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Study of the properties of lead tetraacetate as oxidizing and oxidimetric reagent in analytical chemistry
MICROCHEMICAL
JOURNAL
Study
183-192
11,
of the
(1966)
Properties
as Oxidizing
and
of
Oxidimetric
in Analytical III.
Equilibria
of
the
Hydrochloric VLADIMIR Department
DVOGK,
of Analytical
IVAN Chemistry, Received
Lead
Tetraacetate Reagent
Chemistry System
Pb(lV)/Pb(ll)
Acid
Medium
NI~MEC,
AND
Charles August
University,
JAROSLAV Prague,
in
Z+KA Czechoslovakia
30, 1965
The results described in the preceding communication (see communication II of this series) have formed the necessarybase for an attempt of a partial solution of the equilibrium of the system Pb(IV)
+ 2 Cl- + Pb(I1) + Cln.
It was mainly the observation, that polarography with a rotating platinum electrode allows the determination of the tetravalent lead concentration in the presenceof chlorine, in a medium of hydrochloric acid; moreover the observation, that lead tetraacetate converts in a medium of hydrochloric acid into the [PbC1612- ion was also a very significant aid. It has been known for a long time from preparative inorganic chemistry, that a suspensionof PbClz in HCl converts into yellow-colored solutions when a stream of chlorine is passed through it, a suitable cation when added precipitating a solid chloroplumbate(IV) of the composition X2[PbCls] (X+ = NH4f, K+, Csf, and others). Still, however, the problem of the Pb(IV)/Pb( II) equilibrium in hydrochloric acid media was not fully solved and ‘was,therefore, the object of the study presented. EXPERIMENTAL
Reagents
An approximately 0.05 M lead tetraacetate solution in glacial acetic acid was prepared in the way described in the preceding papers (see, e.g., communication II of this series) and its concentration determined iodometritally (0.50 ml of the lead tetraacetate solution was added to an excessof 183
184
v.
DvoKAK,
I.
NEMEC,
AND
J.
Z~KA
potassium iodide solution and the iodine liberated was titrated by 0.01 N W&O3 ) . All other chemicals used were of reagent-grade purity; the hydrochloric acid was twice azeotropically distilled. Apparatus All pipettes and burettes used were officially calibrated. The method of polarographic measurements is described in part II of this series. The first to study this problem quantitatively was Wescott (8). He studied the equilibrium between solid PbC12 and gaseous chlorine, which was located above the solution at a given pressure, in a closed system in the acidity range of 0.1 to 0.8 M HCl. The author mentioned considered the following equations for the reaction taking place in the system described : PbClz (s) + Cl2 (g) + 2 Cl- + [PbCle]“(1) PbCls (s) + Cl2 (g) +
Cl- Z+ [PbClj12-
(2)
with the respective equilibria Kl=
[PbC1,2-]/[Cl-]2*
[pcl,],
K2 = [PbC1,2-]/[C1-].
[pClz],
where [ ] denotes the concentration of the respective ions in the solution in molarities, and [pC12] the partial pressure of chlorine in the closed system (in atmospheres). The concentration of bivalent lead in the solution was taken by the author to be constant and was, therefore, included in the reaction constant. On the basis of experimental data the author came to the conclusion that in dilute hydrochloric acid solutions the reaction proceeds mainly according to Eq. 2 ; he obtained the value of K2 = 0.055 for the equilibrium constant of this reaction. The author concludes from these observations that in dilute HCl solutions tetravalent lead exists mainly in the form of the [PbCIS] - ion. Szychlinski et al. (7) have studied qualitatively the equilibrium Pb(IV)
+ 2 Cl- + Pb(I1)
+ Cl2
in a medium of 6 to 1OM HCl. They have studied the variation of the concentrations of bivalent and tetravalent lead in hydrochloric acid solution through which a stream of chlorine passed, in dependence on the acidity of the solution and on the temperature. They have found that the ratio of the concentrations Pb(II)/Pb(IV) increases regularly with the increasing acidity (in 6 to 1OM HCl the value is 0.1 to 0.3), in 10 M HCI
III
185
it increases with the increasing temperature from 10” to 50” in of 0.2 to 2.0. The authors prove from this increase of the ratio dence on the increasing temperature, that the reaction Pb (IV) Pb(I1) is endothermic under these conditions. Heal and May (5) have studied spectrophotometrically the equilibrium constant expressed by the relation
the range in depen+ 2 e F?
STUDY
K'=
OF
THE
PROPERTIES
[total Pb(II)].
OF
LEAD
TETRAACETATE.
[total C12],/[total
apparent
Pb(IV)].
The symbols [total concentration] denote here the molar concentration of all possibly present forms of ions and complexes of chlorine, bivalent and tetravalent lead in the solution. The authors have determined the apparent equilibrium constants for 1 J4 HCl or for mixtures of HCl and HClOd in a total concentration of 1 M. They have found that for a temperature of 19.5’C the following relations are valid: 0.1 M 0.35il4 0.7il4 l.OM
HCl . . . K'= HCl . . . K'= HCl . . . K'= HCl . . . K'=
1700.lo-” 83. lo-’ 22.6. 1O-4 12.3*10-4.
In our work we have attempted to determine the value of the apparent constant K', on the one hand for concentrations lower than 1 1M HCl, mainly however for higher concentrations, where this constant has not yet been determined. We have started out from the assumption that after the conversion of lead tetraacetate to the chlorocomplex of tetravalent lead, acetate ions added with the reagent will not participate in the course of the equilibrium itself. The chlorocomplex formed reacts with chloride ions in the solution, chlorine and bivalent lead being formed. If this reaction takes place in a closed system, equilibrium is established after some time. A knowledge of the equilibrium concentrations of tetravalent and bivalent lead and of the chlorine concentration in the solution then allows the determination of the above defined constant K'. The closed system was realized by means of a flask, blackened on the outside and fitted with a well-sealing ground joint. The flask was filled to the top, so that after being closed it contained nothing but the liquid. Equilibrium was generally achieved within a time shorter than 1 hour, but still the mixture was kept in the flask in a thermostat at 20.0” + O.l”C for several hours. We have determined the equilibrium concentration of tetravalent lead in the solution polarographically, using a rotating platinum electrode. It
186
v.
DvoiiA~,
I.
N~EC,
AND
J.
2+~.4
has been found advantageous to anodize the platinum electrode first at +2.0 v in a solution of 1O-4 M Pb(IV) in 11 M HC104 (see the preceding communication II of this series) for a period of 8 minutes, since then the tetravalent lead wave is better developed on the polarographic curves, and is easier to measure. In an equilibrium solution (about 60 ml) at least four curves were recorded, one immediately after the other in each case. Although the magnitude of the sum of limiting currents, corresponding to the reduction of chlorine and of tetravalent lead decreased with time, it has been found that the magnitude of the limiting current corresponding to the tetravalent lead concentration is practically constant during a period of 12 minutes at a temperature of 20°C and all the variation is due to the chlorine wave only. The precise height of the wave, corresponding to tetravalent lead in the equilibrium solution, has been obtained from four measured values by extrapolating to the time corresponding to the opening of the flask containing the equilibrium solution. After this polarographic recording, an electrode pre-polarized in the same way as in the preceding case was used to record a wave corresponding to the known concentration of the tetravalent lead chlorocomplex in 50 ml of a hydrochloric acid solution of the same concentration as in the equilibrium mixture under study. (In practice, 0.50 ml of a saturated lead tetraacetate solution of known concentration was added to 49.5 ml of a solution of HCI of a given concentration in a beaker with a rotating platinum and saturated calomel electrodes, and a polarogram was registered immediately after the components had been mixed. By comparison of the wave-heights of the equilibrium mixture and of the known amount of the tetravalent lead chlorocomplex (determined iodometrically), the concentration of tetravalent lead in the equilibrium solution was determined. The following holds for the tetravalent lead concentration in the equilibrium solution: [Pb(IV)]
= +j
lo-4M,
where a equals wave-height in mm, corresponding to 0.50 ml saturated lead tetraacetate in 50.00 ml of a solution of the given HCI concentration; b equals wave-height in mm, corresponding to tetravalent lead in the solution; d equals consumption (ml) of a 0.01 N Na2S203 solution for 0.50 ml saturated lead tetraacetate; f equals titer of the 0.01 N Na&03 volumetric solution. The chlorine concentration in the equilibrium solution was determined as follows: 50.00 ml of the equilibrium solution were added to excess
STUDY
OF
THE
PROPERTIES
OF
LEAD
TETRAACETATE.
III
187
iodide, and the sum of tetravalent lead and chlorine concentrations determined by titrating with 0.01 N NazSaOa. The equilibrium chlorine concentration is obtained after subtracting the respective tetravalent lead concentration: [CL] = (+;)~lO-‘92, where c equals consumption (ml) of a 0.01 N Na&Oa solution for 50.00 ml of the equilibrium solution; a, b, f, d equal known expressionsdefined in the preceding paragraph. The bivalent lead determination was based on the following consideration: since the saturated lead tetraacetate solution used for these aims contained no bivalent lead, all the bivalent lead in the equilibrium must originate in the reaction of tetravalent lead with the Cl- ions. In general the bivalent lead concentration, with respect to the reaction scheme Pb(IV) + 2 Cl- + Pb(I1) + Cl, must be equal to the chlorine concentration in the solution. However, although a bi-distilled hydrochloric acid azeotrope was used for these experiments, and the lead tetraacetate solution was prepared from the substance four times crystallized and from distilled acetic acid, still the solution in a filled equilibration flask contained after 2 to 5 hours a sum of oxidizing agents lower by several per cent compared to the solution originally prepared from lead tetraacetate. This difference, which is probably causedby a slight degree of loss of chlorine through the ground joint, or by a possiblereaction of chlorine or Pb(IV) with traces of oxidizable substancesin the solution, must be equal to the increase of concentration of bivalent lead in the equilibrated solution. This means, that the equilibrium concentration of bivalent lead is equal to the sum of the equilibrium concentration of chlorine and the decreaseof the oxidant concentration. Since the equilibrium solution was always prepared in such a way that 250.0 ml of the hydrochloric acid solution contained 2.00 ml saturated lead tetraacetate, and the sum of tetravalent lead and chlorine was always titrated in 50.00 ml of this solution (after 2 to 5 hours), the respective concentration decreaseof oxidizing substancesis given by the expression (d - 1.25 C) X f X 10h4M, and the total bivalent lead concentration in the equilibrium solution is given by bd
[Pb(II) ] =
d - 0.2% - -
a
f
10-4M.
188
V. DVOkhK,
I.
NEMEC,
AND
J.
Z+KA
After insertion of these concentrations into the relation for the apparent equilibrium constant we obtain the relation
K’ zz
(
d - 0.25~ -
x-9 bd a
f10-4M.
All constants K’ have been calculated from this formula under the experimental conditions described, by insertion of the respective values. The procedure of the determination of the constants was as follows: By pipette 2.00 ml is measured from a tempered and precisely filled volumetric flask of 250 ml volume, containing a HCl solution of the given concentration; 2.00 ml saturated lead tetraacetate in acetic acid are then added. -After mixing the solution is immediately transferred into a blackened volumetric flask of 100-ml volume. The flask is closed with a ground stopper and kept in a thermostat at 20.0” 4 O.l”C for 2 to 5 hours. After this time, the flask is opened and 50.00 ml of the solution are transferred by pipette as quickly as possible into excess potassium iodide solution, and the rest from the flask is transferred into a beaker with platinum electrode and saturated calomel electrode (the beaker serving as polarographic vessel), and a polarographic curve is registered. The value of b is determined from four polarographic ‘curves, recorded one after the other, after reading the height of waves corresponding to Pb(IV) reduction and extrapolating them to the time at which the flask has been opened. The consumption of 0.01 N thiosulfate solution for titration of the iodine, liberated after transferring 50.00 ml of the equilibrium solution into excess iodide, gives the value of c. A titration of 0.50 ml of saturated lead tetraacetate in excess iodide by means of thiosulfate gives the value of d. The value of a is obtained from polarographing 50.0 ml of hydrochloric acid of the same concentration as in the case of the equilibrium, and containing 0.50 ml of saturated lead tetraacetate solution (the recording must be done always immediately after mixing the components). The preliminary experiments have shown that in concentrations lower than 1 M HCI the wave corresponding to tetravalent lead is relatively small; with the decreasing concentration of HCI in the equilibrium solution this wave decreases, so that in the evaluation large errors were necessarily incurred. We have therefore measured the constants mainly in equilibrium solutions of a higher acid concentration that 1 M. Only a few
STUDY
OF
THE
PROPERTIES
OF
LEAD
TETRAACETATE.
III
189
constants have been determined at concentrations lower than 1 M, in order to permit comparison with literature data (5). For instance, for the equilibrium mixture in 0.8 M HCl (the solution being 0.2 M HC104 at the same time) we have obtained as mean value of three measurements K’ = 14.7 * lOwa, which agrees by order of magnitude with the results of the authors (5). The results obtained for concentrations of 1 to 4.5M HCl are given in Table 1. It is to be seen from the table, that the mean values of the constants do not vary in the range of 2 to 4.5 M HCI (within the limits of experimental error). Measurements at concentrations higher than 4.5 M ‘were impossible, since under these conditions it was impossible to resolve precisely the polarographic waves corresponding to the reduction of tetravalent lead from the chlorine waves. All measurements were carried out at a temperature of 20.0’ f O.l”C. Measurements could not be carried out at higher temperatures, since under the given experimental arrangement (polarographing in an open beaker) it was impossible to measure the height of the lead wave in the equilibrium solution, which at higher temperatures decreases even during the first few minutes of recording a polarogram, due to the far more rapid establishment of equilibrium and greater degree of volatility of the chlorine. We did not succeed in constructing an apparatus which would have eliminated quantitatively this difficulty-chlorine volatilization in the course of recording the polarograms. Due to the influence of perchloric acid on the tetravalent lead wave in the equilibrium mixture (the wave becomes less distinct, and its height is difficult to evaluate), constants have not been measured in a constant ionic strength solution (4.5 M). However, we have found orientatively that in a medium of 2 M HCl + 2.5 M HCIOl the mean value of the equilibrium constant was 5.4. lOua; i.e., a value which permits the assumption that the values of the constants mentioned in the acidity range of 2 to 4.5 M HCl will not differ substantially with a variation of the ionic strength of the solution of acidities in the range of 2 to 4.5 M. From the invariance of the K’ constant in the range of 2 to 4.5 M HCl, it follows that in the reaction considered chloride ions do not participate. Before attempting to explain the scheme of equilibrium, let us consider the forms in which ions may be present in this acidity region. From the absorption spectra of the equilibrium solutions, which were fully identical with the spectra described for [PbC1612- (see communication I of this series), it follows that all the tetravalent lead will be probably in the form of [PbC16]“ions in this acidity range. Under the condition that the constant of equilibrium between C12/C13-
TABLE EQUILIBRIUM
1
CONSTANTS IN TRE SYSTEM Pb(IV)-Pb(II)-Cl2
MEDIUM OF 1 TO 4.94 (K’=
~Pb(II)I~[CI,l/IPb(IV)I, equilibrium concentrations
HCI
Time
(M)
(hrs)
1.0
1.5
2 .o
2.5
3.0
3.5
4.0
4.5
where [Pb(II)l, [Pb(IV)I, [Cl,1 are the in d4.104 at a temperature of 20.0 2 0.1 Co)
Pb(IV) (M
x 104)
IN A
HCI
(M
Cl2 x 104)
Pb(I1) (M
x 104)
K’ x 104
4 5 4 3 5
1.09 1.04 1.06 0.996 1.09
3.34 3.12 3.39 3.46 3.54
3.62 3.77 3.74 3.80 3.95
4 3 5 3 2
1.38 1.50 r.81 1.63 1.47
3.23 3.13 2.84 3.43 3.42
3.41 3.28 4.13 3.68 3.50
8.00 6.84 6.48 7.74 8.14
7.44
3 2 4 5 2
1.52 1.58 1.18 1.19 1.76
3.15 3.13 2.59 2.57 3.20
3.21 3.11 265 2.60 3.19
6.65 6.16 5.82 5.61 5.80
6.01
3 4 3 2 5
1.63 1.19 1.22 1.66 1.27
3.07 2.62 2.57 3.18 2.48
3.03 2.63 2.60 3.16 2.56
5.70 5.79 5.48 6.05 5.00
5.60
5 2 2 3 3
1.92 1.65 1.61 1.60 1.26
3.29 3.15 3.16 3.08 2.53
3.38 3.02 3.06 3.03 2.52
5.79 5.76 6.00 5.83 5.06
5.69
2 3 3 5 4
1.67 1.73 1.26 1.08 1.20
3.01 2.93 2.51 2.74 2.74
3.15 3.00 2.55 2.70 2.55
5.68 5.08 5.20 6.85 5.82
5.73
3 3 2 5 4
1.54 1.51 1.77 1.12 1.17
3.11 3.15 2.96 2.74 2.60
3.13 3.16 2.94 2.65 2.59
6.32 6.59 4.92 6.48 5.75
6.02
2 2 5 3 4
1.68 1.66 1.25 1.72 1.79
3.25 3.22 2.53 3.20 3.13
3.29 3.16 2.51 3.15 3.08
6.36 6.13 5.08 5.86 5.38
5.76
190
11.2 11.3 12.0 13.2 12.8
K’ x 104 mean value
12.1
STUDY
OF
THE
PROPERTIES
OF
LEAD
TETRAACETATE.
III
191
holds for this acidity region, and that it is uninfluenced by acetate ions (present in the equilibrium solution in a concentration of approximately 0.1 n/r) this means, that for instance in 3 M HCl chlorine is present to about 60% in the form of Clz molecules and to 40% in the form of CLions. Therefore, we must consider both forms of chlorine in this range. An unambiguous decision about the form of complex ion in which bivalent lead is present is difficult. It is known, that in concentrations higher than 6 M HCl lead exists mainly in the form of the [PbCL]‘ion (3, 6). At concentrations lower than 6 M, however, complexes of the type [ PbC13] - as well as complexes with a smaller number of chloride ions may occur (1, 2, 3, 6). More detailed quantitative relations between these forms of chlorocomplexes in HCl media are not exactly known. The whole problem of bivalent lead is complicated because at acetate ion concentrations of the order of 10-l M (the case of the equilibria studied) acetate complexes of the type [PbAc] - exist already (4). On the other hand, we must also keep in mind that these acetate complexes are considered and proved in pure acetate solutions (4), i.e., in the absence of chlorides. Since in our case we have a ZO-fold to 40-fold excess of chloride ions compared to the acetate ions, and because the great complexing capacity of chloride ions is well known, it may be assumed that these acetate complexes will probably not participate. It is impossible, however, to decide on the measure in which the ratios of the complexes [PbClr12and [PbCla] - will exist in the given acidity range of 2 to 4.5 M HCl. Based on these considerations, the conditions for the presence of possible other forms of the components present in the equilibrium solution, and under the condition that in the equilibrium studied no free chloride ions will participate (this condition follows from the constants determined), the equilibrium may be expressed in general by means of the following equations: [PbClo12- e [PbC14]*- + Cl2 (a) (b)
[PbClc]‘-
z+ [PbCla]“-
+ Cl:<-.
A comparison of the mean value of the constant K’ in 1 M HC1 in Table 1 with the value given by the authors of (5) shows clearly, that both constants agree. The increase of values of constants for the range of 2 to 1 M HCl observed in the present work and mainly at concentrations lower than 1 M shows that in the equilibrium mentioned chloride ions start to participate at concentrations lower than 1.5 M HCl. In this case the above scheme of the equilibrium (Eqs. (a) and (b) ) is invalid;
192
V.
DVOHAK,
I.
NeMEC,
AND
J.
ZkKA
chloride complexes of bivalent lkad probably are formed with less than three coordinated chloride ions, and probably even tetravalent lead begins to occur in the form of [PbC&] -, as proved by Heal and May (5) as well as by Wescott (8). SUMMARY Variations of the equilibrium constant K’ = lPb(I1) 1. [Cl,I/[Pb(IV) I have been studied in dependence on the acidity of the solution, based on the polarographic behavior and determination of the concentration of tetravalent lead in the presence of chlorine. A possible scheme of the equilibrium is discussed, based on the invariance of this constant in the range of 2 to 4.5 M HCI and comparison with literature data. REFERENCES 1. 2.
3.
4.
I., PARTON, H. N., AND ROBINSON, R. A., The constitution of the lead halides in aqueous solution. J. Am. Chem. Sot. 77, 5844-5848 (1955). FROMHERZ, H., AND LIH, K. H., Spektroskopische Untersuchung der Dissoziationsverhaltnisse von Blei- und Thallohalogeniden in wassriger LBsung. Z. Phys. Chew A 153, 321-375 (1931). GLASNER, A., AND AVINUR, P., Spectrophotometric methods for the determination of impurities in pure and analytical reagents. II. Some absorption spectra in concentrated chlorines and their applications. Talanta 11, 761-773 (1964). GOBOM, S., The complex formation between lead(H) ions and acetate ions. Acta Broos,
Chem. 5. 6.
7. 8.
A.
&and.
17,
2181-2189
(1963).
H. G., AND MAY, J., A spectrophotometric study of the stability of lead(IV) in hydrochloric acid solutions. J. Am. Chew Sot. 80, 2374-2377 (1958). MERRITT, C., JR., HERSHENSON, H. M., AND ROGERS, L. B., Spectrometric determination of bismuth lead and thallium with hydrochloric acid. Anal. Ckem. 25, 572-577 (1953). SZYCHLINSKI, J., LATOWSKI, T., AND KOREWA, R., A study of the equilibria in the Pb4f - Clsystem. Roczniky Chem. 32, 1013-1023 (1958) (in Polish). WESCOTT, E. W., The equilibrium between chlorine and plumbous and plumbic chlorides in aqueous solution. J. Am. Chem. Sot. 42, 1335-1349 (1920). HEAL,
JOURNAL
Study
183-192
11,
of the
(1966)
Properties
as Oxidizing
and
of
Oxidimetric
in Analytical III.
Equilibria
of
the
Hydrochloric VLADIMIR Department
DVOGK,
of Analytical
IVAN Chemistry, Received
Lead
Tetraacetate Reagent
Chemistry System
Pb(lV)/Pb(ll)
Acid
Medium
NI~MEC,
AND
Charles August
University,
JAROSLAV Prague,
in
Z+KA Czechoslovakia
30, 1965
The results described in the preceding communication (see communication II of this series) have formed the necessarybase for an attempt of a partial solution of the equilibrium of the system Pb(IV)
+ 2 Cl- + Pb(I1) + Cln.
It was mainly the observation, that polarography with a rotating platinum electrode allows the determination of the tetravalent lead concentration in the presenceof chlorine, in a medium of hydrochloric acid; moreover the observation, that lead tetraacetate converts in a medium of hydrochloric acid into the [PbC1612- ion was also a very significant aid. It has been known for a long time from preparative inorganic chemistry, that a suspensionof PbClz in HCl converts into yellow-colored solutions when a stream of chlorine is passed through it, a suitable cation when added precipitating a solid chloroplumbate(IV) of the composition X2[PbCls] (X+ = NH4f, K+, Csf, and others). Still, however, the problem of the Pb(IV)/Pb( II) equilibrium in hydrochloric acid media was not fully solved and ‘was,therefore, the object of the study presented. EXPERIMENTAL
Reagents
An approximately 0.05 M lead tetraacetate solution in glacial acetic acid was prepared in the way described in the preceding papers (see, e.g., communication II of this series) and its concentration determined iodometritally (0.50 ml of the lead tetraacetate solution was added to an excessof 183
184
v.
DvoKAK,
I.
NEMEC,
AND
J.
Z~KA
potassium iodide solution and the iodine liberated was titrated by 0.01 N W&O3 ) . All other chemicals used were of reagent-grade purity; the hydrochloric acid was twice azeotropically distilled. Apparatus All pipettes and burettes used were officially calibrated. The method of polarographic measurements is described in part II of this series. The first to study this problem quantitatively was Wescott (8). He studied the equilibrium between solid PbC12 and gaseous chlorine, which was located above the solution at a given pressure, in a closed system in the acidity range of 0.1 to 0.8 M HCl. The author mentioned considered the following equations for the reaction taking place in the system described : PbClz (s) + Cl2 (g) + 2 Cl- + [PbCle]“(1) PbCls (s) + Cl2 (g) +
Cl- Z+ [PbClj12-
(2)
with the respective equilibria Kl=
[PbC1,2-]/[Cl-]2*
[pcl,],
K2 = [PbC1,2-]/[C1-].
[pClz],
where [ ] denotes the concentration of the respective ions in the solution in molarities, and [pC12] the partial pressure of chlorine in the closed system (in atmospheres). The concentration of bivalent lead in the solution was taken by the author to be constant and was, therefore, included in the reaction constant. On the basis of experimental data the author came to the conclusion that in dilute hydrochloric acid solutions the reaction proceeds mainly according to Eq. 2 ; he obtained the value of K2 = 0.055 for the equilibrium constant of this reaction. The author concludes from these observations that in dilute HCl solutions tetravalent lead exists mainly in the form of the [PbCIS] - ion. Szychlinski et al. (7) have studied qualitatively the equilibrium Pb(IV)
+ 2 Cl- + Pb(I1)
+ Cl2
in a medium of 6 to 1OM HCl. They have studied the variation of the concentrations of bivalent and tetravalent lead in hydrochloric acid solution through which a stream of chlorine passed, in dependence on the acidity of the solution and on the temperature. They have found that the ratio of the concentrations Pb(II)/Pb(IV) increases regularly with the increasing acidity (in 6 to 1OM HCl the value is 0.1 to 0.3), in 10 M HCI
III
185
it increases with the increasing temperature from 10” to 50” in of 0.2 to 2.0. The authors prove from this increase of the ratio dence on the increasing temperature, that the reaction Pb (IV) Pb(I1) is endothermic under these conditions. Heal and May (5) have studied spectrophotometrically the equilibrium constant expressed by the relation
the range in depen+ 2 e F?
STUDY
K'=
OF
THE
PROPERTIES
[total Pb(II)].
OF
LEAD
TETRAACETATE.
[total C12],/[total
apparent
Pb(IV)].
The symbols [total concentration] denote here the molar concentration of all possibly present forms of ions and complexes of chlorine, bivalent and tetravalent lead in the solution. The authors have determined the apparent equilibrium constants for 1 J4 HCl or for mixtures of HCl and HClOd in a total concentration of 1 M. They have found that for a temperature of 19.5’C the following relations are valid: 0.1 M 0.35il4 0.7il4 l.OM
HCl . . . K'= HCl . . . K'= HCl . . . K'= HCl . . . K'=
1700.lo-” 83. lo-’ 22.6. 1O-4 12.3*10-4.
In our work we have attempted to determine the value of the apparent constant K', on the one hand for concentrations lower than 1 1M HCl, mainly however for higher concentrations, where this constant has not yet been determined. We have started out from the assumption that after the conversion of lead tetraacetate to the chlorocomplex of tetravalent lead, acetate ions added with the reagent will not participate in the course of the equilibrium itself. The chlorocomplex formed reacts with chloride ions in the solution, chlorine and bivalent lead being formed. If this reaction takes place in a closed system, equilibrium is established after some time. A knowledge of the equilibrium concentrations of tetravalent and bivalent lead and of the chlorine concentration in the solution then allows the determination of the above defined constant K'. The closed system was realized by means of a flask, blackened on the outside and fitted with a well-sealing ground joint. The flask was filled to the top, so that after being closed it contained nothing but the liquid. Equilibrium was generally achieved within a time shorter than 1 hour, but still the mixture was kept in the flask in a thermostat at 20.0” + O.l”C for several hours. We have determined the equilibrium concentration of tetravalent lead in the solution polarographically, using a rotating platinum electrode. It
186
v.
DvoiiA~,
I.
N~EC,
AND
J.
2+~.4
has been found advantageous to anodize the platinum electrode first at +2.0 v in a solution of 1O-4 M Pb(IV) in 11 M HC104 (see the preceding communication II of this series) for a period of 8 minutes, since then the tetravalent lead wave is better developed on the polarographic curves, and is easier to measure. In an equilibrium solution (about 60 ml) at least four curves were recorded, one immediately after the other in each case. Although the magnitude of the sum of limiting currents, corresponding to the reduction of chlorine and of tetravalent lead decreased with time, it has been found that the magnitude of the limiting current corresponding to the tetravalent lead concentration is practically constant during a period of 12 minutes at a temperature of 20°C and all the variation is due to the chlorine wave only. The precise height of the wave, corresponding to tetravalent lead in the equilibrium solution, has been obtained from four measured values by extrapolating to the time corresponding to the opening of the flask containing the equilibrium solution. After this polarographic recording, an electrode pre-polarized in the same way as in the preceding case was used to record a wave corresponding to the known concentration of the tetravalent lead chlorocomplex in 50 ml of a hydrochloric acid solution of the same concentration as in the equilibrium mixture under study. (In practice, 0.50 ml of a saturated lead tetraacetate solution of known concentration was added to 49.5 ml of a solution of HCI of a given concentration in a beaker with a rotating platinum and saturated calomel electrodes, and a polarogram was registered immediately after the components had been mixed. By comparison of the wave-heights of the equilibrium mixture and of the known amount of the tetravalent lead chlorocomplex (determined iodometrically), the concentration of tetravalent lead in the equilibrium solution was determined. The following holds for the tetravalent lead concentration in the equilibrium solution: [Pb(IV)]
= +j
lo-4M,
where a equals wave-height in mm, corresponding to 0.50 ml saturated lead tetraacetate in 50.00 ml of a solution of the given HCI concentration; b equals wave-height in mm, corresponding to tetravalent lead in the solution; d equals consumption (ml) of a 0.01 N Na2S203 solution for 0.50 ml saturated lead tetraacetate; f equals titer of the 0.01 N Na&03 volumetric solution. The chlorine concentration in the equilibrium solution was determined as follows: 50.00 ml of the equilibrium solution were added to excess
STUDY
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TETRAACETATE.
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187
iodide, and the sum of tetravalent lead and chlorine concentrations determined by titrating with 0.01 N NazSaOa. The equilibrium chlorine concentration is obtained after subtracting the respective tetravalent lead concentration: [CL] = (+;)~lO-‘92, where c equals consumption (ml) of a 0.01 N Na&Oa solution for 50.00 ml of the equilibrium solution; a, b, f, d equal known expressionsdefined in the preceding paragraph. The bivalent lead determination was based on the following consideration: since the saturated lead tetraacetate solution used for these aims contained no bivalent lead, all the bivalent lead in the equilibrium must originate in the reaction of tetravalent lead with the Cl- ions. In general the bivalent lead concentration, with respect to the reaction scheme Pb(IV) + 2 Cl- + Pb(I1) + Cl, must be equal to the chlorine concentration in the solution. However, although a bi-distilled hydrochloric acid azeotrope was used for these experiments, and the lead tetraacetate solution was prepared from the substance four times crystallized and from distilled acetic acid, still the solution in a filled equilibration flask contained after 2 to 5 hours a sum of oxidizing agents lower by several per cent compared to the solution originally prepared from lead tetraacetate. This difference, which is probably causedby a slight degree of loss of chlorine through the ground joint, or by a possiblereaction of chlorine or Pb(IV) with traces of oxidizable substancesin the solution, must be equal to the increase of concentration of bivalent lead in the equilibrated solution. This means, that the equilibrium concentration of bivalent lead is equal to the sum of the equilibrium concentration of chlorine and the decreaseof the oxidant concentration. Since the equilibrium solution was always prepared in such a way that 250.0 ml of the hydrochloric acid solution contained 2.00 ml saturated lead tetraacetate, and the sum of tetravalent lead and chlorine was always titrated in 50.00 ml of this solution (after 2 to 5 hours), the respective concentration decreaseof oxidizing substancesis given by the expression (d - 1.25 C) X f X 10h4M, and the total bivalent lead concentration in the equilibrium solution is given by bd
[Pb(II) ] =
d - 0.2% - -
a
f
10-4M.
188
V. DVOkhK,
I.
NEMEC,
AND
J.
Z+KA
After insertion of these concentrations into the relation for the apparent equilibrium constant we obtain the relation
K’ zz
(
d - 0.25~ -
x-9 bd a
f10-4M.
All constants K’ have been calculated from this formula under the experimental conditions described, by insertion of the respective values. The procedure of the determination of the constants was as follows: By pipette 2.00 ml is measured from a tempered and precisely filled volumetric flask of 250 ml volume, containing a HCl solution of the given concentration; 2.00 ml saturated lead tetraacetate in acetic acid are then added. -After mixing the solution is immediately transferred into a blackened volumetric flask of 100-ml volume. The flask is closed with a ground stopper and kept in a thermostat at 20.0” 4 O.l”C for 2 to 5 hours. After this time, the flask is opened and 50.00 ml of the solution are transferred by pipette as quickly as possible into excess potassium iodide solution, and the rest from the flask is transferred into a beaker with platinum electrode and saturated calomel electrode (the beaker serving as polarographic vessel), and a polarographic curve is registered. The value of b is determined from four polarographic ‘curves, recorded one after the other, after reading the height of waves corresponding to Pb(IV) reduction and extrapolating them to the time at which the flask has been opened. The consumption of 0.01 N thiosulfate solution for titration of the iodine, liberated after transferring 50.00 ml of the equilibrium solution into excess iodide, gives the value of c. A titration of 0.50 ml of saturated lead tetraacetate in excess iodide by means of thiosulfate gives the value of d. The value of a is obtained from polarographing 50.0 ml of hydrochloric acid of the same concentration as in the case of the equilibrium, and containing 0.50 ml of saturated lead tetraacetate solution (the recording must be done always immediately after mixing the components). The preliminary experiments have shown that in concentrations lower than 1 M HCI the wave corresponding to tetravalent lead is relatively small; with the decreasing concentration of HCI in the equilibrium solution this wave decreases, so that in the evaluation large errors were necessarily incurred. We have therefore measured the constants mainly in equilibrium solutions of a higher acid concentration that 1 M. Only a few
STUDY
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TETRAACETATE.
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189
constants have been determined at concentrations lower than 1 M, in order to permit comparison with literature data (5). For instance, for the equilibrium mixture in 0.8 M HCl (the solution being 0.2 M HC104 at the same time) we have obtained as mean value of three measurements K’ = 14.7 * lOwa, which agrees by order of magnitude with the results of the authors (5). The results obtained for concentrations of 1 to 4.5M HCl are given in Table 1. It is to be seen from the table, that the mean values of the constants do not vary in the range of 2 to 4.5 M HCI (within the limits of experimental error). Measurements at concentrations higher than 4.5 M ‘were impossible, since under these conditions it was impossible to resolve precisely the polarographic waves corresponding to the reduction of tetravalent lead from the chlorine waves. All measurements were carried out at a temperature of 20.0’ f O.l”C. Measurements could not be carried out at higher temperatures, since under the given experimental arrangement (polarographing in an open beaker) it was impossible to measure the height of the lead wave in the equilibrium solution, which at higher temperatures decreases even during the first few minutes of recording a polarogram, due to the far more rapid establishment of equilibrium and greater degree of volatility of the chlorine. We did not succeed in constructing an apparatus which would have eliminated quantitatively this difficulty-chlorine volatilization in the course of recording the polarograms. Due to the influence of perchloric acid on the tetravalent lead wave in the equilibrium mixture (the wave becomes less distinct, and its height is difficult to evaluate), constants have not been measured in a constant ionic strength solution (4.5 M). However, we have found orientatively that in a medium of 2 M HCl + 2.5 M HCIOl the mean value of the equilibrium constant was 5.4. lOua; i.e., a value which permits the assumption that the values of the constants mentioned in the acidity range of 2 to 4.5 M HCl will not differ substantially with a variation of the ionic strength of the solution of acidities in the range of 2 to 4.5 M. From the invariance of the K’ constant in the range of 2 to 4.5 M HCl, it follows that in the reaction considered chloride ions do not participate. Before attempting to explain the scheme of equilibrium, let us consider the forms in which ions may be present in this acidity region. From the absorption spectra of the equilibrium solutions, which were fully identical with the spectra described for [PbC1612- (see communication I of this series), it follows that all the tetravalent lead will be probably in the form of [PbC16]“ions in this acidity range. Under the condition that the constant of equilibrium between C12/C13-
TABLE EQUILIBRIUM
1
CONSTANTS IN TRE SYSTEM Pb(IV)-Pb(II)-Cl2
MEDIUM OF 1 TO 4.94 (K’=
~Pb(II)I~[CI,l/IPb(IV)I, equilibrium concentrations
HCI
Time
(M)
(hrs)
1.0
1.5
2 .o
2.5
3.0
3.5
4.0
4.5
where [Pb(II)l, [Pb(IV)I, [Cl,1 are the in d4.104 at a temperature of 20.0 2 0.1 Co)
Pb(IV) (M
x 104)
IN A
HCI
(M
Cl2 x 104)
Pb(I1) (M
x 104)
K’ x 104
4 5 4 3 5
1.09 1.04 1.06 0.996 1.09
3.34 3.12 3.39 3.46 3.54
3.62 3.77 3.74 3.80 3.95
4 3 5 3 2
1.38 1.50 r.81 1.63 1.47
3.23 3.13 2.84 3.43 3.42
3.41 3.28 4.13 3.68 3.50
8.00 6.84 6.48 7.74 8.14
7.44
3 2 4 5 2
1.52 1.58 1.18 1.19 1.76
3.15 3.13 2.59 2.57 3.20
3.21 3.11 265 2.60 3.19
6.65 6.16 5.82 5.61 5.80
6.01
3 4 3 2 5
1.63 1.19 1.22 1.66 1.27
3.07 2.62 2.57 3.18 2.48
3.03 2.63 2.60 3.16 2.56
5.70 5.79 5.48 6.05 5.00
5.60
5 2 2 3 3
1.92 1.65 1.61 1.60 1.26
3.29 3.15 3.16 3.08 2.53
3.38 3.02 3.06 3.03 2.52
5.79 5.76 6.00 5.83 5.06
5.69
2 3 3 5 4
1.67 1.73 1.26 1.08 1.20
3.01 2.93 2.51 2.74 2.74
3.15 3.00 2.55 2.70 2.55
5.68 5.08 5.20 6.85 5.82
5.73
3 3 2 5 4
1.54 1.51 1.77 1.12 1.17
3.11 3.15 2.96 2.74 2.60
3.13 3.16 2.94 2.65 2.59
6.32 6.59 4.92 6.48 5.75
6.02
2 2 5 3 4
1.68 1.66 1.25 1.72 1.79
3.25 3.22 2.53 3.20 3.13
3.29 3.16 2.51 3.15 3.08
6.36 6.13 5.08 5.86 5.38
5.76
190
11.2 11.3 12.0 13.2 12.8
K’ x 104 mean value
12.1
STUDY
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TETRAACETATE.
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191
holds for this acidity region, and that it is uninfluenced by acetate ions (present in the equilibrium solution in a concentration of approximately 0.1 n/r) this means, that for instance in 3 M HCl chlorine is present to about 60% in the form of Clz molecules and to 40% in the form of CLions. Therefore, we must consider both forms of chlorine in this range. An unambiguous decision about the form of complex ion in which bivalent lead is present is difficult. It is known, that in concentrations higher than 6 M HCl lead exists mainly in the form of the [PbCL]‘ion (3, 6). At concentrations lower than 6 M, however, complexes of the type [ PbC13] - as well as complexes with a smaller number of chloride ions may occur (1, 2, 3, 6). More detailed quantitative relations between these forms of chlorocomplexes in HCl media are not exactly known. The whole problem of bivalent lead is complicated because at acetate ion concentrations of the order of 10-l M (the case of the equilibria studied) acetate complexes of the type [PbAc] - exist already (4). On the other hand, we must also keep in mind that these acetate complexes are considered and proved in pure acetate solutions (4), i.e., in the absence of chlorides. Since in our case we have a ZO-fold to 40-fold excess of chloride ions compared to the acetate ions, and because the great complexing capacity of chloride ions is well known, it may be assumed that these acetate complexes will probably not participate. It is impossible, however, to decide on the measure in which the ratios of the complexes [PbClr12and [PbCla] - will exist in the given acidity range of 2 to 4.5 M HCl. Based on these considerations, the conditions for the presence of possible other forms of the components present in the equilibrium solution, and under the condition that in the equilibrium studied no free chloride ions will participate (this condition follows from the constants determined), the equilibrium may be expressed in general by means of the following equations: [PbClo12- e [PbC14]*- + Cl2 (a) (b)
[PbClc]‘-
z+ [PbCla]“-
+ Cl:<-.
A comparison of the mean value of the constant K’ in 1 M HC1 in Table 1 with the value given by the authors of (5) shows clearly, that both constants agree. The increase of values of constants for the range of 2 to 1 M HCl observed in the present work and mainly at concentrations lower than 1 M shows that in the equilibrium mentioned chloride ions start to participate at concentrations lower than 1.5 M HCl. In this case the above scheme of the equilibrium (Eqs. (a) and (b) ) is invalid;
192
V.
DVOHAK,
I.
NeMEC,
AND
J.
ZkKA
chloride complexes of bivalent lkad probably are formed with less than three coordinated chloride ions, and probably even tetravalent lead begins to occur in the form of [PbC&] -, as proved by Heal and May (5) as well as by Wescott (8). SUMMARY Variations of the equilibrium constant K’ = lPb(I1) 1. [Cl,I/[Pb(IV) I have been studied in dependence on the acidity of the solution, based on the polarographic behavior and determination of the concentration of tetravalent lead in the presence of chlorine. A possible scheme of the equilibrium is discussed, based on the invariance of this constant in the range of 2 to 4.5 M HCI and comparison with literature data. REFERENCES 1. 2.
3.
4.
I., PARTON, H. N., AND ROBINSON, R. A., The constitution of the lead halides in aqueous solution. J. Am. Chem. Sot. 77, 5844-5848 (1955). FROMHERZ, H., AND LIH, K. H., Spektroskopische Untersuchung der Dissoziationsverhaltnisse von Blei- und Thallohalogeniden in wassriger LBsung. Z. Phys. Chew A 153, 321-375 (1931). GLASNER, A., AND AVINUR, P., Spectrophotometric methods for the determination of impurities in pure and analytical reagents. II. Some absorption spectra in concentrated chlorines and their applications. Talanta 11, 761-773 (1964). GOBOM, S., The complex formation between lead(H) ions and acetate ions. Acta Broos,
Chem. 5. 6.
7. 8.
A.
&and.
17,
2181-2189
(1963).
H. G., AND MAY, J., A spectrophotometric study of the stability of lead(IV) in hydrochloric acid solutions. J. Am. Chew Sot. 80, 2374-2377 (1958). MERRITT, C., JR., HERSHENSON, H. M., AND ROGERS, L. B., Spectrometric determination of bismuth lead and thallium with hydrochloric acid. Anal. Ckem. 25, 572-577 (1953). SZYCHLINSKI, J., LATOWSKI, T., AND KOREWA, R., A study of the equilibria in the Pb4f - Clsystem. Roczniky Chem. 32, 1013-1023 (1958) (in Polish). WESCOTT, E. W., The equilibrium between chlorine and plumbous and plumbic chlorides in aqueous solution. J. Am. Chem. Sot. 42, 1335-1349 (1920). HEAL,