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Structure and behavior of organic analytical reagents
VOL.
ANALYTICA CHIMICA ACTA
11 (rg51)
STRUCTURE VI.
HEATS
AND
20 I
BEHAVIOR OF ORGANIC REAGENTS
ANALYTICAL
AND ENTROPIES OF FORMATION OF DIVALENT METAL CHELATES OF z- AND a-METHYL-S-HYDROXY QUINOLINE’
SEVERAL
by W. Cort/rlbutlon
D.
JOHNSTONS
of llrc Uepnrtnaent
of Chenaisiry,
AND
H.
FREISER
Untvcvsity
01
Ptltsburgh,
Pa.
(U.S.
A .)
It has been well established that any substitution in the z-position of 8-hydroxyquinoline prevents the precipitation of an aluminum complex. This effect was first noted by MBRRITT AND WALKBH~ in the case of a-methyl-8-hydroxyquinoline. IRVINC~ and co-workers established the general na turc of this effect by a-study of the analytical behavior of z,4-dimcthyl- and z-phenyl-8-hydroxyyuinoline as and g-hydroxy-I,z,3,+tetrahydroacridine. They well as x-hydroxyacridine reasoned that this effect was due to the steric hindrance provided by the group in the z-position. ALBERTA has noted that any substituent in the z-position of 8-hydroxyquinoline lowered its antibacterial action. In a previous paper the authors O have reported formation constants for a number of metal chelates of 8-hydroxyquinoline and z-methyl-8-hydroxyquinoline. It was found that the chelates of z-methyl-8-hydroxyquinoline were slightly less However, the nickel(I1) chelate of stable than those of 8-hydroxyquinoline. z-methyl-8-hydroxyquinoline was considerably less stable than the corresponding This unusual effect was so pronounced th t the chelate of 8-hydroxyquinolinc. 7 ated usual metal sequence Cu(I1) > Ni(II) > Co(II), %n(II) > Mn(I1) was vio in that the nickel(I1) chelate was found to be less stable than those of cobalt(II) and zinc(I1). In order to explore these effects in greater detail it was thought appropriate to determine the AZ? and AS values for the formation of chelates of z-methyl-8hydroxyquinoline. The AI;, AI3 and AS values for ths chelation reactions of f-methyl-8-hydroxyquinoline were obtained as a comparison. ,Formation constants of chelates of z-phenyl-8-hydroxyquinoline were also determined to extend our study of sterically hindered 8-hydroxyquinolines. EXPERIMENTAL Materials. The method of preparation of 2-methyl-8-hydroxyquinoline was described previouslp. 4-Methyl-8-hydroxyquinoline was prepared from methyl vinyl ketone and o-aminophenol by the DOEBigER-VONMILLERreaction7. After recrystalliRefcvences
p.
2x3.
202
W.
D. JOHNSTON, H. FREISER
VOL. 11 (1954)
zation from alcohol and water it melted at 140-142~ (reported’ 141”). 2-Phenyl-8hydroxyquinoline was prepared by the method of GILMAN from 8-hydroxyquinoline and phenyllithium 7. After decolorization with charcoal and recrystallization from alcohol and water it melted at 58.5-5g” (reported’ 59”). The method of purification of dioxane and the procedures used in standardizing metal perchlorate solutions have been previously describcds. A$$aratus and Procedure. The titration apparatus and procedure have been previously described s. Temperature was controlled at 0.7 & 0.1~ C by immersing the titration apparatus in a water-ice bath. At 25 and 40” C temperature was controlled to f 0.2.~ C by circulating water from a constant temperature bath through the water jacket of the titration vessel. An atmosphere of nitrogen, previously saturated with a 50% water-dioxane mixture at the appropriate temperature, was maintained over the titration mixture. The CALVIN-B JERRUM method of determining stability constants Calculations. was applied as previously described *. As mentioned in paper IV, it has been the general practice in this type of study to take pH meter readings as an accurate indication of hydrogen ion concentration. In many cases this is a very satisfactory approximation but where more accurate formation constants are required, such as in the determination of temperature coefficients of stability constants, it becomes necessary to make small corrections. These are particularly important at the higher temperatures. By comparison of pH meter readings with theoretically cc!sulated hydrogen ion concentrations, correction factors were determined which were independent of PH but which varied with temperature. Thus it was possible to correct PH meter readings to hydrogen ion concentrations by adding respectively 0.01, 0.07, and 0.14 to the PI-I meter reading at 0.7” C, 25” C and 40” C. In order to be strictly comparable, the formation and acid dissociation constants at 25” C for 2-methyl-8-hydroxyquinoline chelates were redetermined at the conditions of the present work. This involved slightly larger concentrations of the reacting species. This factor combined with the correction to pH readings which were not previously applied gives stability constants that are somewhat lower than those previously reported. Results. Table 1 lists the stepwise formation constants for chelates of z-methyland 4-methyl-8-hydroxyquinoline determined at 0.7, 25 and 40” C. The free energies of formation at these temperatures are also listed. The free energy terms refer to complete chelate formation and are calculated from the values given for For comparison, the formation constants for chelates 2 log KW (KW = fK,K,). of 2-phenyl-8-hydroquinoline are included. The acid dissociation constants of the three reagents are also listed. All data were obtained in pH regions which previously had been shown to be free of metal hydrolysiss. Hydrolysis of the metal ion or precipitation of the metal chelate precluded precise calculation of stability constants in a number of instances. In some of these cases calculated or extrapolated values are listed parenthetically. References p. arg.
VOL.
ORGANIC
11 (r954)
ANALYTICAL
REAGENTS
VI
203
The stability data listed in Table I were used in calculating AH of complete chelation by measuring the slopes of the curves shown in Figures Ia, b, and c. The AS values of the formation reactions were. calculated from the AF and AH values. The heats and entropies of chelation are given in Table II.
-_.q-Methyl-Hhydroxycluinollnc -._-. ___. -.. _-
-
-_.
I’Cnlp,
“C
log
_--_
0.7
I’.36
Co(1 I) Zn(Il) E(K)
0.7 0.7 0.7 ‘5
I I .29 II.25
25 25 2s 25 25 “5 4” 40 40 40 -lo
R((IYj Cd(l I) Fu”(‘I’$ Ni(I1) Co II) Zn \II) Mn(I1) .- _-. . 2-Methyl-8hydroxyquinohnc
. --_.
- --_. --Temp. “C
-._-_-
Cu(I1) Ni II) co I II) Zn(I1)
_---
o-7 0.7 0.7 o-7 o-7 25 25 25 25 25
%$I’, ?A .I ::,, Zn(I1) g!lflf~ zg$Yx))
25
:;‘(‘I:;)
40 25
E
40 40
I,:,, .i C& [II)!)
Referewes
p.
4: 213.
_-_-- _ I-’ kcal.
_
c(Y::
z(,::; 7. Zn(II)
--_A
Ii, ’ 4.32 -
(7.60) 13.52 (ro.7.z) (9.45) (9.57) 7.24 8.13 -
1 I .57 IO.55 to.67 8.31 II.1 1
9.44
6.45 I x.15
5.46 12.8% 10.17
LO.22 IO.25 H. IL!
G9) (6.94)
-. _ ___----
12.94 9.67 9.97 10.07 7.75 12.48 9.41 ;z 7144 10.30 9.00 5.24 7.7’
x2.12
9.07 9.37 9.47 7.40
2 I .4o 16.22 (;;.g)
’ .
20.00
20.24 ‘5.50 I 8.76 (‘7.20) I I.90 (26.46) 21.34 (19.28) 19.44 15.00
_-_. _-__-. --
log
A-,
2 log
I&
.- ..- -_-
.._
-- _ -..----.
37.33 29.93 26.63 26.76 20.28 37.58 30.40 27.29 27.62 21.15 25.60 23.47 16.24 37.92 30.58 27.63 27.86 21.58
IO.08
(10.28)
8.63 -
.____.._ log Ii,
(2;.&6)
I I.74
8.71 9.17 9.27) I6.85) II.32 8.35 8.87 6.55 (8.20)
-
4.48 IO.82
8.22 8.74 6.50
24.74 18.44 19.16 19.32 14.56 23.84 17.84 ‘8.44 I 8.72 14.00 18.50 (16.60) 9.72 23.00 17.36 18.14 (18.28) 13.90
-_-_---A
P
kcal.
-_ --.--30.93 23.06 23.96 24.16 I 8.20 32.53 24.34 25.16 25.54 1g.10 25.24 22.65 I 3.26 32.96 24.88 26.00 26.20 ‘9.92
.
W.
20~
D.
JOHNSTON,
Ii.
FI
TABLE I (Continued) .-----_ ___z-Phenyl-8Temp. log K, log K, hyclroxyquinoline OC -.__-----__ ..-_-__ -. IO.89 Cu(II) II.40 25 Ni II) 6.93 7.57 25 cu II) 25 (@ zn i II) 8.52 25 Mn(l1) ((3.22) 25
VOL.
11
(1954)
_--
-__
---
._-.. ._---
SUMMARY
_--. ..4-Methyl-8hydroxyquinollne 4-Methyl-S hydroxyquinoline 4-Methyl-ghydroxy uinoline 2-M& it yl-8hydroxy uinoline z-Met 4, yl-8hydroxy uinolinc 2-Met is.yl-8hydroxyquinoline 2-Phcnyl-8hydroxyquinoline ---- --_-_
2
-A
F
kcal.
22.26 ‘4.50 f 7.54
:;*“,z . 23.93
-
-
_-.._
_. -_..-
._
2 log Kav
017
ACID
_-
DISSOCIATION
Temp. @
@(OH
CONSTANTS
kcal.
AI;NH
A
FOH
o .7
5.18
x2.22
6.48
IS.27
25
4.67
11.62
6.37
IS.85
40
4.45
11.31
6.38
16.20
0.7
5.00
12.18
6.25
X5.23
25
4.58
II.71
G.24
15.98
40
4.3 *
11 a44
G.18
16.39
2.83
16.19 .---.
25
_
2.07 __- ..__~
I I .87 - --..__.---._
_- _..
Log
HA V.
24
I
23
/J
22
: .
i
ra:Temperature
References
fi.
_&/
20
i
Pig.
2x3.
kcal.
dependence
./
of chelnte stability.
4-methyl-Shydroxyquinoline.
11 (1954)
VOL.
ORGANIC
ANALYTICAL
REAGENTS
VI
25 24 2LO9
2J
‘CAV
Fig. 1b. Tcmpcrature
The
constants
dcpcndcncc
reported
in
4-Methyl-8-Hydroxyqu~nolme
of chelatc stability,
Table P-Methyl-8
II
are
r-nlcthyl-8-hyclroxycluinolinc.
not
-Uydroxyqurnolrno
122r 12.7 .
12.2
II.9
.
12.0
128
.
72.7 -PKOH
11.9
71.7. 720 71.6 11.5 :.
II.8
,,!
71.7 716
7l-3.
71.4
115 11.4 5.1 5.0 4.9 4.8 4.7 4.6 4.5 4.4
4’4
32
3.4
4.3
3.6 VT
Fig.
IC. Temperature
References
p.
2x3.
x 10
dependence constants.
of reagent
ionlzntion
thermodynamic constank in that no activity coefficients have been employed. The low ionic strengths encountcxed (0.005) minimize the necessity of such a correction when using the stability data at one temperature for purposes of comparison. The neglect of activity coefficients when heats and entropies of reactions are calculated results in errors only insofar as the neglected activity coefficients vary with temperature. It is estimated that thisvariation results in errors which are negligible compared with normal cxperimental errors.
W.
206
D.
JOHNSTON,
H.
TABLE
II
IfEATS ANI> liXTROPIES OP CliELAT15 . __ ____ ____ __-_ -._ _ _.-._._. .__---
-AI/ kcal. 4-Methyl-Shydroxyqtimolinc ._ .- __-_. _ -cu (II) 34.1 26.0 Ni II) 20.3 I 2 (‘I:‘, 19.6 I I.6 MnlIIj_: --.. _A~-Methyl-% hydroxyquinoline -‘-.- -_- --_-._- __ I7..{ .-_ Cu(IT) Ni(TT) IO.7 cO(I[j ‘0.3 h(Irj IO.3 6.0 Mn(II) ___-_~ - __ 4-Methyl-8hyclroxyquinoline _.. __.._ ..‘-7. __----. lm‘311no I-i* 7.’ 0. I ohenollc I-I+ _ z-Methyl-8hyclroxyquinolinc --_. _ _-~~~~;ai-~~__. _ 6.9 phencAic EI* 7-s _._._ --_ . .._-_ 1_-_
__
..-
-_
VOL. 11 (1954)
FREISEH
YORMATIOS .___-...
AS
c. II. _ .
12.
15. -3.
‘,
26. 3’.
50.
.&G. so51. ‘) I . .. ..-__.. -_
_.
-2.
23. _
__.
-2:
28.
_-__
-
(“1 exprcssetl as formation constants fr’o,.r
-
_._HR.__. (H+) (I<-)
DISCUSSION
OF
RESULTS
As can be seen from Table I the acid dissociation constants of the two methyl-8hydroxyquinolines are. very similar. This would indicate that the total of the inductive and hyperconjugative effects of the methyl group on reagent basicity is nearly independent of its position. Thus, unless there is some type of steric hindrance involved, these two reagents would be expected to react with a given metal to form chelates of equal stability. Table I shows that the overall formation constants for all chelates of z-methyl-& hydroxyquinoline are from 1.5 to 5 log units less than those for the corresponding chelates of 4-methyl-8-hydroxyquinoline. This points immediately to the operation of some type of steric hindrance in the chelates of 2-mchyl-S-hydroxyquinoline. This stability decrease is more graphically illustrated in Fig. 2 where the overall formation constants at 25” C of the chelates of the two reagents are plotted against the second ionization potential of the gaseous metal atom. Fig. 2 also points out the previously reported fact that the stability of the nickel(H) chelate of z-methyl-8-hydroxyquinoline is considerably less than would Refcrenctx
p.
2x3.
VOL.
11
(1954)
ORGANIC
ANALYTICAL
REAGENTS
VI
207
be expected. Such an effect is also found in the nickel(I1) chelate of zz-phenyl-8hydroxyquinohne. In fact in both cases the usual metal stability sequence is violated such that the nickel(I1) chelates are less stable than the corresponding cobaIt(I1) and zinc(H) chelates. Work in this laboratory has definitely established that this relative decrease in the stability of a nickel(I1) chelate is always found in cases where steric hindrance is possible. In a previous report on the stabilities of chelates of dimethylglyoxime the usual metal stability sequence is adhered to while in the case of its 0-monomethyl ether the hindered metal order is foundr.
151 1.3’
’ 14
15
-
A
16
17
’
’
18
’
79
’
20
. 22 ’ 2.3 * 24 * 25 ’ 26 ’ 27 ’ 28 21 2L'?7KAV
Fig. 2.
This unusual order has also been found in chelates of z-(o-liydroxyphenyl)-benzothiazoline and z-(o-hydroxyphenyl)-benzoxazole O. In every case when there is a substituent group on the carbon in alpha position to the heterocyclic nitrogen the hindered stability sequence has been found. This effect is still not clearly understood but it seems reasonable to assume that the tendency of nickel(I1) to form planar complexes is being hindered, giving rise to either a strained or distorted planar configuration or possibly even a tetrahedron. A molecular model of the nickel(I1) chelate of 2-methyl-8-hydroxyquinoline shows that a symmetrical planar configuration would be nearly impossible to attain since there would be considerable interference between the methyl group of one reagent and the oxygen of the other. This point of view is difficult to reconcile with the findings of MELLOR AND CRAIG~O who reported the nickel(I1) &elate of 8-hydroxyquinoline to be paramagnetic and hence presumably tetrahedral. On the other hand, the nickel(I1) chelate of o-aminophenol was found to be diamagnetic and therefore of square coplanar configuration. From these magnetic data one would predict that the bonds in the o-aminophenol chelate would be stronger, giving rise to greater References
$.
213.
W.
208
VOL. 11
D. JOHNSTON, H. FREISER
chelate stability, than those of the 8-hydroxyquinoline reverse has been found to be true ‘JJl. Further, the nickel(II) N was exceptional
in that it was the only one of the class
(1954)
chelate. Actually the 8-hydroxyquinolinate 0
‘Ni’
chelates
(where 0
0’ ‘N originally formed part of a phenolic hydroxide group) which displayed paramagnetism in the MELLOR AND CRAIG investigation. If we assume that the magnetic properties in nickel(I1) chelates of alkylated 8-hydroxyquinolines are similar, another anomaly becomes apparent. The fact that the nickel chelates should be tetrahedral, according to magnetic data, is not in agreement with the argument relating to the unusual position of nickel(II) in the hindered metal stability sequence. These inconsistencies together with the fact that the nickel(I1) chelate of 8-hydroxyquinoline is an exception to the MRLLOR AND CRAIG classification indicate the need for either a re-examination of the magnetic susceptibility of this compound or the re-interprctationzof the magnetic data. 4-Methy/ - 8 - Hydroxyquinoline 34
-AH
2 -Methyl-8
- Hydroxyquinoh’ne
t
32 * 30. ae2624 222018. 161472L*
18 AH
22
*
24
*
26
“0’
26
30
32 34
-A Fwc
*
36
1.
38
CU
1614-
61mA20
’
22
’
24
’
’ 30’ 32* 34’
26 28
-
Fig.
3,
-A&PC
Fig. 3 shows a plot of AF ZISAH for the chelates of the two reagents. It can be seen that these curves are roughly linear. This is to say that the entropy changes of the reactions have caused no unusual effects within the chelates of a given reagent, It is reasonable to assume that this will generally be true. That is, the entropy of a series of chelation reactions with a given reagent will either be roughly constant or will vary regularly. This accounts for the widespread success of the log K, second ionization potential graphs. Each time this relationship References
p.
2r3.
VOL.
ORGANIC
11 (1954)
ANALYTICAL
REAGENTS
cog
VI
4 is used it is implicitly
assumed that only bond strength or AH is involved. Since the second ionization potential may be taken as a rough estimation of the average electron attracting power of a divalent metal ion it will also be a measure of the attracting power of that divalent metal ion for a source of electrons such as are found in the chelate ligand groups. Hence it is more nearly related to bond energy and AH of chelation than to AF as measured by the log of the chelate formation constant. It should also be noted that since in many cases changes in AH are accompanied by opposing changes in AS, AH should be a more sensitive index of bond strength than AF. Fig. 4 shows a plot of the second ionization potential of the gaseous metal atom against AH of chelation. As can be seen the relationship is satisfactory.
0 ._ 8 .- 77 B s 16
.z18t
751 6’
’
8
’ ’ ’ ’ 70 12 14 76 78 20 22 24 26 28 30 32 34 36 -bka1. ’
’
’
*
’
’
’
’
-
’
Fig. 4. Very few determinations of AH have been published for neutral chclate species. However, a number of such investigations have been made on the metal ammine type of complexes. Of particular interest are the very comprehensive studies by JONASSEN et nZ.l*on complexes of diethylenetriamine and triethylenetetramine. The heat of reaction sequence for the complexes of both of these reagents follow the normal stability order rather closely. MCINTYRE'~ gives a rather complete compilation of the thermodynamic functions for metal amminc formation. In practically all cases listed the heats of reaction for complexes of a given reagent follow the normal metal stability sequence. A comparison of the AH values listed in Table II points out the importance
of steric hindrance in all chelates of 2-methyl-S-hydroxyquinoline. In all cases the heats of chelation for chelates of z-methyl-8-hydroxyquinoline are only about half that found associated with the respective chelates of the 4-methyl compound. Refere?rcesp. 2r3.
W.
2 IO
I>.
JOHNSTON,
Ii.
FREISER
VOL.
11
(xg!jd$)
This points to a very large decrease in bond strength due to the hindrance of the a-methyl group. It is reasonable to assume that the presence of the 2-methyl group prevents the close approach of the two reagent molecules around the metal atom regardless of the preferred configuration and thus causes the chelate bonds considerable strain. Therefore, one might reasonably expect the magnitude of the steric effect to increase with decreasing metal size. For this purpose, the thermodynamic functions of the exchange reaction, MQ, + 2L- t ML, + eQwhere Q- and L- arc the anions of z-methyl- and 4-methyl-8-hydroxyquinoline, respectively, were plotted against the divalent crystal radii of the metals involved. Although there arc a number of conflicting sets of values of crystal radii, those of PAULING~~ were used mainly because these are the values used also by POWELL AND LATINIER~~ in similar thermodynamic calculations. As can be seen from Fig. 5 the TAS term varies regularly with the crystal radius while with AF and
;I ;----_:. , *-?zIc
.
0.70
Fig.
5.
Comparison
of
0.74 0.76 . 0.72 Crystal radius in A0 (after Pouting)
thermodynamic functions of alkylzlted 8-hydroxyquinolinc chelates.
AH there are sharp discontinuties in the case of nickel(H). In agreement with our expectations, the exchange reactions become more exothermic as the crystal radii decrease. It appears that in the case of nickel(I1) a minimum radius has been reached and steric hindrance is excessive. This helps to explain the decreased stability that has been observed in the case of hindered nickel(I1) chelates. Rcfevences
+.
213.
VOL.
11
ORGANIC
(1954)
ANALYTICAL
REAGENTS
VI
2x1
The relatively large entropy increase associated with the formation of chelates of 2-methyl&hydroxyquinoline was rather surprising but not unprecedented. FRANK AIUD EVANS’s have interpreted the following data in terms of steric hindrance and the structure-making influence on water of the reagents and complexes. AH = -rg kcal., AS = -22 Ni+2(aq.) + 6NH,(aq.) ---L [Ni(Nl-I,),Jf2(aq.); e. u. Ni+z(aq.) + GCH,NH,(aq.) =+ [Ni(CH,NH&jC2 (aq.); AH = tg.7 kcal., AS = +73 e. u. Although the complexes formed are of about equal stability spatial requirements prevent the close approach of the six methylamines about the nickel(H) ion according to arguments presented by CALVIN and discussed by FRANK AND EVANS. This easily accounts for the more positive heat of reaction in the hindered case. This is in qualitative agreement with the heat effects presented in this paper. The relatively large positive entropy term in the hindered reaction above is also in agreement with the results prcscntcd hcrc. The interpretations of these entropy effects are rather difficult. FRANK AND EVANS
attributed
only
a small
part
of
the
greater
entropy
of
formation
of
the
greater size relative to that of the ~i(NH,),Jf2 complex. PJW-WH2M They speculated that the larger part of the entropy contribution might be due to the imposition of iccbcrg-forming tendencies on the methylaminc molecule by the methyl group which would lower its partial molal entropy compared with that of ammonia. If the water structure breaking tcndcncies of both complexes were similar this would account for the greater entropy of reaction in the second case. Although not strictly comparable, the thermodynamic data of SPIKE AND YARRF which describe the formation of [Cd(NH,),lC2 and [Cd(CH,NH2),]+2 tend to cast doubt on this explanation. No steric hindrance was found in the cadmium(II) complexes as evidenced by similar values of AH. Of greater importance is the fact that the AS values of SPIKE AND PARRY do not reflect a trend
+2 to
similar
to
The
that
large
its
reported
difference
by
methyl-&hydroxyquinoline methyl
groups
able
to
partially
such
shielding
would
reduce
atoms
and
remnant An
in
the
the
a first
interesting
for
chelates
is
probably
of
polar
a chelate
oxygen, in the
forces
the
formation best
of
greater
of
explained
z-methylby
and
metal
and
noting
S-hydroxyquinolinc
nitrogen,
corresponding
between
much
coordination
alternative
EVANS.
values
be possible solvent
AND
AS
z-positions
the attractive
allow of
the shield
would
FRANK
in the
atoms
4-methyl-chelate.
4-
that
might
be
while
no
Shielding
the solvent
molecules
and
these
polar
frceclom
by
virtually
eliminating
any
sought
in
sphere. explanation
can
be
the
inverse
relation
-AH and AS. In the case of the exchange reactions mentioned above and also in the case of formation of 4-methyl-S-hydroxyquinolinc chelates there seems to be a linear inverse relation between -AH and AS. It is probable that as the bond strength decreases, the freedom of motion involving the atoms so of
Referewes
p.
2x3.
W. D. JOHNSTON,
2x2
H.
FREISER
VOL.
‘1
(r9s-I )
bonded, increases. This increased freedom of motion may well result in a significant entropy increase. This relationship involving intramolecular forces may be considered analogous to that of BARCLAY AND BUTLER~~ (linear relation of AH and AS of vaporization of certain non-ionized solutes) in which intermolecular, non-valence, forces are involved. It would bc interesting to study the entropies and heats of chelation for chelates II
the
financial
support
of the U.
S. Atomic
SUMMARY The CALVIN-RJERRUM titration technique for the determination of chelate formation constants has been applied to the Cu(II), Ni II). Co(II), Zn(I1) and Mn(I1) were made at chelates of z- and 4-methyl-8-hydroxyquinoline. L easurements several temperatures in order to evaluate Al-f and AS values of chelation. The results obtamed were interpreted in terms of steric hindrance of the z-methyl group. In all cases the heats of formation of the chelates of 2-methyl-8-hydrox quinoline were remarkably more ositive than those for the corresponding chelates o Y +methyl8-hydroxyquinoline, 3%. IS large difference in the strengths of the metal-chelate bonds is apparently due to the hindrance of the methyl groups which prevent the close grouping of the two reagent molecules around the metal in chelates of z-methyl8-hydrpxyguinoline, The lower bond strength in chelates of 2-methyl-8-hydroxycompensated by a relatively larger entropy of formation. umohne IS partially % his is attributed to decreased solvent chelate interaction caused by the shielding of the polar 0, N, and metal atoms by the z-methyl groups. The determination of chelate formation constants of z-phenyl-8-hydroxyqumolme has been carried out to further cxtcnd our study of stcric effects in metal chelates. Rl%UMfi Les auteurs ont appliqud la technique de titrage de CALVIN-BJERRUM B la determination dcs constantes de formation des complexes des 2- et +m&hyl-8hydroxyquinol&nes avec Cu, Ni, Co, Zn et Mn. Les mesures ont Bt6 faites tLdiverses temperatures afin d’8valucr les valeurs de chelation AH et AS. Les rQultats obtenus ont &l! interprCtds en fonction de l’empechement stdrique provenant du Les chaleurs de formation des complexes de la 2-methylou ement 2-m&hyl. Er -hy x roxyquinol&ne sont toujours nettemcnt plus positives que celles des comReferences
p.
ar3.
VOL.
11
(19%)
ORGANIC ANALYTICAL
REAGENTS
VI
2 I.3
plexes de la 4-methyl-8-hydroxyquinoleine. Cette grande difference dans les forces de liaison entre le metal et la molecule organique semble Btre due B l’empechement sterique des groupes m&h 1 rendant t&s difficile la formation de chelate. La faible force de liaison dans les c Helates de la z-methyl-8-hydroxyquinol&ine est partiellement corn ensee par une relativement grande entropie. Cette derniere est attribuee a un affai +Zlissement de l’interaction entre le dissolvant et le complexe, les groupements methyliques empechant le reactif de reagir avcc l’oxygene, l’azote et les atomes metalliques polariscs. Les autcurs ont determine egalement les constantes de formation des complexes dc la z-phcnyl-S-hydroxy uinoleine pour faire mieux comprendre l’importance des facteurs stdriques dans 3 e domaine des complexes. ZUSAMMENFASSUNG Die CALVIN-BJnRRUMSche Titrationsmethode zur Bestimmung der Komplexbildungs-Konstanten ist auf die Cu II)-, Ni(II)-, cO(II)-. Zn(II)- und Mn(II)-Komplexe \ins angewandt worden. Es wurden Messungen bei des 2- und 4-Methyl-8-oxychino verschiedenen Temperaturen ausgefiihrt urn die Wiirmettinung und die Entropic Die Ergcbnisse sind im Sinne sterischer der Komplexbildung zu bestimmen. Hinderung der 2-Methylgruppe gedeutct worden. In ailen Fallen war die Warmetonung dcr Komplexbildung der Komplexe dcs 2-M&h 1-8-oxychinolins bcmerkenswert positiver als die derselben Komplexe des 4- d ethyl-8-oxychinolins. Dieser der Metall-Komplexbindungen ist angrosse Untcrschied der Bindungsenergie scheinend auf die iMethyl ruppen, die einc dichte Gruppierung der zwei Rcagenz2-Methyl-8-oxychinolinkomplcx verhindern, molckiile um das Meta P1 in dem Die kleincre Bindungsenergie der Komplexe des 2-Methyl-gzuriickzuftihren. kompensiert. Dies wird oxychinolins ist teilweise durch eine grossere Entropie gedeutet als eineverminderung der \Vcchselwirkung zwischen Komplex und Losung?: mittel durch Abschirmung der polarcn 0, N, und Metallatomen durch die Methylgruppen. Die Bildungskonstantcn der Komplcxe des 2-Phen I-8-oxychinolins wurdcn bestimmt urn unsere Forschung der sterischen Effekte in X9etallkomplexcn weiter ausendchnen. REFERENCES 1
7
Ii. G. CHAKLES AND H. FREISER, Anal. C/rim. Ada, 11 (1954) IOI. Abstracted from the thesis submitted by \V. D. JOI-INSTON in partial fulfillment of Pittsburgh, of the requircmcnts for the Ph. D. degree at the University address : \Vcstinghousc Research Laboratories, East August 1953. Present Pittsburgh, Pa., U.S.A. L. L. MRRRKTT AND J. Ii. ~VALKDR, Ind. E?rg. Clrcrn., Anal. Ed., 16 (1944) 387. E. J. BUTLER AND M. F. RING, J. C/rein. Sot., (x949) 1489. H. IRVING, New York, rg5 1. A. ALBERT, SelEcttvc To.uicLty, John \Vlley & Suns Inc., \\‘. D. JO~INSTON AND l-l. FREISER, J. Al/r. C/rem. Sot., 74 (x952) 5239. 1. P. PJIILLIXS, 11. L. ELI~J~NGEIZ AND I,. L. MoRRIr*r Jr., J. Am. Chem. Sot..
8
1949) 3986. REISDR, 12. G. ‘H’.B
2
3 4 5 0
0 10 11 12 la
CHARLES AND \V. 13. JOIKNSTON,J. Am. C/rem. Sot., 74 (rg52) 1383. Ii. G. CRARLES AND I-1. FR&ISI:R, Anal. Chtm. Ada, IK (1954) I. D. I?. MELLOR AND D. P. CRAIG, J. ~'YOC. Roy. SOL. N. S. Wales, 74 (1940) 475. Ii. G. CHARLES AND H. FRRIsRR, J. Ant. C/rem. Sot., 74 ISIBOIIM,J. PlbyS. I-I-Ieza. JONASSEN, G. G. HURST, l<. 13. LRBLANC AND A. VV. 56 (1953) 16. G. H:' MCINTYRB, Ph. D. Thesis, The Pennsylvania State College (x953 . lVatzwe of the Chemical Bad, Cornell University Press, It h aca,
17 C. G. 18 1. M.
AND \V. M. LATIMER, J. Chem. Whys., Ig (1951) 1139. M. \V. EVANS, J. C/bent. Plbys., 13 (1945) 507. SPIKE AND R. \V. PARRY, J. Am. Clbewz. Sot., 75 (1953) 2726 BARCLAY AND J . A. V. BUTLER, l’umas. Faraday So&, 34 (1938) 1451. Received January ~rth, rg54
ANALYTICA CHIMICA ACTA
11 (rg51)
STRUCTURE VI.
HEATS
AND
20 I
BEHAVIOR OF ORGANIC REAGENTS
ANALYTICAL
AND ENTROPIES OF FORMATION OF DIVALENT METAL CHELATES OF z- AND a-METHYL-S-HYDROXY QUINOLINE’
SEVERAL
by W. Cort/rlbutlon
D.
JOHNSTONS
of llrc Uepnrtnaent
of Chenaisiry,
AND
H.
FREISER
Untvcvsity
01
Ptltsburgh,
Pa.
(U.S.
A .)
It has been well established that any substitution in the z-position of 8-hydroxyquinoline prevents the precipitation of an aluminum complex. This effect was first noted by MBRRITT AND WALKBH~ in the case of a-methyl-8-hydroxyquinoline. IRVINC~ and co-workers established the general na turc of this effect by a-study of the analytical behavior of z,4-dimcthyl- and z-phenyl-8-hydroxyyuinoline as and g-hydroxy-I,z,3,+tetrahydroacridine. They well as x-hydroxyacridine reasoned that this effect was due to the steric hindrance provided by the group in the z-position. ALBERTA has noted that any substituent in the z-position of 8-hydroxyquinoline lowered its antibacterial action. In a previous paper the authors O have reported formation constants for a number of metal chelates of 8-hydroxyquinoline and z-methyl-8-hydroxyquinoline. It was found that the chelates of z-methyl-8-hydroxyquinoline were slightly less However, the nickel(I1) chelate of stable than those of 8-hydroxyquinoline. z-methyl-8-hydroxyquinoline was considerably less stable than the corresponding This unusual effect was so pronounced th t the chelate of 8-hydroxyquinolinc. 7 ated usual metal sequence Cu(I1) > Ni(II) > Co(II), %n(II) > Mn(I1) was vio in that the nickel(I1) chelate was found to be less stable than those of cobalt(II) and zinc(I1). In order to explore these effects in greater detail it was thought appropriate to determine the AZ? and AS values for the formation of chelates of z-methyl-8hydroxyquinoline. The AI;, AI3 and AS values for ths chelation reactions of f-methyl-8-hydroxyquinoline were obtained as a comparison. ,Formation constants of chelates of z-phenyl-8-hydroxyquinoline were also determined to extend our study of sterically hindered 8-hydroxyquinolines. EXPERIMENTAL Materials. The method of preparation of 2-methyl-8-hydroxyquinoline was described previouslp. 4-Methyl-8-hydroxyquinoline was prepared from methyl vinyl ketone and o-aminophenol by the DOEBigER-VONMILLERreaction7. After recrystalliRefcvences
p.
2x3.
202
W.
D. JOHNSTON, H. FREISER
VOL. 11 (1954)
zation from alcohol and water it melted at 140-142~ (reported’ 141”). 2-Phenyl-8hydroxyquinoline was prepared by the method of GILMAN from 8-hydroxyquinoline and phenyllithium 7. After decolorization with charcoal and recrystallization from alcohol and water it melted at 58.5-5g” (reported’ 59”). The method of purification of dioxane and the procedures used in standardizing metal perchlorate solutions have been previously describcds. A$$aratus and Procedure. The titration apparatus and procedure have been previously described s. Temperature was controlled at 0.7 & 0.1~ C by immersing the titration apparatus in a water-ice bath. At 25 and 40” C temperature was controlled to f 0.2.~ C by circulating water from a constant temperature bath through the water jacket of the titration vessel. An atmosphere of nitrogen, previously saturated with a 50% water-dioxane mixture at the appropriate temperature, was maintained over the titration mixture. The CALVIN-B JERRUM method of determining stability constants Calculations. was applied as previously described *. As mentioned in paper IV, it has been the general practice in this type of study to take pH meter readings as an accurate indication of hydrogen ion concentration. In many cases this is a very satisfactory approximation but where more accurate formation constants are required, such as in the determination of temperature coefficients of stability constants, it becomes necessary to make small corrections. These are particularly important at the higher temperatures. By comparison of pH meter readings with theoretically cc!sulated hydrogen ion concentrations, correction factors were determined which were independent of PH but which varied with temperature. Thus it was possible to correct PH meter readings to hydrogen ion concentrations by adding respectively 0.01, 0.07, and 0.14 to the PI-I meter reading at 0.7” C, 25” C and 40” C. In order to be strictly comparable, the formation and acid dissociation constants at 25” C for 2-methyl-8-hydroxyquinoline chelates were redetermined at the conditions of the present work. This involved slightly larger concentrations of the reacting species. This factor combined with the correction to pH readings which were not previously applied gives stability constants that are somewhat lower than those previously reported. Results. Table 1 lists the stepwise formation constants for chelates of z-methyland 4-methyl-8-hydroxyquinoline determined at 0.7, 25 and 40” C. The free energies of formation at these temperatures are also listed. The free energy terms refer to complete chelate formation and are calculated from the values given for For comparison, the formation constants for chelates 2 log KW (KW = fK,K,). of 2-phenyl-8-hydroquinoline are included. The acid dissociation constants of the three reagents are also listed. All data were obtained in pH regions which previously had been shown to be free of metal hydrolysiss. Hydrolysis of the metal ion or precipitation of the metal chelate precluded precise calculation of stability constants in a number of instances. In some of these cases calculated or extrapolated values are listed parenthetically. References p. arg.
VOL.
ORGANIC
11 (r954)
ANALYTICAL
REAGENTS
VI
203
The stability data listed in Table I were used in calculating AH of complete chelation by measuring the slopes of the curves shown in Figures Ia, b, and c. The AS values of the formation reactions were. calculated from the AF and AH values. The heats and entropies of chelation are given in Table II.
-_.q-Methyl-Hhydroxycluinollnc -._-. ___. -.. _-
-
-_.
I’Cnlp,
“C
log
_--_
0.7
I’.36
Co(1 I) Zn(Il) E(K)
0.7 0.7 0.7 ‘5
I I .29 II.25
25 25 2s 25 25 “5 4” 40 40 40 -lo
R((IYj Cd(l I) Fu”(‘I’$ Ni(I1) Co II) Zn \II) Mn(I1) .- _-. . 2-Methyl-8hydroxyquinohnc
. --_.
- --_. --Temp. “C
-._-_-
Cu(I1) Ni II) co I II) Zn(I1)
_---
o-7 0.7 0.7 o-7 o-7 25 25 25 25 25
%$I’, ?A .I ::,, Zn(I1) g!lflf~ zg$Yx))
25
:;‘(‘I:;)
40 25
E
40 40
I,:,, .i C& [II)!)
Referewes
p.
4: 213.
_-_-- _ I-’ kcal.
_
c(Y::
z(,::; 7. Zn(II)
--_A
Ii, ’ 4.32 -
(7.60) 13.52 (ro.7.z) (9.45) (9.57) 7.24 8.13 -
1 I .57 IO.55 to.67 8.31 II.1 1
9.44
6.45 I x.15
5.46 12.8% 10.17
LO.22 IO.25 H. IL!
G9) (6.94)
-. _ ___----
12.94 9.67 9.97 10.07 7.75 12.48 9.41 ;z 7144 10.30 9.00 5.24 7.7’
x2.12
9.07 9.37 9.47 7.40
2 I .4o 16.22 (;;.g)
’ .
20.00
20.24 ‘5.50 I 8.76 (‘7.20) I I.90 (26.46) 21.34 (19.28) 19.44 15.00
_-_. _-__-. --
log
A-,
2 log
I&
.- ..- -_-
.._
-- _ -..----.
37.33 29.93 26.63 26.76 20.28 37.58 30.40 27.29 27.62 21.15 25.60 23.47 16.24 37.92 30.58 27.63 27.86 21.58
IO.08
(10.28)
8.63 -
.____.._ log Ii,
(2;.&6)
I I.74
8.71 9.17 9.27) I6.85) II.32 8.35 8.87 6.55 (8.20)
-
4.48 IO.82
8.22 8.74 6.50
24.74 18.44 19.16 19.32 14.56 23.84 17.84 ‘8.44 I 8.72 14.00 18.50 (16.60) 9.72 23.00 17.36 18.14 (18.28) 13.90
-_-_---A
P
kcal.
-_ --.--30.93 23.06 23.96 24.16 I 8.20 32.53 24.34 25.16 25.54 1g.10 25.24 22.65 I 3.26 32.96 24.88 26.00 26.20 ‘9.92
.
W.
20~
D.
JOHNSTON,
Ii.
FI
TABLE I (Continued) .-----_ ___z-Phenyl-8Temp. log K, log K, hyclroxyquinoline OC -.__-----__ ..-_-__ -. IO.89 Cu(II) II.40 25 Ni II) 6.93 7.57 25 cu II) 25 (@ zn i II) 8.52 25 Mn(l1) ((3.22) 25
VOL.
11
(1954)
_--
-__
---
._-.. ._---
SUMMARY
_--. ..4-Methyl-8hydroxyquinollne 4-Methyl-S hydroxyquinoline 4-Methyl-ghydroxy uinoline 2-M& it yl-8hydroxy uinoline z-Met 4, yl-8hydroxy uinolinc 2-Met is.yl-8hydroxyquinoline 2-Phcnyl-8hydroxyquinoline ---- --_-_
2
-A
F
kcal.
22.26 ‘4.50 f 7.54
:;*“,z . 23.93
-
-
_-.._
_. -_..-
._
2 log Kav
017
ACID
_-
DISSOCIATION
Temp. @
@(OH
CONSTANTS
kcal.
AI;NH
A
FOH
o .7
5.18
x2.22
6.48
IS.27
25
4.67
11.62
6.37
IS.85
40
4.45
11.31
6.38
16.20
0.7
5.00
12.18
6.25
X5.23
25
4.58
II.71
G.24
15.98
40
4.3 *
11 a44
G.18
16.39
2.83
16.19 .---.
25
_
2.07 __- ..__~
I I .87 - --..__.---._
_- _..
Log
HA V.
24
I
23
/J
22
: .
i
ra:Temperature
References
fi.
_&/
20
i
Pig.
2x3.
kcal.
dependence
./
of chelnte stability.
4-methyl-Shydroxyquinoline.
11 (1954)
VOL.
ORGANIC
ANALYTICAL
REAGENTS
VI
25 24 2LO9
2J
‘CAV
Fig. 1b. Tcmpcrature
The
constants
dcpcndcncc
reported
in
4-Methyl-8-Hydroxyqu~nolme
of chelatc stability,
Table P-Methyl-8
II
are
r-nlcthyl-8-hyclroxycluinolinc.
not
-Uydroxyqurnolrno
122r 12.7 .
12.2
II.9
.
12.0
128
.
72.7 -PKOH
11.9
71.7. 720 71.6 11.5 :.
II.8
,,!
71.7 716
7l-3.
71.4
115 11.4 5.1 5.0 4.9 4.8 4.7 4.6 4.5 4.4
4’4
32
3.4
4.3
3.6 VT
Fig.
IC. Temperature
References
p.
2x3.
x 10
dependence constants.
of reagent
ionlzntion
thermodynamic constank in that no activity coefficients have been employed. The low ionic strengths encountcxed (0.005) minimize the necessity of such a correction when using the stability data at one temperature for purposes of comparison. The neglect of activity coefficients when heats and entropies of reactions are calculated results in errors only insofar as the neglected activity coefficients vary with temperature. It is estimated that thisvariation results in errors which are negligible compared with normal cxperimental errors.
W.
206
D.
JOHNSTON,
H.
TABLE
II
IfEATS ANI> liXTROPIES OP CliELAT15 . __ ____ ____ __-_ -._ _ _.-._._. .__---
-AI/ kcal. 4-Methyl-Shydroxyqtimolinc ._ .- __-_. _ -cu (II) 34.1 26.0 Ni II) 20.3 I 2 (‘I:‘, 19.6 I I.6 MnlIIj_: --.. _A~-Methyl-% hydroxyquinoline -‘-.- -_- --_-._- __ I7..{ .-_ Cu(IT) Ni(TT) IO.7 cO(I[j ‘0.3 h(Irj IO.3 6.0 Mn(II) ___-_~ - __ 4-Methyl-8hyclroxyquinoline _.. __.._ ..‘-7. __----. lm‘311no I-i* 7.’ 0. I ohenollc I-I+ _ z-Methyl-8hyclroxyquinolinc --_. _ _-~~~~;ai-~~__. _ 6.9 phencAic EI* 7-s _._._ --_ . .._-_ 1_-_
__
..-
-_
VOL. 11 (1954)
FREISEH
YORMATIOS .___-...
AS
c. II. _ .
12.
15. -3.
‘,
26. 3’.
50.
.&G. so51. ‘) I . .. ..-__.. -_
_.
-2.
23. _
__.
-2:
28.
_-__
-
(“1 exprcssetl as formation constants fr’o,.r
-
_._HR.__. (H+) (I<-)
DISCUSSION
OF
RESULTS
As can be seen from Table I the acid dissociation constants of the two methyl-8hydroxyquinolines are. very similar. This would indicate that the total of the inductive and hyperconjugative effects of the methyl group on reagent basicity is nearly independent of its position. Thus, unless there is some type of steric hindrance involved, these two reagents would be expected to react with a given metal to form chelates of equal stability. Table I shows that the overall formation constants for all chelates of z-methyl-& hydroxyquinoline are from 1.5 to 5 log units less than those for the corresponding chelates of 4-methyl-8-hydroxyquinoline. This points immediately to the operation of some type of steric hindrance in the chelates of 2-mchyl-S-hydroxyquinoline. This stability decrease is more graphically illustrated in Fig. 2 where the overall formation constants at 25” C of the chelates of the two reagents are plotted against the second ionization potential of the gaseous metal atom. Fig. 2 also points out the previously reported fact that the stability of the nickel(H) chelate of z-methyl-8-hydroxyquinoline is considerably less than would Refcrenctx
p.
2x3.
VOL.
11
(1954)
ORGANIC
ANALYTICAL
REAGENTS
VI
207
be expected. Such an effect is also found in the nickel(I1) chelate of zz-phenyl-8hydroxyquinohne. In fact in both cases the usual metal stability sequence is violated such that the nickel(I1) chelates are less stable than the corresponding cobaIt(I1) and zinc(H) chelates. Work in this laboratory has definitely established that this relative decrease in the stability of a nickel(I1) chelate is always found in cases where steric hindrance is possible. In a previous report on the stabilities of chelates of dimethylglyoxime the usual metal stability sequence is adhered to while in the case of its 0-monomethyl ether the hindered metal order is foundr.
151 1.3’
’ 14
15
-
A
16
17
’
’
18
’
79
’
20
. 22 ’ 2.3 * 24 * 25 ’ 26 ’ 27 ’ 28 21 2L'?7KAV
Fig. 2.
This unusual order has also been found in chelates of z-(o-liydroxyphenyl)-benzothiazoline and z-(o-hydroxyphenyl)-benzoxazole O. In every case when there is a substituent group on the carbon in alpha position to the heterocyclic nitrogen the hindered stability sequence has been found. This effect is still not clearly understood but it seems reasonable to assume that the tendency of nickel(I1) to form planar complexes is being hindered, giving rise to either a strained or distorted planar configuration or possibly even a tetrahedron. A molecular model of the nickel(I1) chelate of 2-methyl-8-hydroxyquinoline shows that a symmetrical planar configuration would be nearly impossible to attain since there would be considerable interference between the methyl group of one reagent and the oxygen of the other. This point of view is difficult to reconcile with the findings of MELLOR AND CRAIG~O who reported the nickel(I1) &elate of 8-hydroxyquinoline to be paramagnetic and hence presumably tetrahedral. On the other hand, the nickel(I1) chelate of o-aminophenol was found to be diamagnetic and therefore of square coplanar configuration. From these magnetic data one would predict that the bonds in the o-aminophenol chelate would be stronger, giving rise to greater References
$.
213.
W.
208
VOL. 11
D. JOHNSTON, H. FREISER
chelate stability, than those of the 8-hydroxyquinoline reverse has been found to be true ‘JJl. Further, the nickel(II) N was exceptional
in that it was the only one of the class
(1954)
chelate. Actually the 8-hydroxyquinolinate 0
‘Ni’
chelates
(where 0
0’ ‘N originally formed part of a phenolic hydroxide group) which displayed paramagnetism in the MELLOR AND CRAIG investigation. If we assume that the magnetic properties in nickel(I1) chelates of alkylated 8-hydroxyquinolines are similar, another anomaly becomes apparent. The fact that the nickel chelates should be tetrahedral, according to magnetic data, is not in agreement with the argument relating to the unusual position of nickel(II) in the hindered metal stability sequence. These inconsistencies together with the fact that the nickel(I1) chelate of 8-hydroxyquinoline is an exception to the MRLLOR AND CRAIG classification indicate the need for either a re-examination of the magnetic susceptibility of this compound or the re-interprctationzof the magnetic data. 4-Methy/ - 8 - Hydroxyquinoline 34
-AH
2 -Methyl-8
- Hydroxyquinoh’ne
t
32 * 30. ae2624 222018. 161472L*
18 AH
22
*
24
*
26
“0’
26
30
32 34
-A Fwc
*
36
1.
38
CU
1614-
61mA20
’
22
’
24
’
’ 30’ 32* 34’
26 28
-
Fig.
3,
-A&PC
Fig. 3 shows a plot of AF ZISAH for the chelates of the two reagents. It can be seen that these curves are roughly linear. This is to say that the entropy changes of the reactions have caused no unusual effects within the chelates of a given reagent, It is reasonable to assume that this will generally be true. That is, the entropy of a series of chelation reactions with a given reagent will either be roughly constant or will vary regularly. This accounts for the widespread success of the log K, second ionization potential graphs. Each time this relationship References
p.
2r3.
VOL.
ORGANIC
11 (1954)
ANALYTICAL
REAGENTS
cog
VI
4 is used it is implicitly
assumed that only bond strength or AH is involved. Since the second ionization potential may be taken as a rough estimation of the average electron attracting power of a divalent metal ion it will also be a measure of the attracting power of that divalent metal ion for a source of electrons such as are found in the chelate ligand groups. Hence it is more nearly related to bond energy and AH of chelation than to AF as measured by the log of the chelate formation constant. It should also be noted that since in many cases changes in AH are accompanied by opposing changes in AS, AH should be a more sensitive index of bond strength than AF. Fig. 4 shows a plot of the second ionization potential of the gaseous metal atom against AH of chelation. As can be seen the relationship is satisfactory.
0 ._ 8 .- 77 B s 16
.z18t
751 6’
’
8
’ ’ ’ ’ 70 12 14 76 78 20 22 24 26 28 30 32 34 36 -bka1. ’
’
’
*
’
’
’
’
-
’
Fig. 4. Very few determinations of AH have been published for neutral chclate species. However, a number of such investigations have been made on the metal ammine type of complexes. Of particular interest are the very comprehensive studies by JONASSEN et nZ.l*on complexes of diethylenetriamine and triethylenetetramine. The heat of reaction sequence for the complexes of both of these reagents follow the normal stability order rather closely. MCINTYRE'~ gives a rather complete compilation of the thermodynamic functions for metal amminc formation. In practically all cases listed the heats of reaction for complexes of a given reagent follow the normal metal stability sequence. A comparison of the AH values listed in Table II points out the importance
of steric hindrance in all chelates of 2-methyl-S-hydroxyquinoline. In all cases the heats of chelation for chelates of z-methyl-8-hydroxyquinoline are only about half that found associated with the respective chelates of the 4-methyl compound. Refere?rcesp. 2r3.
W.
2 IO
I>.
JOHNSTON,
Ii.
FREISER
VOL.
11
(xg!jd$)
This points to a very large decrease in bond strength due to the hindrance of the a-methyl group. It is reasonable to assume that the presence of the 2-methyl group prevents the close approach of the two reagent molecules around the metal atom regardless of the preferred configuration and thus causes the chelate bonds considerable strain. Therefore, one might reasonably expect the magnitude of the steric effect to increase with decreasing metal size. For this purpose, the thermodynamic functions of the exchange reaction, MQ, + 2L- t ML, + eQwhere Q- and L- arc the anions of z-methyl- and 4-methyl-8-hydroxyquinoline, respectively, were plotted against the divalent crystal radii of the metals involved. Although there arc a number of conflicting sets of values of crystal radii, those of PAULING~~ were used mainly because these are the values used also by POWELL AND LATINIER~~ in similar thermodynamic calculations. As can be seen from Fig. 5 the TAS term varies regularly with the crystal radius while with AF and
;I ;----_:. , *-?zIc
.
0.70
Fig.
5.
Comparison
of
0.74 0.76 . 0.72 Crystal radius in A0 (after Pouting)
thermodynamic functions of alkylzlted 8-hydroxyquinolinc chelates.
AH there are sharp discontinuties in the case of nickel(H). In agreement with our expectations, the exchange reactions become more exothermic as the crystal radii decrease. It appears that in the case of nickel(I1) a minimum radius has been reached and steric hindrance is excessive. This helps to explain the decreased stability that has been observed in the case of hindered nickel(I1) chelates. Rcfevences
+.
213.
VOL.
11
ORGANIC
(1954)
ANALYTICAL
REAGENTS
VI
2x1
The relatively large entropy increase associated with the formation of chelates of 2-methyl&hydroxyquinoline was rather surprising but not unprecedented. FRANK AIUD EVANS’s have interpreted the following data in terms of steric hindrance and the structure-making influence on water of the reagents and complexes. AH = -rg kcal., AS = -22 Ni+2(aq.) + 6NH,(aq.) ---L [Ni(Nl-I,),Jf2(aq.); e. u. Ni+z(aq.) + GCH,NH,(aq.) =+ [Ni(CH,NH&jC2 (aq.); AH = tg.7 kcal., AS = +73 e. u. Although the complexes formed are of about equal stability spatial requirements prevent the close approach of the six methylamines about the nickel(H) ion according to arguments presented by CALVIN and discussed by FRANK AND EVANS. This easily accounts for the more positive heat of reaction in the hindered case. This is in qualitative agreement with the heat effects presented in this paper. The relatively large positive entropy term in the hindered reaction above is also in agreement with the results prcscntcd hcrc. The interpretations of these entropy effects are rather difficult. FRANK AND EVANS
attributed
only
a small
part
of
the
greater
entropy
of
formation
of
the
greater size relative to that of the ~i(NH,),Jf2 complex. PJW-WH2M They speculated that the larger part of the entropy contribution might be due to the imposition of iccbcrg-forming tendencies on the methylaminc molecule by the methyl group which would lower its partial molal entropy compared with that of ammonia. If the water structure breaking tcndcncies of both complexes were similar this would account for the greater entropy of reaction in the second case. Although not strictly comparable, the thermodynamic data of SPIKE AND YARRF which describe the formation of [Cd(NH,),lC2 and [Cd(CH,NH2),]+2 tend to cast doubt on this explanation. No steric hindrance was found in the cadmium(II) complexes as evidenced by similar values of AH. Of greater importance is the fact that the AS values of SPIKE AND PARRY do not reflect a trend
+2 to
similar
to
The
that
large
its
reported
difference
by
methyl-&hydroxyquinoline methyl
groups
able
to
partially
such
shielding
would
reduce
atoms
and
remnant An
in
the
the
a first
interesting
for
chelates
is
probably
of
polar
a chelate
oxygen, in the
forces
the
formation best
of
greater
of
explained
z-methylby
and
metal
and
noting
S-hydroxyquinolinc
nitrogen,
corresponding
between
much
coordination
alternative
EVANS.
values
be possible solvent
AND
AS
z-positions
the attractive
allow of
the shield
would
FRANK
in the
atoms
4-methyl-chelate.
4-
that
might
be
while
no
Shielding
the solvent
molecules
and
these
polar
frceclom
by
virtually
eliminating
any
sought
in
sphere. explanation
can
be
the
inverse
relation
-AH and AS. In the case of the exchange reactions mentioned above and also in the case of formation of 4-methyl-S-hydroxyquinolinc chelates there seems to be a linear inverse relation between -AH and AS. It is probable that as the bond strength decreases, the freedom of motion involving the atoms so of
Referewes
p.
2x3.
W. D. JOHNSTON,
2x2
H.
FREISER
VOL.
‘1
(r9s-I )
bonded, increases. This increased freedom of motion may well result in a significant entropy increase. This relationship involving intramolecular forces may be considered analogous to that of BARCLAY AND BUTLER~~ (linear relation of AH and AS of vaporization of certain non-ionized solutes) in which intermolecular, non-valence, forces are involved. It would bc interesting to study the entropies and heats of chelation for chelates II
the
financial
support
of the U.
S. Atomic
SUMMARY The CALVIN-RJERRUM titration technique for the determination of chelate formation constants has been applied to the Cu(II), Ni II). Co(II), Zn(I1) and Mn(I1) were made at chelates of z- and 4-methyl-8-hydroxyquinoline. L easurements several temperatures in order to evaluate Al-f and AS values of chelation. The results obtamed were interpreted in terms of steric hindrance of the z-methyl group. In all cases the heats of formation of the chelates of 2-methyl-8-hydrox quinoline were remarkably more ositive than those for the corresponding chelates o Y +methyl8-hydroxyquinoline, 3%. IS large difference in the strengths of the metal-chelate bonds is apparently due to the hindrance of the methyl groups which prevent the close grouping of the two reagent molecules around the metal in chelates of z-methyl8-hydrpxyguinoline, The lower bond strength in chelates of 2-methyl-8-hydroxycompensated by a relatively larger entropy of formation. umohne IS partially % his is attributed to decreased solvent chelate interaction caused by the shielding of the polar 0, N, and metal atoms by the z-methyl groups. The determination of chelate formation constants of z-phenyl-8-hydroxyqumolme has been carried out to further cxtcnd our study of stcric effects in metal chelates. Rl%UMfi Les auteurs ont appliqud la technique de titrage de CALVIN-BJERRUM B la determination dcs constantes de formation des complexes des 2- et +m&hyl-8hydroxyquinol&nes avec Cu, Ni, Co, Zn et Mn. Les mesures ont Bt6 faites tLdiverses temperatures afin d’8valucr les valeurs de chelation AH et AS. Les rQultats obtenus ont &l! interprCtds en fonction de l’empechement stdrique provenant du Les chaleurs de formation des complexes de la 2-methylou ement 2-m&hyl. Er -hy x roxyquinol&ne sont toujours nettemcnt plus positives que celles des comReferences
p.
ar3.
VOL.
11
(19%)
ORGANIC ANALYTICAL
REAGENTS
VI
2 I.3
plexes de la 4-methyl-8-hydroxyquinoleine. Cette grande difference dans les forces de liaison entre le metal et la molecule organique semble Btre due B l’empechement sterique des groupes m&h 1 rendant t&s difficile la formation de chelate. La faible force de liaison dans les c Helates de la z-methyl-8-hydroxyquinol&ine est partiellement corn ensee par une relativement grande entropie. Cette derniere est attribuee a un affai +Zlissement de l’interaction entre le dissolvant et le complexe, les groupements methyliques empechant le reactif de reagir avcc l’oxygene, l’azote et les atomes metalliques polariscs. Les autcurs ont determine egalement les constantes de formation des complexes dc la z-phcnyl-S-hydroxy uinoleine pour faire mieux comprendre l’importance des facteurs stdriques dans 3 e domaine des complexes. ZUSAMMENFASSUNG Die CALVIN-BJnRRUMSche Titrationsmethode zur Bestimmung der Komplexbildungs-Konstanten ist auf die Cu II)-, Ni(II)-, cO(II)-. Zn(II)- und Mn(II)-Komplexe \ins angewandt worden. Es wurden Messungen bei des 2- und 4-Methyl-8-oxychino verschiedenen Temperaturen ausgefiihrt urn die Wiirmettinung und die Entropic Die Ergcbnisse sind im Sinne sterischer der Komplexbildung zu bestimmen. Hinderung der 2-Methylgruppe gedeutct worden. In ailen Fallen war die Warmetonung dcr Komplexbildung der Komplexe dcs 2-M&h 1-8-oxychinolins bcmerkenswert positiver als die derselben Komplexe des 4- d ethyl-8-oxychinolins. Dieser der Metall-Komplexbindungen ist angrosse Untcrschied der Bindungsenergie scheinend auf die iMethyl ruppen, die einc dichte Gruppierung der zwei Rcagenz2-Methyl-8-oxychinolinkomplcx verhindern, molckiile um das Meta P1 in dem Die kleincre Bindungsenergie der Komplexe des 2-Methyl-gzuriickzuftihren. kompensiert. Dies wird oxychinolins ist teilweise durch eine grossere Entropie gedeutet als eineverminderung der \Vcchselwirkung zwischen Komplex und Losung?: mittel durch Abschirmung der polarcn 0, N, und Metallatomen durch die Methylgruppen. Die Bildungskonstantcn der Komplcxe des 2-Phen I-8-oxychinolins wurdcn bestimmt urn unsere Forschung der sterischen Effekte in X9etallkomplexcn weiter ausendchnen. REFERENCES 1
7
Ii. G. CHAKLES AND H. FREISER, Anal. C/rim. Ada, 11 (1954) IOI. Abstracted from the thesis submitted by \V. D. JOI-INSTON in partial fulfillment of Pittsburgh, of the requircmcnts for the Ph. D. degree at the University address : \Vcstinghousc Research Laboratories, East August 1953. Present Pittsburgh, Pa., U.S.A. L. L. MRRRKTT AND J. Ii. ~VALKDR, Ind. E?rg. Clrcrn., Anal. Ed., 16 (1944) 387. E. J. BUTLER AND M. F. RING, J. C/rein. Sot., (x949) 1489. H. IRVING, New York, rg5 1. A. ALBERT, SelEcttvc To.uicLty, John \Vlley & Suns Inc., \\‘. D. JO~INSTON AND l-l. FREISER, J. Al/r. C/rem. Sot., 74 (x952) 5239. 1. P. PJIILLIXS, 11. L. ELI~J~NGEIZ AND I,. L. MoRRIr*r Jr., J. Am. Chem. Sot..
8
1949) 3986. REISDR, 12. G. ‘H’.B
2
3 4 5 0
0 10 11 12 la
CHARLES AND \V. 13. JOIKNSTON,J. Am. C/rem. Sot., 74 (rg52) 1383. Ii. G. CRARLES AND I-1. FR&ISI:R, Anal. Chtm. Ada, IK (1954) I. D. I?. MELLOR AND D. P. CRAIG, J. ~'YOC. Roy. SOL. N. S. Wales, 74 (1940) 475. Ii. G. CHARLES AND H. FRRIsRR, J. Ant. C/rem. Sot., 74 ISIBOIIM,J. PlbyS. I-I-Ieza. JONASSEN, G. G. HURST, l<. 13. LRBLANC AND A. VV. 56 (1953) 16. G. H:' MCINTYRB, Ph. D. Thesis, The Pennsylvania State College (x953 . lVatzwe of the Chemical Bad, Cornell University Press, It h aca,
17 C. G. 18 1. M.
AND \V. M. LATIMER, J. Chem. Whys., Ig (1951) 1139. M. \V. EVANS, J. C/bent. Plbys., 13 (1945) 507. SPIKE AND R. \V. PARRY, J. Am. Clbewz. Sot., 75 (1953) 2726 BARCLAY AND J . A. V. BUTLER, l’umas. Faraday So&, 34 (1938) 1451. Received January ~rth, rg54