508(100) 35(53) -916(59) -160(42 H(261) 500(0) H(262) 500(0)
TABLE Al{II) X-ray diffraction data for [Pe(gDd)31 (BF4
)2
BrJl,d Lengths
N-Fe C(2)-N C(-1)-C(1) H(-1)-C(1)
HCJ)-CC2)
"1.9S2<::2) "1 . '~59 ('~) 1 . '123 ( 6) 0.97'1(29) 0.91(4)
C(1)-N F(1)-B
--I, , O. DODO, H(2)-C(2) H (I,) - C en
Bond AngLes
C(1)-N-Fe C (2) --N-C (-1)
N-Fe-N N-Fe-N H(2)-C(2)-N H(3)-C(2)-H(2) H(4)-C(2)-H(2)
114.6(2) C(?)-N-Fe "119.0(3) N-Fe-N 95.4(1) N-Fe-N 89.4(1) H(l)-C(l)-N -108.0 (-19) H (3) - C (2) -1'1 104.5(29) H(4)-C(2)-N 113(3) H(4)-C(2)-H(3)
Atomic thermaL parameters (M"1U~*4)
Atom U or U-l1 U--;>--:> U33 U23
Fe 42S (L,) 425(4) 62lH 6) O(m C(D 597 <'18) 47"1(15) 835(22) -42 (1'~) C(2) 856(26) 682(22) 81:::-(26) "197 (2m N 554(14) /~98(1"1) 685(4) 37(10) H(l) 720(84) H(2) 892(103) H(3) 991(1"15) H(4) 878 (11'~) B 623(20) 623(20) 88[J( 42) 0(0) F(l) 1092 (17) 1092 (-17) 80't- (22) 0(0) F(2) 853 <-15) 13,4 (20) 1399(19) 28 (-16)
FractionaL atomic co-or'd i nates
Atom M y z
tt -1) 8·QPgQ~i9) • .Lt:> t 8: 9~99((59) 8:~j~~~~~J
1 • 27 ~ (3) 1 .365 ( 7,
0.0000, 1.00UO 0.98(,3, 0.92 CII
un
O(m 171(-15) 132 (20) 136(12)
0(0) 0(0)
288(14)
126.2<::::) 95.4(1) 80.0(t) -1-18 • 7 ( 1 7)
-110.6(22) 108.5("21) 1"12(3)
U12
2-13 (2)
-167 (1'.)
410(21) 264('12)
312 (-10) 5'~6(81
590 (1',)
C(2) 0.0',46(5) 0.239' .. (5) 0.19"181<-11) 1'1 0.1 OlVt-6 ("25) 0.19507(24) O. 222'~2 (6) H(l) 0.300(3) 0.396(3) 0.2231(8) H(2) -0.042(4) O. 25"1 (I~) 0.2008(8) H(3) 0.116(4) 0.336(5) 0.1835 (10) H(4) 0.015(1t-) 0.16't(4) o . -1 7 It-"! ( 8 ) B O. 3:~:S33 (0) 0.66667(0) 0.26208(18) F(l) 0.33333 (0) 0.66667 (0) O. 22'~99 (-10) F(2) -0."19071(27) 0.3650(3) o . 226"llt (7)
Non-bonded Contacts
C(l) ... Fe 2.739 C(2) ... Fe 3.050 H(2) ..• Fe 3.254 H(L,).~.Fe 3.194 H (1) ••• 1'1 1.935 H(2) ... 1'1 -1.989
HCJ) ... 1'1 1.973 HU,) ..• N 1.957
N ... N 2.889 --. ~, 0.0000, 0.0000, 0.0000
N ... N 2.889 3, 0.0000, O.OOOU, O. OfI(JO 2.511 -,~ , 0.0000, 0.0000, l.OOOO N ..• N 2.278 -', , o.nooo, 0.0000, 1.0nOO C (-1) ••• 1'1 2.747 -5, 0.0000, 0.0000, 1.00g0 N ... N "1.3',9 -II, -"l.ooon, O.OUOO, 1 • nooo F(2) •.• B 2.206 -' .. , -1.0000, 0.0000, 1.0000 F(2) •.. F(1)
HC?) ... F(2) 2.4J4 H(l) ••. F(l) 2.522 l.OOOO, 0.0000 2.2-1J -.., 0.0000, F(2) ... F(2) ~,
0.0000, O.OfIUU J, -1.0000, F(2) •.• F(2) 2.2"1J 2. '.12
C(2) .•. C(l) 2.355 H(3) ... C(1) 0.0000, 1.0000 2."1'.2 -4, 0.0000, H(l) •.. C(-1)
H (J) ... H (1) 2.169 C("2) ... H(l) 2.48J
1.578 H(3) ••. H(2) -1.491 H(4) ... H(2)
H(4) ... HCJ) "1.521
- 1')1 -
TABLE Al(lll) X-ray diffraction data for [Fe(bmi)3](CIO~)2
Bond lengths
N(l)-Fe C(2)-N(1) C(l)-C(l) H(21)-C(2) H(23)-C(2) H(32)-C(3) 0(1)-Cl 0(3)-Cl 0(3)-0(3) 0(3)-0(3) 0(4)-0(4) 0(4)-0(4)
1.956(2) 1.474(4) 1.477(6) 1.081(16) 1.071(15) 1.074(16) 1.372(7) 1.363(14) 0.951(23) 0.951(23) 1.32(5) 1.32(5)
C(l)-N(l) C(3)-C(1) -4, 0.0000,
H(22)-C(2) H(31)-C(3) H(33)-C(3) 0(2)-Cl 0(4)-CL
2, 1.0000, 3, 0.0000, 2, 1.0000, 3, 0.0000,
0.0000,
1.0000, 1.0000, 1.0000, 1.0000,
Bond Angles
C(1)-N(1)-Fe C(2)-N(1)-C(1) N(1)-Fe-N(1) N(l)-Fe-N(l) H(21)-C(2)-N(1) H(22)-C(2)-H(21) H(23)-C(2)-H(21) H(31)-C(3)-C(1) H(32)-C(3)-H(31) H(33)-C(3)-H(31) 0(2)-Cl-0(1) O(3)-Cl-0(2) 0(4)-Cl-O(2) O(l)-CL-O(l) 0(2)-CL-O(2)
,0(3)-CL-0(3) 0(4)-CL-O(4)
116.9(2) 119.9(2) 96.1(1) 88.6(1) 107.2(15) 109.6(19) 108.8(19) 108.5(11) 108.6(18) 109.3(18) 145.6(10) 81.6(12) 88.6(16) 102.8(6) 98.9(15) 40.8(10) 57.4(21)
C(2)-N(1)-Fe N(1)-Fe-N(1) N(1)-Fe-N(1) C(3)-C(1)-N(1) H(22)-C(2'-N(1) H(23)-C(2)-N(1) H(23)-C(2)-H(22) H(32)-C(3)-C(1) H(33)-C(3)-C(1) H(33)-C(3)-H(32) 0(3)-CL-0(1) O(4)-CL-0(1) 0(4)-CL-O(3) O(l)-CL-O(l) 0(2)-CL-0(2) O(3)-CL-O(3) 0(4)-CL-O(4)
ntomic thermaL parameters (x10~*4)
Atom U or U11 U22 U33 U23
Fe 322(4) 322(4) 326(5) 0(0)
C(1) 334(15) 402(17) 543(18) 128(14)
C(2) 569(22) 648(25) 574(20) -2(19)
C(3) 396(19) 553(23) 999(34) 171(22)
N(1) 414(14) 391(15) 420(12) 46(12)
H(21) 1343(116) H(22) 1343(116) H(23) 1343(116) H(31) 1668(179) H(32) 1668(179) H(33) 1668(179)
0(0) 523(9)
U13
0(0) 46(14)
194(18) 158(19) 34(10)
0(0)
1.292(4) 1.490(4) 1.0000
1.042(15) 1.085(16) 1.063(16) 1.401(13) 1.38(3) 0.0000 0.0000 0.0000 0.0000
123.1(2) 96.1(1) 79.5(2) 126.1(3) 110.1(16) 108.4(15) 112.6(19) 109.5(12) 110.3(12) 110.6(18) 91.9(7) 97.7(15) 169.8(14) 102.8(6) 98.9(15) 40.8(10) 57.4(21)
U12
161(2) 185(15) 284(19 1
163(18) 211(14)
299(3) 599(6) CL 599(6) -658(101) 8S~(89)
0(1) 2604(152) 755(57) 1810(92) -5J7(63) 1533(225) 1535(155)
0(2) 2645(240) 1274(119) 3500(261) 458(171)
O(J) 765(50) 0(4) 725(82)
- lq~ -
TABLE Al(III) (continued) ...
FractionaL atomic co-ordinates
Atom
Fe C(l) C(2) cc:n tt (1 )
H (21) H(22) H(23) HCH) H(32) H(33) CL (I (1)
0(2) O(3)
0(4)
x
0.00000 (0) 0.29'.5(3) 0.2293 (it)
0.4'.96 (4) 0.1932(3) 0.129 (3)
0.323(3) 0.249(5) 0.5363 (13) 0.46l(3) O. '.66 (3)
0.33333(0) 0.2916(t5) 0.2811(25) 0.3196(27) 0.345(5)
y z
0.00000 (0) 0.25000(0) 0.2427 (I.) o . 2 -116 2 (21 ) 0.0370(5) 0.10977 (2'.) 0.3437(5) 0.17272(29) O. -1071 (3) 0.18438(15)
-0.015 ('t) 0.0680(18) 0.121(3) o . 07 '19 ( 19)
-0.05t(3) 0.1343(19) 0.363(5) 0.2215(20) 0.'.528(23) 0.155(3) 0.290(4) 0.-ll69(20) 0.66667 (0) 0.07449(-10) 0.5276(t1) 0.0357 (7)
0.7546 (-17) 0.-1t85( 19) 0."6063(18) O. -1562 (10) 0.748(4) -0.0007 (21)
Non-bonded Contacts
C (-1) ... Fe 2.790 C(2) ... Fe H(2-1) ... Fe 3.099 H(2t) ... NO) H (22) ... N (-1) 2.078 H (23) ... N ( t ) C(3) ... N(-1) 2. '.81 H ( 33:' ... N ( 1 ) N (-1) ... N (l) 2.909 2, 0.0000, 0.0000, N (-1) • • • N (-1) 2.909 3, 0.0000, 0.0000, N(t) ... N(1) 2.501 -4, 0.0000, 0.0000, C(l) ... N(1) 2.316 -4, 0.0000, 0.0000, N(1) ... N(l) 2.733 -6, 0.0000, 0.0000, C(2) ... C(1) 2.397 H <::22) ... C ( 1 ) H ( 31) . . . C <: -1 ) 2.103 H (32) ... C (-1)
C(3) ... CO) 2.577 -4, 0.0000, 0.0000, H <: 33) . . . C (-1) 2. -109 C(3) ... CC2) H(33) ... C(2) 2.457 H(22) ... H(21) H<:23) ... HC~1) l.7'.9 H(23) ... H(22) C (3) ... H (22:> 2.1./.4 H(33) ... H(2.2) o ('1) . . . H ( 22 ) 2.454 -3, 0.0000, 0.0000, 0(2) ... H(23) 2.l:33 1, 0.0000, 1.0000, 0(4) ... C(3) 3.009 -") k, 0.0000, 1.0000, C(3) .•. C(3) 2.992 -4, 0.0000, 0.0000, H(32) ... H(31) 1.753 H(33) ..• H(3l) 0(3) ... H(31) 2.116 -4, 0.0000, 0.0000, o ( 3) . . . H (3-1) 2. '.87 -6, 0.0000, '1.0000, H(33) ... H(32) 1.757 0(3) ... H(32) 0(4) ... H(33) 2. -ltD -2, O.OOElO, 1.0000, 0('0 ... H(33) 2.538 -3, 0.0000, 0.0000, 0(2) ... 0(-1) 2.649 0(3) ... (I (-1)
(I ('1) ... 0 (1) 2.145 --, 1.0000, 1.0000, .:., 0(2) ... 0(-1) 2.258 --, 1.0000, 1.0000, .:. , 0(3) ... 0(1) 2. '.09 ~) 1.0000, 1.0000, -, 0(1.) ... OU) 1.326 .-, '1.0000, '1.0000, ..::. , 0(1:> ... 0 ( 1 ) 2.145 3, 0.0000, 1.0000,
0('.2) ... (0) 1.344 3, 0.0000, 1.0000,
0(3) ... 0(1) 2.447 3, 0.0000, 1.0000,
0(4) ... 0(1) 1.109 3, 0.0000, 1.0000,
(I ( 't) ••• 0 ( 1 ) 2.073 0(3) ... 0(2)
0(2) ... 0(2) 2.129 ..... 1.0000, 1.0000, .::.,
O(;J) ••• 0(2) 0.964 .-. 1.0000, 1.0000, ..t..,
o (It) ••• 0(2) 2.651 ..... 1.0000, 1.0000, .::., 0(2) ... 0(2)- 2.129 3, 0.0000, 1.0000,
0(3) ... 0('2) 1.497 3, 0.0000, 1.0000,
(l(/.) ••• 0(2:' 2.330 3, 0.0000, 1.0000,
0('.) ... oe::) -1.94::' 0(4) ... 0(3)
o ('t) .•. 0 (3) 2.505 ") 1.0000, 1.0000, ..... , 0(4) ... 0(3) 2.480 3, 0.0000, -1.0000,
- Fl~ -
3.024 2.070 2.078 2.621
0.0000 0.0000 1.0000 -1.0000 1.0000 2.502 2.107
1.0000 2.907 1.736 1.758 1.700
0.0000 0.0000 0.0000 1.0000
'1.752 1.0000 1.0000 2.560
0.0000 0.0000
-1.966 0.0000 0.0000 0.0000 0.0000 0.0000 0.0000 0.0000 0.0000 1.806
O.OOOG 0.0000 0.0000 0.0000 0.0000 0.0000
2.7=-'.2 0.0000 0.0000
~
\.0 (J\
-:J:J,....~
5pp.,.-,
2pt:'....,
--------- --_. -------- _._-- -----~--- ----...".-_._-----. -.. - ----,-- ------~--; :;':'
... ~ ~-. -'; <; ,~,
--- --- -- --'-"_.--'-'-' _ .. --' ---:--:-7"-"-
FIGURE Al(I) ~--.--.-:..::~:~--- ----------. -- c· :.. - -. __ .. - -.. -----------
Proton nmr spectrum in da-acetonitrile of [Fe(gDd)a]2+ at 298.2 K ~
- --- -------------- ----- -.:...----~ . - .. --L -' ----..:'--...!.-:.:..:.
~ ---:--:---. ---.-- ------..-.-------r - .. ---~.-.~=-: .. :- - -- .. -~-. ~ - --"~--- --~~--~:-.::....--
r- -----~-.J. : ..
- ----. -. ------------- --- - ~-.:- -....:.---=------- f----,--.--. ~ -- :-: - ---- -- -------. -.( :.... .. -.- .!- - . ., .. - _._--------" :..:.- ..
: : . ~. ~ ,------- -- ---:.-----~-- -----..- - - --
_t. __ .: .
--- -t--·
_. - '-,- _.- -- --. -_._;---. ----- - .. .. - ., :-: .. = .:; ~
H.. H... . , -------:-. __ . - ---. -,-.~~.,.....- -"--T-
-. -- .... -----1:--·-
1---- " / C-c I \.
Ie-I I-1Ie
I.
- ----------------
, ------.-~-,.:..;--. -: ._f---' .-------
.:. :-.-::~':' CIla _ :.... __ 0 -:-,;:-: ;- ~L __ _ _.- - - I --'--'-.i_-. --'-~
i .!
----1 . .-.- -,,--:: :;'--:-~~:
- ~)~:'i ~---::. -';~_-i:~ I ----- - ;'- ----::... .. :---
~-_r~-~IL. ---- .... III
..... __ ... --'':--':"""":". =:::-:";::--. "
. _____ : __ ~...l.
- • ___ ..... _. ____ 0 __ 0
- -
.. --- _ ... -. . .
--------.--~---
• ___ . - '-0'>. __
--------'-
_' _ .. I .,
I
-=-'- \:C-FF ~--···-~-~-·-l . --·-i
______ -_j ___ . __ ~.-: _1 ,- ,
._ .... _ - ___ --1._ --:-i
:1....;.
,.: ... : . .:..;- .. -- -----~-------:---.- """';-
------, . • .1 -:~ --'.-1
--.----~-,.~----- - -. ; ... ~:~~.-.:~-~:--.j;~) '- ~ - - ----
• • I __ : __
,....: .. -:-,----~~.-:- .. . -=~~=- -~- . i -r--
; ~~.-
.! .' . -1
---.-.:
---'-- -- -----------------------------------
rr...., I~ 'i:l ,~ b r :., j :J
\,() -..J
-;C';)r-
5:'r:-m
2rmm
.. ' :'" ,---__ -- -____ ._. _·_ .. --_--__ --0- .~_
:._;. ,~. --:;:.~:: -----.. ---'-'---'--c---'-----'- -',---:-:----------... --
r)' ~ ;-'- ::,' J .--------- ----. _., - 4>':'~
------_._-------------_._--. ------- ----=--- ...!_---_. -
1,'.'
r--
FIGURE Al <I 1)
.-..
'. r:---::=-'-
- ~~, ;---=.:.. ..;..-
-'-----_.- ---_._-'-'---. -. ..... ..
~--.- .... - -- ._-----.... --.
-:.., r,
..;." ~
,-.-.-----------------~
; ; ----- - •. --:-r-----..
. -~---:----:-----:--.
_._-----------------. Proton IllI1r spectrum in d3-acetonitrile of [Fe (bmi) 3] 2+ at 298.2 K
, . .
-----... _- .•• '-.-.:-"7"" ~
L ..
; :~ ~ i r-----'---~--.-- -------.. r -.--- -. -----.-----~--------.' . . , ~ - .-:- -~ -~ - .. ~
_.---'- - -- ·--:-1' - . ',,"-
-. I .
:-'.--"
,------
~ .. -=-_1
. -- -'-r--;---:-'-: ~ --" -'-' - -:,---. '-'-. -'- --, --:-----:-------:--
._..1-_ ·.:i
--~:...:.......:.-... -. ,
--'~--.-~ -'-__ ---' ___ ------=.:....l _--'-----'_.
._- -- --_. --
](e ](e
'" / C-c ! \.
Ie-I I-lie ~--- .-. -
_. __ .. -
__ . ___ . _____ - - •. _. __ 0::-- ___ . ~. __ ... __
- .--- -.... - .-- -'-'"- --. --- - ---.. -. -.--- -.----.--------------i--.-
__ .1_ I
"
._ o. ---------. --;---.-----o __ •• ... -'_.""-_ .. -.-
,. '---; ----, ,- - ----_._----:----_ . j.
,: " CI6I.' :.+:-':"".""';'-i ___ . ';; -: --~- '-- _ .. ---.
. __ ._--- ._--_._---=------ ,:: :1 .. ,...1' ...L,_. __ ._:_;_., ----.-...:..,--: -_. -
" -- ----.--.- :--r.-== -:.--"----~~ - ... .
'-- -_._--- --~--' _____ ·_·1;-,· ______ , ___
---,---~, ---- .. --. ---------'-,--lf~
""':.-;-. ~:-.. - ----:--.
- --_._'- -...:-- ~- -. - ... _'- - -. ----- ---. --. ..:-. , .
-T-'-
"
-"-.;......:......;--_. --' -'-----_. --'-----------------------
ppm (0) 1.'\ " [< ~ b ::, :'
\.0 co
FIGURE Al<III>
Proton nmr spectrum in d3-acetonitrile of (Fe(cDd)3]2+ at 298.2 K
elk
cyclohexane protons n C-c / \.
Ie-I I-JIe
I : Z 3 ' L ~ E ::: ; . ~ ~ <: r. ~ 2 ' 2 ~ 2 , II B E 2 6 t: 4 ;' ; , 2 ~ r. (1 , s" s e " f\ " ;' " ' G
~
\0 \0
t'
---'
I~ \ ------ \...
~. L r
. ,:. ....~ r;"~
#' .~. _':..
"
FIGURE Al (1) ...... :
.'" .. ',: -.::
" . . .
" Proton nmr spectrum in da-acetonitrile of [Fe(cxcage)]2+ at 298.2 K .. ,: .....
protons from CaB3 rings cyclohexane protons
'--- ---- -.l
~ r :J -, ""'-~-...,..----"-~--r--__ ...--__ ':"''''''~,_-_---~-~-- I ,---r- - p-----.--.----, r-' - -r ~--~ ~--.-~-- ~------ -
'E S]c s.rt Abr LB' '.Ar '}\~ "if' t~ 1.L( ~.q ~?t !rr ~~(' .le :.~c ~~r rr r:.r :(~ 0;'''1
~, ,'r 'd t,
N 0 0
[Fe(gDi)s] {Cl04)2 [Fe (phen)s] (Cl04)2 [Fe{bipy)s] (Cl04)2 ------------------------------------------------------------------------------------------------------------------------Vtt Urea o 10 20 30 40 o 10 20 30 40 o 10 20 30 4 ------------------------------------------------------------------------------------------------------------------------Absorbance of SS
103So1. lmol run-3
om~e{Salt)/kJ DOl- 1
(a)
Volt Glycerol
Absorbance of SS
103S0 1./mol dnr3
oM~e(Salt)/kJ mol- 1
(b)
424 624 11'1'1 1674 2159 9.48 16.2 29.6 47.7 '15.2 18.56 30.38
49.3 72.6 137 195 250 0.824 1. 41 2.57 4.15 6.54 2.13 3.49
-2.87 -7.59 -10.2 -12.1 -3.'19 -8.47 -12.01 -15.40 -3.66
[Fe (gmi) 3] (Cl04) 2 [Fe(bipY)3] (CI04)2
o 10 20 30 40 50 60 o 10 20 30
468 423 369 327 279 231 189 18.1 19.0 18.7 18.4-
54-.4 49.2 42.9 38.0 32.4 26.8 22.0 2.08 2.19 2.15 2.12
+0.75 +1.77 +2.66 +3.85 +5.25 +6.74 -0.36 -0.26 -0.11
APPElIDII 2
TABLE A2 <I)
Solubility and derived transfer chemical potentials for Fe(!!) diimine perchlorate salts in (a) urea and (b) glycerol binary aqueous ndxtures at 298.2 K.
40
18.0
2.07
+0.04-
47.85 71. 7 10
5.50 8.24 1
-7.04 -10.05 -1
50 60
18.6 21. 1
2.13 2.43
-0.17 -1.13
TABLE A2 <ii)
Derivation of transfer chemical potentials (kJ mol- 1 ) f OH- ion in aqueous ethanol and aqueous acetone. at 298~~ K.
Ethanol Acetone
Wt~ A B Wt~ A B
0 0 3.77 0.10 +0.57 3.79 0.11 +0.63 7.27 0.17 +0.97 7.30 0.21 +1.20
10.5 0.23 +1. 31 10.6 0.30 +1. 71 13.5 0.29 +1.65 13.6 0.42 +2.39 16.4 0.35 +1.99 16.5 0.51 +2.91 21. 5 0.43 +2.45 21.6 0.68 +3.88 26.2 0.49 +2.79 26.2 0.84 +4.79 30.1 0.54 +3.08 30.2 1. 00 +5.70 33.7 0.61 +3.48 33.9 1.12 +6.38 36.9 0.65 +3.71 37.1 1.24 +7.07 39.9 0.71 +4.05 40.1 1. 36 +7.75 42.6 0.75 +4.27 42.8 1. 48 +8.44 45.1 0.79 +4.50 45.3 1. 58 +9.01 47.3 0.82 +4.67 47.5 1.69 +9.63 49.3 0.85 +4.84 49.6 1. 78 +10.15 51.3 0.89 +5.07
A = pKw (mix) - pK...,(aq) calculated from ref. 30. Capter 4 B = 6'mp9(H+On-) = RTln<10HpK...,(mix) - pK..,(aq)]
Ethanol Acetone Wt~ --------------------- --------------------------------------
5 10 +1. 30 +0.50 15 20 +2.28 +0.50 25 30 +3.05 -0.20 35 40 +4.02 -1. 80 45 50 +4.95 -5.50
(a) ref. 16.27; (b)
011-
to.80
+1. 78
+3.25
+5.82
+8.45
ref. 32
Hel (b)
+0.17 +0.33 +0.50 +0.70 +0.95 +1. 30 +1.55 +1. 90 +2.30
(c)
- 201 -
OH-(c)
+1.25 -1.08 to.80 tl.88 +2.55 -2.22 tl.65 t3.87 t4.00 -3.50 t2.55 t6.05 +5.52 -4.82 +3.50 t8.32 +7.27 -6.32 t4.50 +10.82 +8.95 -7.65 t5.52 +13.17
+10.70 -9.15 t6.60 t15.75 t12.30 -10.40 t7.70 +18.10 t13.90 -11. 60 +8.95 t20.55
ref. 29. Cbapter 4
"D o "D
10ppm ""OOHz 75 Q! ~\..I 5
Sppm 50 375 ::=::39<. lSO:FF --+--=7
2ppm 180. ~-=::
150 - 20 90 -00 30
APPERDII 3
FIGURE A3 (1) CI6 -
Proton nmr spectrum in dG-acetone of [Fe(tsbh)2]2+ at 298.2 K
---_::::LO_- .
1=--
1'-=. ~-
~
, . !
Xe ~
~ J
. ~~ , ~ prri~~~;,p~~g H'-I
t== _:L:_
.---r-- - -~,.......:..:...
H -~~ ..
F-'-:-'--' -.
=:-=-'== ... -- X'~ . . ~
~ ~ --::'.
I , ! ' , I . , I , I , I I I ' I I • , • I I I I I I I I r I I ' ! I I , I I
ppm (0) 10 9 8 7 6
J(e
H
=
, , I I I , , , i . I I I I I , I, ! I , I I ! I I I , I ,
5 4 3 2
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Proton nmr spectrum in do-acetone and ~O of [Fe(tsbh)2]2- at 298.2 K =+=== I ±:=~ = ' = . E .C16 ,·
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II FIGURE A.3 <I II)
__ Proton nmr spectrum in d&- acetone of [Fe(~)3J ~- at
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APPEIDIX 4
TABLE A4
Solubility and transfer cheBdcal potentials for hydroquinone in aqueous cosolvents at 298.2 K
]leaH EtOH i-PrOH Vol -------- -------~ A B A B A B
0 0.650 0.651 0.654 10 0.841 -0.64 0.950 -0.93 0.964 -0.94 20 0.988 -1. 04 1. 35'1 -1. 82 1. 324 -1. '15 30 1. 24-4 -1. 61 1. 35'1 -1. 82 1. 324 -2.49 40 1. 64'1 -2.31 2.212 -3.03 1. 896 -2.64 50 2.012 -2.80 2.5'19 -3.41 2.280 -3.09 60 2.380 -3.22 2.'115 -3.54 2.125 -2.92 70 2.604 -3.44 2.983 -3.'1'1 2.201 -3.01 80 2.935 -3.74 2.855 -3.66 I 2.096 -2.88 90 2.948 -3.75 3.093 -3.86 I 1.980 -2.'14
100 3.034 -3.82 2.880 -3.68 I 1.615 -2.24
A = Solubility/mol ~3 B = o~~e/kJ mol- 1
t-BuOH Acetone -----
A B A B
0.62'1 0.612 0.929 -0.9'1 1.161 -1.58 0.540 +0.36 1. 863 -2. '16 0.619 +0.03 2.519 -3.50 : 0.'!'1'1 -0.53 3.158 -4.06 0.908 -0.92 3.111 -4.03 1. 038 -1. 25 3.078 -4.00 1. 093 -1. 38 3.108 -4.03 1. 098 -1. 39 2.9'13 -3.92 0.'113 -0.31 2.556 -3.54 0.184 +2.9'1 1. 398 -2.04