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Studies on the stability of beryllium chelates of catecholdisulphonic and chromotropic acids in aqueous solutions
123
STUDIES ON THE STABILITY OF BEKYLLIUM CHELATES OF CATECHOLDISULPHONIC AND CHROMOTROPIC ACIDS IN AQUEOUS SOLUTIONS
The complexation reactions of beryllium with the disodium salts disulphonic and chromotropic acids have been studied by potcntiometric tometric methods. In the acidic range, the formation of mono-derivatives only was both the qstems even in the presence of a fargc excess of cllelatin~ ~~uilil~riunl constants “K” of the reaction : Be+-+&As-
of catecholand conducindicated in agents. The
F’ BeA”-fzH+-
and the logaritl~ms of the formation constant fk’lj of the complex anion /Be.“l+] in catech~l~sulphonic and chromotropit acids systems were found to be 9.33 x TO-~ and 13.23, and 2.95 x IO-~ and 16.43, respectively. The l(~~arithms of the formation constants (log Kn} in both the systems for I:I and I:Z complexes have also been determined by the alternative methods to be 13.52 +O.IZ; 12_55-t_o.14 and 16.89& 0.09; rg.()I+o.xI respectively.
The disodium salts of catecl~oldisulI~honic ~tiran) and cl~ro~lr~tropi~ acids are good chelating agents and form strong chelates with a number of metallic ions’-8. From a perusal of the literature, it was evident that no ~~llysic~~-chemical measurements have been made on the beryllium-catecholdisulphttnic and berylliumclrrtrmotropic acids systems. It was therefore considered of interest to make a detaiicd stutl~ on these systems by potentiometric and conductometric techniques.
Solutions of B.D.H. Analar beryllium sulphate, BeSO 14 Hz0 and E. &kck btryHium nitrate, Be(NO& - 4 HsO were prepared by direct weighing and the - ---. .Kurukshetrs l.:nirersitv, ~~uruksb~tr~, Punjab * I’rcscnt address: Chemistry ~epartmcnt, (India).
S. N. DUBEY, R. C. MEHROTRA
124
strengths were checked gravimetrically by precipitation with ammonium hydroxide and subsequent ignition to oxide. The disodium salts of catecholdisulphonic (C~~~Na~O~S~ ’ H20) and chromoE. Merck products. tropic acids [&Hu (SOaNa)z(OH)z * zH20] were recrystallized The solutions of these chelating agents were also prepared by direct weighing. All solutions were made up with conductivity water.
pH measurements were carried out at room temperature with a Cambridge Bench pH meter standardized against 0.05 M potassium hydrogen phthlate solution. A Philips magic-eye-type (P.R. 9500) conductivity bridge was used for conductivity measurements. POTE~TIOMETRI~ STUDIES The potentiometric titration of beryllium sulphate with a solution of tiron is shown by curves x,2 and 3 (Fig. I). The chelation between beryllium ions and tiron may be represented as :
2 40 --w--x-
2.00
1
1
2
3
4
5I
3.oo--r.-
Molesof
Moles of tiron added
NaCtI added Cm)
5’
6
I
7
Fig.
I. Curve I, potentiometric titration of beryllium sulphate (0.005 34, so ml) with tiron (0.10 M) ; curves 2 and 3, above titrations after the addition of I and a ml of 0.10 A4 sodium hydroxide to
the beryllium sulphate solution.
Fig. 2. Curves o, I, 2 and 3 represent potentiometric titrations of beryllium sulphate (0.005 M, 20 ml) with sodium hydroxide (0.1 M) in the presence of o, I, 2 and 4 ml of tiron (0.10 M) ; m = mole of base added per mole of beryllium ion. J. Less-Commow
Metals,
g (1965)
123-132
STABILITY
OF BERYLLIUM
CHELATES
I25
Curve o {Fig. z), for the potentiometri~ titration of beryllium sulphate with caustic soda, is similar to that obtained by RRITTON~. The potentiometric titration of be ryllium sulphate with caustic soda in the presence of equimolar concentration of tiron (curve I, Fig. 2) shows an inflexion at m= 2 (m =moles of base added per mole of beryllium ion) indicating the formation of beryllium mano-catecholdisulphonate derivative of the type (A). Be”-’ +
--OaS/\OH
---(&S//yO\ +zOH-F=*
0’
u
-2>Be j + z I-120
so3 (4
A second inflexion at m=3, anion of the type (B).
is probably due to the formation of the hydrolysable
(W The titration of beryllium sulphate with sodium hydroxide in the presence of moles of tiron (curve 2, Fig. 2) shows inflexion at m=2 and m=4, indicating the formation of I :I and I :2 compfexes which can be explained by the following simple reaction : 2
-03s ReZ-b
OH
+
OH
0 SOS-
Similar titration in the presence of 4 moles of tiron (curve 3, Fig. 2) shows inflexion at m =2, indicating the formation of I :I complex. A second inflexion at m=6, occurring at pH about 9.50, corresponds to the neutralization of excess of tiron present in the system. The potentiometric titrations of beryllium sulphate with chromotropic acids shown in curves I, 2 and 3 (Fig. 3) indicate chelation which may be represented as: B@, HO
Re
7%
OH
Be”+ +
+ * Ht
-0sS
The potentiometric titration of basic beryllium sulphate with chromotropic acid (curve 2, Fig. 3) shows inflexion at m==o.s (m represents moles of chromotropic acid added per mole of be~ll~urn ion), indicating the following reaction : j. ihs-CoWZt?ZOW
M&k,
9 (1965)
I23--132
S. N. DUBEY, R. C. MEHROTRA
126 Be+ HO
Pe@Wl+ t
flexion type :
OH
A -OsS\
‘I
/c
sos--
-_.
sH&so
3
_+HzO
3
Similar titration of beryllium hydroxide with chromotropic acid shows inat rn=r, indicating the formation of monochromotropate derivative of the / HO Be@%
OH ~o~s~s*~~
+_oz;asOs_+
The reactions representing derivatives are represented below : (4 HO
OH
the
+ 2 HzO
formation
of mano-
and
_ /JYo
12$
Bee+ +
dichromotropate
2 I-r20
SOS_
03-S 0J
(Curves I, 2 and 3 ; Fig. 4)
Moles
of chromotrc,nC
acid
Moles Of MOH
added
(ml
Fig. 3. Curve x, potentiometric titration of beryllium sulphate (o.005 M, so ml) with chromotropic acid (0.10 M); curves 2 and 3, above titrations after the addition of I and z ml of 0.10 M sodium hydroxide to the beryllium sulphate solution. Fig. 4. Curves o, I, 2 and 3 represent potentiometric titrations of beryllium sulphate (0.005 M, 20 ml) with sodium hydroxide (0.10 iM) in the presence of o, I, z and 3 ml of chromotropic acid (0.10 M) ; m = mole of base added per mole of beryllium ion. J. Less-Common Metals,
g (1965) 123-132
STABILITY
OF BERYLLIUM
CHELATES
=27 OH
0))
3
I Be
f
OH-
-
(Curve I, Fig. 4) HO
Be”+ + 2
r~,
OH
+4
()H-
+
/SO,-
-0ss
(Curve 2, Fig. 4)
CONDUCTOMETRIC
STUDIES
The conductometric titrations of beryllium sulphate with sodium hydroxide in the presence of o, I, z and 4 moles of tiron (curves o, I, z and 3; Fig. 5) and o, I, z and 3 moles of chromotropic acid (curves o, I, z and 3; Fig. 6) showed inflexions in corroboration with those obtained from potentiometric titrations. When sodium hydroxide was replaced by the weak base, ammonium hydrox-
00 I
1
2
3
4
5L
6
7/
Moles of alkoII added Cm)
Fig. 5. Curves o, I, 2 and 3 represent conductometric titrations of beryllium sulphate (20 ml, 0.005 M) with 0.1 111NaOH in the presence of o, I, L and 4 ml 0.1 jzI tiron; curves o’, I’, L’ and 3’ represent conductometric titrations of beryllium sulphate (zo ml, 0.005 AW)with 0.1 M NH&H in the presence of o, I, 2 and 4 ml 0.1 M tiron. ,I. Less-Common Metals,
g (1965) 123-132
128
S. N. DIJBEY,
R. C. MEHROTRA
ide, the above titrations yielded very sharp inflexion points (curves 0’, I’, 2’ and 3’; Figs. 5 and G), as the excess of ammonium hydroxide does not affect the observed conductivity of the systems appreciably.
Fig. 6. Curves 0, I, 2 and 3 represent conductometric titrations of beryllium sulphate (0.005 hf, 20 ml) with 0.10 M NaOH in the presence of o, I, zz and 3 ml (0.10 M) chromotropic acid; curves o‘, I’, 2’ and 3’ represent conductometric titrations of beryllium sulphate (0.005 M, 20 ml) with 0.10 M NH40H in the presence of o, I, 2 and 3 ml (0.10 M) chromotropic acid. EQUILIBRIUM
AND FORMATION
CONSTANTS
(I) Of all the possible equilibria for the formation of monocatecholdisulphonate and monochromotropate derivatives, only the following equilibrium Bezt- + HsA2- +BeAz- + zH+ (I) was found to give constant pK values. For the above calculations, the ionic strength was kept approximately constant by using 0.10 M potassium nitrate medium and low concentrations of ligand and metal ions. The equilib~um constant “I-i” for the reaction (I) is given by: K
=
[BeA2-I @+I2 [I3e’J+] [HzAZ-]
and the values of “K” were calculated from the various concentrations of [BeAz-] [H+] [Besf] and [HsA’+] present in the equilibrium mixture. Over the pH range studied, the concentrations of [HAS-] [A4-] [OH-] and [H+] were negligible compared with the other ionic species present in the equilibrium mixture. The values of pK were calculated at various points on the curves I and z (Figs. 7 and 8) from m= o to WC=1.2 and from m= o to m= 1.8respectively, where nz represents mole of base added per mole of the beryllium ion. The values of pK calculated from the data of curves I and z (Figs. 7 and 8) are given in Tables I and II. After determining the equilibrium constants “K” of the reaction (I), the forJ. Less-Common
Metals,
g (1965) 123-132
STABILITY
OF
BERYLLIUM
ml of
NaOH
CHELATES
129
ml
(O.OlM)
of
kCCh
!OOl MI
Fig. 7. Potentiomctric titrations of beryllium sulphatc with sodium hydroxide in the presence of tiron at 0.10 Mionic strength: curve I, 5 x 10-4~lf in bcryIlium sulphate and 5 x IO-~ Al in tiron; curve 2, 5 x 10-4 M in beryllium sulphate and I i 10-3 II in tiron; total volume of the initial mixture = 50 ml.
Fig. 8. I’otentiometric titrations of beryllium sulphatc with sodium hydroxide in the presence of chromotropic acid at 0.10 M ionic strength : curve I, 5 x ro-.i M in beryllium sulphatc and 5 x 10-4 ilf in chromotropic acid; curve 2, 5 x IO- 4 111 in beryllium sulphatc and I x 10-3 Al in chromotropic acid; total volume of initial mixture = 50 ml.
TAHLI:
1 EQUILIURIUhI CONSTANTS
T.1BLE FOR THE
II CONSTANTS
EQUILIBRI”M
FOR
BERYLLIUh~FMONOCATECHOLDISUL-
BERRYLLII!~l~~lONOCHROMOTROPATE
PHONATE
DERIVATIVE
1.01. of
iYUOH (FEl)
DERIVATIVE
tivmz vcctio (c1we I)
Fov 1:~ Betivo,r m/i0 (cwue 2)
PH
PK
PH
PS
4.49 4.65 4.76
7.07
F;ou I:1
Be-
0.2
+02
7.07
0.0
4.73
7.00
I.0
$‘,2
1.4 I.8 2.0
5.02
5.13 5.19
2.4 5.31 2.8 5.43 3.2 5.57 .\lean pK value
* Neglcctcd.
7.00
7.00 j.00 7.00 7.03 7.06 ___ (7.12)* 7.01
4.8i
4.96 5.01 .j.IO
5.20 5.31
7.04 7.05 7.04 7.04 7.0(’ 7.08 (7.‘2)* (7.16)* 7.05
0.2
1.73
0.0
3.80
I.0
3.8
1.4 I.8
3.95 4.03
2.0
1.07
2.4 2.Y 3.2 3.6 4 0 4.4 &Lean ph’
.+.rC, 4.26 4.38 4.53
4.58 4.56 4.56 4.55 4.55 4.54 -l..ib 4.57 4.49 (4.64)*
3.62 3.68 3.74 3.79 3.84 3.87 3.93 4.00 4.08 4.17
zf.27 4.41 value
4.56
THE
4.58 4.58 4.57 4.55 4.52 4.52 4.50 4.50 4.50 4.50 4.48 (4.44)* 4.53
S. N. DDBEY,
130
R. C. MEHROTRA
mation constant Kl of the complex anion (Be&z-) was determined from the following equation :
K
=
Kl
Ku1 - Kaz
Substituting the values of K (9.33 x IO-~), Ku1 10 (2.188 x IO-*) and Ku2 for beryllium-catecholdisulphonic acid system in eqn. (2), the value of log KI = 13.23 was obtained. Similarly, by substituting the values of K (2.95 x IO-~), Ku1 11 (4.365 x 10-6) and Ku2 (2.512 x 10-16) for the beryllium-chromotropic acid system in eqn. (2), the value of log KI= 16.43 was obtained. (2.512
x 10-13)
(2)
Co~~~tat~a~ offo~~zat~o~ ~o~sta~~tsfro~ the functions Fi ami [A *-I The formation constants KI and KZ for the systems (Be2ffA4+ BeAzand BeA2- i- Ad- + BeA+) were calculated from the titrations of tiron or chromotropic acid and nitric acid in the absence and presence of beryllium ions (curves I and 2, Fig. 9; curves I and 2, Fig. IO respectively); the ionic strength of the medium was
2.5i
2
4 6 8 ml NaOH (0.02M)
Fig. 9. Curve I, potentiometric
curve 2, sameascurve I plus 4
10
I 12
titration of 50 ml (5 x
2.0c
10-3
x 10-4 M beryllium nitrate.
;
8 6 4 mi NaOii(O.02M)
M tiron + 2.5 x
110-3
M
L 1012
nitric acid)
;
Fig. IO. Curve I, potentiometric titration of 50 ml (5 x IO-~ M chromotropic acid + 2.5 x 10-3 M nitric acid) ; curve 2, same as curve I plus 4 x IO-* M beryllium nitrate.
kept approximately constant by using 0.10 M potassium nitrate solution. The values of G (average number of ligands per metal ion) were plotted against the values of pA (Figs. II and 12). The values of log X1 and log Kz were calculated by (i) interpolation at half ~5 values, (ii) spreading factor method, (iii) correction term method and (iv) least square treatments. The final values of log RI and log KZ are given in Table III.
STABILITY OF BERYLLIUM CHELriTES
11.5
12.0
125
130
13.5
131
140
15.5
-log (A‘%
16.0
16.5
170
1225
-iog~A~-)
l;ig. I I. Degree of formation of Z as a function of -log (ligand) for berplliun-catecholdisulphonate chelate. Fig. 12. Degree of formation of n as a function of -log (ligand) for beryllium-chromotropate TABLE
chelate.
I11
COMPUTATION
OFFORMATION
CONSTANTS
IV&hods employed
(I) Interpolation at half n valuesl*~‘3 (iii) Spreading factor method’? (iii) Correction term method’” (iv) Least square treatment’” (a) Curve fitting method (b) Simultaneous equation method _
Be~catecholdisulphonic acid system
Be-chromotropic acid system
log KI
log Kz
log KI
log Kz
‘3.57 ‘3.40 13.49
12.46 12.08 I2._52
16.97 16.81 ~6.85
15.89 15.97 15.92
13.63 13.46
r2.40 IL.00
16.98 16.81
15.80 ‘5.97
ACKNOWLEDGEMENT
One of the authors (S.N.D.) is indebted to the University New Delhi, for providing a research fellowship.
Grants Commission,
KEFEliENCES I Y.>~URAKAMI AND A. E.MARTELL,J. Am. Chem.Soc., 82 (1960) 5605. 2 I<. NASANEN, &omen Kemistilehti, 3rB (1958) 19. 3 B.I. INTORREANDA.E.MAKTELL, J.Am.Chem.Soc.,82 (1960)358. 4 F. WILLI AND J. H. YOE, Anal. Chim. Acta, 8 (1953) 546. 5 X. OKAC AND L. SOMMER, Z. Anal. Chem., 163 (1958) 412. 6 L. SOMMER, Z. Anal. Chem., 164 (1958) zgg. 7 0. JANTTI, Suomen Kemistilehti, 3oB (1957) 136. 8 A.E. MARTELL, S.CHABAREK,JR.,R.C.COURTNEY, S. WESTERBACK AND H.HYYTIAINEN,J. Am. Chem. Sot., 79 (1957) 3036.
J, Less-Common Metals, g (1965) 123-132
S. N. DUBEY, R. C. MEHROTRA
132
9 H. T. S. BRITTON, Hydrogen Ions, Vol. II, Chapman and Hall, London, 1956, p. 72. IO G. SCHWARZBNBACH AND A. WILLI,H~V. Chim. Ada,34 (1951) 528. II J. HELLERANDG.SCHWARZENBACH,~~~V. Chim.Acta,gq (1951)1876. 12 J, BJERRUM, Metal Ammine Formation in Aqueous Solution, P. Haase and Son, Copenhagen, 194’. 13 M. CALVIN AND K. W. WILSON, J. Am. Chem. Sot., 67 (1945) 2003. 14 H.IRVINGAND H.S.ROSSOTTI,J. Chem.Soc., (Ig53)3397. J. L~~s-~~rnrn~
Metals,
9 (x965)
123-132
STUDIES ON THE STABILITY OF BEKYLLIUM CHELATES OF CATECHOLDISULPHONIC AND CHROMOTROPIC ACIDS IN AQUEOUS SOLUTIONS
The complexation reactions of beryllium with the disodium salts disulphonic and chromotropic acids have been studied by potcntiometric tometric methods. In the acidic range, the formation of mono-derivatives only was both the qstems even in the presence of a fargc excess of cllelatin~ ~~uilil~riunl constants “K” of the reaction : Be+-+&As-
of catecholand conducindicated in agents. The
F’ BeA”-fzH+-
and the logaritl~ms of the formation constant fk’lj of the complex anion /Be.“l+] in catech~l~sulphonic and chromotropit acids systems were found to be 9.33 x TO-~ and 13.23, and 2.95 x IO-~ and 16.43, respectively. The l(~~arithms of the formation constants (log Kn} in both the systems for I:I and I:Z complexes have also been determined by the alternative methods to be 13.52 +O.IZ; 12_55-t_o.14 and 16.89& 0.09; rg.()I+o.xI respectively.
The disodium salts of catecl~oldisulI~honic ~tiran) and cl~ro~lr~tropi~ acids are good chelating agents and form strong chelates with a number of metallic ions’-8. From a perusal of the literature, it was evident that no ~~llysic~~-chemical measurements have been made on the beryllium-catecholdisulphttnic and berylliumclrrtrmotropic acids systems. It was therefore considered of interest to make a detaiicd stutl~ on these systems by potentiometric and conductometric techniques.
Solutions of B.D.H. Analar beryllium sulphate, BeSO 14 Hz0 and E. &kck btryHium nitrate, Be(NO& - 4 HsO were prepared by direct weighing and the - ---. .Kurukshetrs l.:nirersitv, ~~uruksb~tr~, Punjab * I’rcscnt address: Chemistry ~epartmcnt, (India).
S. N. DUBEY, R. C. MEHROTRA
124
strengths were checked gravimetrically by precipitation with ammonium hydroxide and subsequent ignition to oxide. The disodium salts of catecholdisulphonic (C~~~Na~O~S~ ’ H20) and chromoE. Merck products. tropic acids [&Hu (SOaNa)z(OH)z * zH20] were recrystallized The solutions of these chelating agents were also prepared by direct weighing. All solutions were made up with conductivity water.
pH measurements were carried out at room temperature with a Cambridge Bench pH meter standardized against 0.05 M potassium hydrogen phthlate solution. A Philips magic-eye-type (P.R. 9500) conductivity bridge was used for conductivity measurements. POTE~TIOMETRI~ STUDIES The potentiometric titration of beryllium sulphate with a solution of tiron is shown by curves x,2 and 3 (Fig. I). The chelation between beryllium ions and tiron may be represented as :
2 40 --w--x-
2.00
1
1
2
3
4
5I
3.oo--r.-
Molesof
Moles of tiron added
NaCtI added Cm)
5’
6
I
7
Fig.
I. Curve I, potentiometric titration of beryllium sulphate (0.005 34, so ml) with tiron (0.10 M) ; curves 2 and 3, above titrations after the addition of I and a ml of 0.10 A4 sodium hydroxide to
the beryllium sulphate solution.
Fig. 2. Curves o, I, 2 and 3 represent potentiometric titrations of beryllium sulphate (0.005 M, 20 ml) with sodium hydroxide (0.1 M) in the presence of o, I, 2 and 4 ml of tiron (0.10 M) ; m = mole of base added per mole of beryllium ion. J. Less-Commow
Metals,
g (1965)
123-132
STABILITY
OF BERYLLIUM
CHELATES
I25
Curve o {Fig. z), for the potentiometri~ titration of beryllium sulphate with caustic soda, is similar to that obtained by RRITTON~. The potentiometric titration of be ryllium sulphate with caustic soda in the presence of equimolar concentration of tiron (curve I, Fig. 2) shows an inflexion at m= 2 (m =moles of base added per mole of beryllium ion) indicating the formation of beryllium mano-catecholdisulphonate derivative of the type (A). Be”-’ +
--OaS/\OH
---(&S//yO\ +zOH-F=*
0’
u
-2>Be j + z I-120
so3 (4
A second inflexion at m=3, anion of the type (B).
is probably due to the formation of the hydrolysable
(W The titration of beryllium sulphate with sodium hydroxide in the presence of moles of tiron (curve 2, Fig. 2) shows inflexion at m=2 and m=4, indicating the formation of I :I and I :2 compfexes which can be explained by the following simple reaction : 2
-03s ReZ-b
OH
+
OH
0 SOS-
Similar titration in the presence of 4 moles of tiron (curve 3, Fig. 2) shows inflexion at m =2, indicating the formation of I :I complex. A second inflexion at m=6, occurring at pH about 9.50, corresponds to the neutralization of excess of tiron present in the system. The potentiometric titrations of beryllium sulphate with chromotropic acids shown in curves I, 2 and 3 (Fig. 3) indicate chelation which may be represented as: B@, HO
Re
7%
OH
Be”+ +
+ * Ht
-0sS
The potentiometric titration of basic beryllium sulphate with chromotropic acid (curve 2, Fig. 3) shows inflexion at m==o.s (m represents moles of chromotropic acid added per mole of be~ll~urn ion), indicating the following reaction : j. ihs-CoWZt?ZOW
M&k,
9 (1965)
I23--132
S. N. DUBEY, R. C. MEHROTRA
126 Be+ HO
Pe@Wl+ t
flexion type :
OH
A -OsS\
‘I
/c
sos--
-_.
sH&so
3
_+HzO
3
Similar titration of beryllium hydroxide with chromotropic acid shows inat rn=r, indicating the formation of monochromotropate derivative of the / HO Be@%
OH ~o~s~s*~~
+_oz;asOs_+
The reactions representing derivatives are represented below : (4 HO
OH
the
+ 2 HzO
formation
of mano-
and
_ /JYo
12$
Bee+ +
dichromotropate
2 I-r20
SOS_
03-S 0J
(Curves I, 2 and 3 ; Fig. 4)
Moles
of chromotrc,nC
acid
Moles Of MOH
added
(ml
Fig. 3. Curve x, potentiometric titration of beryllium sulphate (o.005 M, so ml) with chromotropic acid (0.10 M); curves 2 and 3, above titrations after the addition of I and z ml of 0.10 M sodium hydroxide to the beryllium sulphate solution. Fig. 4. Curves o, I, 2 and 3 represent potentiometric titrations of beryllium sulphate (0.005 M, 20 ml) with sodium hydroxide (0.10 iM) in the presence of o, I, z and 3 ml of chromotropic acid (0.10 M) ; m = mole of base added per mole of beryllium ion. J. Less-Common Metals,
g (1965) 123-132
STABILITY
OF BERYLLIUM
CHELATES
=27 OH
0))
3
I Be
f
OH-
-
(Curve I, Fig. 4) HO
Be”+ + 2
r~,
OH
+4
()H-
+
/SO,-
-0ss
(Curve 2, Fig. 4)
CONDUCTOMETRIC
STUDIES
The conductometric titrations of beryllium sulphate with sodium hydroxide in the presence of o, I, z and 4 moles of tiron (curves o, I, z and 3; Fig. 5) and o, I, z and 3 moles of chromotropic acid (curves o, I, z and 3; Fig. 6) showed inflexions in corroboration with those obtained from potentiometric titrations. When sodium hydroxide was replaced by the weak base, ammonium hydrox-
00 I
1
2
3
4
5L
6
7/
Moles of alkoII added Cm)
Fig. 5. Curves o, I, 2 and 3 represent conductometric titrations of beryllium sulphate (20 ml, 0.005 M) with 0.1 111NaOH in the presence of o, I, L and 4 ml 0.1 jzI tiron; curves o’, I’, L’ and 3’ represent conductometric titrations of beryllium sulphate (zo ml, 0.005 AW)with 0.1 M NH&H in the presence of o, I, 2 and 4 ml 0.1 M tiron. ,I. Less-Common Metals,
g (1965) 123-132
128
S. N. DIJBEY,
R. C. MEHROTRA
ide, the above titrations yielded very sharp inflexion points (curves 0’, I’, 2’ and 3’; Figs. 5 and G), as the excess of ammonium hydroxide does not affect the observed conductivity of the systems appreciably.
Fig. 6. Curves 0, I, 2 and 3 represent conductometric titrations of beryllium sulphate (0.005 hf, 20 ml) with 0.10 M NaOH in the presence of o, I, zz and 3 ml (0.10 M) chromotropic acid; curves o‘, I’, 2’ and 3’ represent conductometric titrations of beryllium sulphate (0.005 M, 20 ml) with 0.10 M NH40H in the presence of o, I, 2 and 3 ml (0.10 M) chromotropic acid. EQUILIBRIUM
AND FORMATION
CONSTANTS
(I) Of all the possible equilibria for the formation of monocatecholdisulphonate and monochromotropate derivatives, only the following equilibrium Bezt- + HsA2- +BeAz- + zH+ (I) was found to give constant pK values. For the above calculations, the ionic strength was kept approximately constant by using 0.10 M potassium nitrate medium and low concentrations of ligand and metal ions. The equilib~um constant “I-i” for the reaction (I) is given by: K
=
[BeA2-I @+I2 [I3e’J+] [HzAZ-]
and the values of “K” were calculated from the various concentrations of [BeAz-] [H+] [Besf] and [HsA’+] present in the equilibrium mixture. Over the pH range studied, the concentrations of [HAS-] [A4-] [OH-] and [H+] were negligible compared with the other ionic species present in the equilibrium mixture. The values of pK were calculated at various points on the curves I and z (Figs. 7 and 8) from m= o to WC=1.2 and from m= o to m= 1.8respectively, where nz represents mole of base added per mole of the beryllium ion. The values of pK calculated from the data of curves I and z (Figs. 7 and 8) are given in Tables I and II. After determining the equilibrium constants “K” of the reaction (I), the forJ. Less-Common
Metals,
g (1965) 123-132
STABILITY
OF
BERYLLIUM
ml of
NaOH
CHELATES
129
ml
(O.OlM)
of
kCCh
!OOl MI
Fig. 7. Potentiomctric titrations of beryllium sulphatc with sodium hydroxide in the presence of tiron at 0.10 Mionic strength: curve I, 5 x 10-4~lf in bcryIlium sulphate and 5 x IO-~ Al in tiron; curve 2, 5 x 10-4 M in beryllium sulphate and I i 10-3 II in tiron; total volume of the initial mixture = 50 ml.
Fig. 8. I’otentiometric titrations of beryllium sulphatc with sodium hydroxide in the presence of chromotropic acid at 0.10 M ionic strength : curve I, 5 x ro-.i M in beryllium sulphatc and 5 x 10-4 ilf in chromotropic acid; curve 2, 5 x IO- 4 111 in beryllium sulphatc and I x 10-3 Al in chromotropic acid; total volume of initial mixture = 50 ml.
TAHLI:
1 EQUILIURIUhI CONSTANTS
T.1BLE FOR THE
II CONSTANTS
EQUILIBRI”M
FOR
BERYLLIUh~FMONOCATECHOLDISUL-
BERRYLLII!~l~~lONOCHROMOTROPATE
PHONATE
DERIVATIVE
1.01. of
iYUOH (FEl)
DERIVATIVE
tivmz vcctio (c1we I)
Fov 1:~ Betivo,r m/i0 (cwue 2)
PH
PK
PH
PS
4.49 4.65 4.76
7.07
F;ou I:1
Be-
0.2
+02
7.07
0.0
4.73
7.00
I.0
$‘,2
1.4 I.8 2.0
5.02
5.13 5.19
2.4 5.31 2.8 5.43 3.2 5.57 .\lean pK value
* Neglcctcd.
7.00
7.00 j.00 7.00 7.03 7.06 ___ (7.12)* 7.01
4.8i
4.96 5.01 .j.IO
5.20 5.31
7.04 7.05 7.04 7.04 7.0(’ 7.08 (7.‘2)* (7.16)* 7.05
0.2
1.73
0.0
3.80
I.0
3.8
1.4 I.8
3.95 4.03
2.0
1.07
2.4 2.Y 3.2 3.6 4 0 4.4 &Lean ph’
.+.rC, 4.26 4.38 4.53
4.58 4.56 4.56 4.55 4.55 4.54 -l..ib 4.57 4.49 (4.64)*
3.62 3.68 3.74 3.79 3.84 3.87 3.93 4.00 4.08 4.17
zf.27 4.41 value
4.56
THE
4.58 4.58 4.57 4.55 4.52 4.52 4.50 4.50 4.50 4.50 4.48 (4.44)* 4.53
S. N. DDBEY,
130
R. C. MEHROTRA
mation constant Kl of the complex anion (Be&z-) was determined from the following equation :
K
=
Kl
Ku1 - Kaz
Substituting the values of K (9.33 x IO-~), Ku1 10 (2.188 x IO-*) and Ku2 for beryllium-catecholdisulphonic acid system in eqn. (2), the value of log KI = 13.23 was obtained. Similarly, by substituting the values of K (2.95 x IO-~), Ku1 11 (4.365 x 10-6) and Ku2 (2.512 x 10-16) for the beryllium-chromotropic acid system in eqn. (2), the value of log KI= 16.43 was obtained. (2.512
x 10-13)
(2)
Co~~~tat~a~ offo~~zat~o~ ~o~sta~~tsfro~ the functions Fi ami [A *-I The formation constants KI and KZ for the systems (Be2ffA4+ BeAzand BeA2- i- Ad- + BeA+) were calculated from the titrations of tiron or chromotropic acid and nitric acid in the absence and presence of beryllium ions (curves I and 2, Fig. 9; curves I and 2, Fig. IO respectively); the ionic strength of the medium was
2.5i
2
4 6 8 ml NaOH (0.02M)
Fig. 9. Curve I, potentiometric
curve 2, sameascurve I plus 4
10
I 12
titration of 50 ml (5 x
2.0c
10-3
x 10-4 M beryllium nitrate.
;
8 6 4 mi NaOii(O.02M)
M tiron + 2.5 x
110-3
M
L 1012
nitric acid)
;
Fig. IO. Curve I, potentiometric titration of 50 ml (5 x IO-~ M chromotropic acid + 2.5 x 10-3 M nitric acid) ; curve 2, same as curve I plus 4 x IO-* M beryllium nitrate.
kept approximately constant by using 0.10 M potassium nitrate solution. The values of G (average number of ligands per metal ion) were plotted against the values of pA (Figs. II and 12). The values of log X1 and log Kz were calculated by (i) interpolation at half ~5 values, (ii) spreading factor method, (iii) correction term method and (iv) least square treatments. The final values of log RI and log KZ are given in Table III.
STABILITY OF BERYLLIUM CHELriTES
11.5
12.0
125
130
13.5
131
140
15.5
-log (A‘%
16.0
16.5
170
1225
-iog~A~-)
l;ig. I I. Degree of formation of Z as a function of -log (ligand) for berplliun-catecholdisulphonate chelate. Fig. 12. Degree of formation of n as a function of -log (ligand) for beryllium-chromotropate TABLE
chelate.
I11
COMPUTATION
OFFORMATION
CONSTANTS
IV&hods employed
(I) Interpolation at half n valuesl*~‘3 (iii) Spreading factor method’? (iii) Correction term method’” (iv) Least square treatment’” (a) Curve fitting method (b) Simultaneous equation method _
Be~catecholdisulphonic acid system
Be-chromotropic acid system
log KI
log Kz
log KI
log Kz
‘3.57 ‘3.40 13.49
12.46 12.08 I2._52
16.97 16.81 ~6.85
15.89 15.97 15.92
13.63 13.46
r2.40 IL.00
16.98 16.81
15.80 ‘5.97
ACKNOWLEDGEMENT
One of the authors (S.N.D.) is indebted to the University New Delhi, for providing a research fellowship.
Grants Commission,
KEFEliENCES I Y.>~URAKAMI AND A. E.MARTELL,J. Am. Chem.Soc., 82 (1960) 5605. 2 I<. NASANEN, &omen Kemistilehti, 3rB (1958) 19. 3 B.I. INTORREANDA.E.MAKTELL, J.Am.Chem.Soc.,82 (1960)358. 4 F. WILLI AND J. H. YOE, Anal. Chim. Acta, 8 (1953) 546. 5 X. OKAC AND L. SOMMER, Z. Anal. Chem., 163 (1958) 412. 6 L. SOMMER, Z. Anal. Chem., 164 (1958) zgg. 7 0. JANTTI, Suomen Kemistilehti, 3oB (1957) 136. 8 A.E. MARTELL, S.CHABAREK,JR.,R.C.COURTNEY, S. WESTERBACK AND H.HYYTIAINEN,J. Am. Chem. Sot., 79 (1957) 3036.
J, Less-Common Metals, g (1965) 123-132
S. N. DUBEY, R. C. MEHROTRA
132
9 H. T. S. BRITTON, Hydrogen Ions, Vol. II, Chapman and Hall, London, 1956, p. 72. IO G. SCHWARZBNBACH AND A. WILLI,H~V. Chim. Ada,34 (1951) 528. II J. HELLERANDG.SCHWARZENBACH,~~~V. Chim.Acta,gq (1951)1876. 12 J, BJERRUM, Metal Ammine Formation in Aqueous Solution, P. Haase and Son, Copenhagen, 194’. 13 M. CALVIN AND K. W. WILSON, J. Am. Chem. Sot., 67 (1945) 2003. 14 H.IRVINGAND H.S.ROSSOTTI,J. Chem.Soc., (Ig53)3397. J. L~~s-~~rnrn~
Metals,
9 (x965)
123-132