Aqueous Iron(II1) Chloride Complexes Formation
The Journal of Physical Chemlstty, Vol. 83, No. 13, 1979
1889
Thermodynamics and Kinetics of Aqueous Iron( 111) Chloride Complexes Formation' Ulrich Strahm, Ramesh C. Patel," and Egon Matljevie" Institute of Colloid and Surface Science and Department of Chemistry, Ciarkson College of Technology, Potsdam, New York 13676 (Received October 23, 1978; Revised Manuscript Received March 5, 1979) Publication costs assisted by the Electric Power Research Institute
The formation of ferric chloride complexes in aqueous solutions were studied by spectrophotometric and temperature jump techniques by utilizing both the relaxation times and the relaxation amplitudes. The results were analyzed by a multiparametric curve-fitting procedure. Equilibrium constants, AG, AH, and A S values were determined for the following reactions at 25 "C and at an ionic strength of 2.6 M: Fe,,3+ + Cl,; F= Fe(HzO)ClaF(KOut= 1.1,AG = -0.06, AH = 2.8, A S = 9.4);Fe(HzO)C1,,2f F= FeCli,,,T ( K , = 5.9, AG = -1.06, AH = 0.50, A S = 5.2); Fe,T + Cia; F= FeC1,F (Kl = 7.6, AG = -1.22, AH = 4.1, A S = 18);FeClin,,qP++ Cl,, F= FeC1z,,q+( K 2= 1.8, AG = -0.34, AH = 2.1, AS = 8.1); and FeC12,aq+ + Cl,; F= FeC13,,q(K3= 6.0). The rate constants for the formation of the inner-sphere complexes of the ferric mono- and dichloro ions are also given. A species distribution diagram based on these thermodynamic data is shown for a constant ferric ion concentration (0.086 M Fe(C104)3)and a wide range of chloride ion concentrations. This work represents the first attempt to obtain all thermodynamic and kinetic data for a series of complexes under the same conditions.
Introduction Recently, it has been shown2 that the solid phases formed from solutions of ferric salts vary considerably in chemical and structural composition, particle size, morphology, color, and magnetic properties. The final products of the precipitation processes depend on the concentration of ferric ions, pH, temperature, aging time, and are also strongly affected by the nature of anions. Basic ferric sulfate hydrosols3 form in the presence of sulfate ions, whereas different ferric hydrous oxides are generated with chloride, nitrate, or perchlorate as the counterion. Either P-FeOOH or a-FezO3 precipitates from solutions containing chloride ions while solids consisting only of a-Fe203 separate from nitrate and perchlorate solutions. Particles having the same chemical composition may vary greatly in shape when obtained under somewhat different experimental conditions.2 Obviously, ferric complexes involving different anions, which act as precursor to precipitation of ferric hydrous oxides, must be responsible for the properties of the latter. Thus, a comprehensive analysis of all solute species in ferric sulfate solutions made it possible to define the complexes and to suggest chemical mechanisms of formation of the solid basic ferric sulfate^.^ This work endeavors to resolve the thermodynamics and kinetics of the formation of ferric chloride complexes in acidic media. It is expected that this knowledge would help the understanding of the mechanism of generation of various ferric hydrous oxides in aqueous chloride solutions. Several investigations have dealt with the solute species combining iron(II1) and chloride ions in solutions containing up to 8 M HC1 by means of spectrophotometry, potentiometry, and ion exchangea5 Most of these studies concerned monochloro complexes; only little information is available on polychloro ferric species. The rate of formation of the inner-sphere complex FeClin2+has been studied by stopped-flow,6-8 pressure j ~ m p , ~ Jand O pulse radiation techniques.ll As a rule, it is difficult to separate various equilibria by static measurements. On the other hand, relaxation techniques can provide resolution of different steps involved in ion complexation, particularly if measurements of both the relaxation time and the amplitude are exploited. Although the theoretical treatment of relaxation amplitudes for systems of coupled reactions has been 0022-3654/79/2083- 1689$01.OO/O
de~eloped,l~ a -comprehensive ~~ application of this powerful technique, resulting in the determination of equilibrium constants and enthalpy changes, has not been made yet. In this work a combination of relaxation time, amplitude, and absorbance measurements, carried out with instrumentation of high sensitivity, made it possible to obtain reliable thermodynamic and kinetic data for the formation of several ferric chloro complexes.
Experimental Section i. Materials. Perchloric acid and sodium chloride, Baker Analyzed reagents, were used without further purification. Ferric perchlorate, recrystallized from concentrated perchloric acid, and anhydrous sodium perchlorate (both G. Frederick Smith Chemical Co.) were dissolved in doubly distilled water and passed through 0.02-pm Millipore filters. The ferric perchlorate solutions were standardized by a spectrophotometric determinati~nl~ of ferric ions as the thiocyanate complex in acidic media. ii. Preparation of the Solutions. To prevent the hydrolysis of the iron(II1) ion, we prepared the solutions as follows: 50 mL of 2 M perchloric acid was pipetted into a 100-mL volumetric flask, and variable amounts of a sodium chloride solution and a sodium perchlorate solution were added to give a final ionic strength of 2.6 M (on the assumption of complete dissociation). The flask was then filled with doubly distilled water until a little volume was left for the iron(II1) perchlorate. After the flask was thoroughly shaken, the ferric salt solution was added, shaken again, filled to mark with doubly distilled water, and, finally, thoroughly mixed. iii. Methods. Spectrophotometry. Static optical measurements were performed on a single beam Zeiss PMQ3 spectrophotometer equipped with a Zeiss PMI photometer-indicator with digital readout. This latter device permitted the determination of the absorbance values to three decimal places. Matched, 1-cm quartz (Hellma) cells were used. The cell holder was temperature controlled at 25 f 0.1 OC by means of a Haake FK external circulating water bath. Temperature Jump. Relaxation measurements were performed on a combined stopped-flow temperature jump apparatus described in detail elsewhere.16 For the ex@ 1979 American Chemical Society
1690
The Journal of Physical Chemistry, Vol. 83, No. 13, 1979
periments carried out in this work, the stopped-flow feature provided a convenient means of rapidly introducing the sample solution into the temperature jump cell by activating the thermostated drive syringes. The time interval which is normally required for temperature reequilibration in the cell could be eliminated in this manner. All solutions in the cell as well as drive syringes were kept at 21.0 "C and a temperature jump of 4.0 " C was applied by discharging a 0.04-pF capacitor which had been charged to 14.4 kV. The temperature rise of 4.0 "C was directly established by observing the optical density change of the temperature-sensitive Tris buffer-phenol red indicator system, which agreed closely with earlier determinations.16 Individual amplitude measurements could be reproduced to within -2% for optimum conditions, with errors in the relaxation times varying between 5 and 10%.
Data Treatment and Calculations i. Spectrophotometric Determination of FeC12+. The equilibrium constant K1 and the extinction coefficient EF&~Z+ for the monochloro iron(II1) complex were calculated at wavelengths 395-420 nm by using absorbance of the solutions in the absence of chloride ions (Ao),and in the presence of chloride ions (A): A0 = €'~~S+[Fe~+]t,t
(1)
where E ' ~ ~ sis+ a proportionality factor, the relationship of which to the extinction coefficient eF+: is discussed later.
Applying the law of mass conservation and the principle of mass action yields
1 (3) (€FeC12+ - €'~~s+)K~[Fe~+lt,t - K i ( A - Ao) Using a curve-fitting program, one can obtain the equilibrium constant, K1, and the difference in the extinction coefficients (€FeClZ+ - dFes+), from which qeC12+ can be determined, by using the known value for dFes+ (eq l). By considering only the first hydrolysis step in the absence of chloride ions
U. Strahm, R. C. Patel, and E. MatljevlB
chloride concentrations ([Fe3+]>> [Cl-1) as described in the previous section. The remaining three parameters, i.e., the extinction coefficients for FeC12+and FeC12+(for the wavelengths of interest) and P2, can be determined by using a multiparameter curve-fitting program1' and an iteration procedure to calculate the free ferric and chloride ion concentrations. iii. Relaxation Times and Relaxation Amplitudes from Temperature J u m p Measurements. Time-voltage data obtained from temperature-jump experiments were all considered to follow single exponential curves:
AV = AV,(l
- e-t/')
(8) in which AV is the voltage change at time t , AV, is the total voltage change, and T is the relaxation time. To avoid the uncertainties in using the initial and final values, we treated the data by the Guggenheim method,ls which requires values at constant time intervals. A linear regression analysis was applied to the log voltage data from which, on multiplication with the appropriate time and voltage scale factors, the relaxation time and the amplitude values were obtained. These parameters can then be used to establish the kinetic and thermodynamic quantities for the observed reaction step. For the special case of interest in this work, the formation of the monochloro iron(II1) complex can be presented by the following system of equilibria: slow
fast
Feaq3+ + Cia; e 'Wt Fe(H20)Cla,2+e 'in FeClin,aqP+ (9) in which Fe(H20)C1,2+ and FeClin,aq2+represent the outer-sphere and the inner-sphere monochloro complex, respectively. For the relaxation time of the slow process13
where KO, is the equilibrium constant for the outer-sphere complex and kin and k-in are forward and backward rate constants for the slow reaction (eq 9). Square brackets designate the equilibrium molar concentrations of ferric and chloride ions. In order to reduce the number of parameters, we can write eq 10 in the form
+
A0 = cFe3+[Fe3+] €F~oH~+KF~oH~+[OH-] [Fe3+] (4) In highly acidic media (1 M HClOJ, [FeOH2+]becomes very small and, consequently, one can write and
ii. Spectrophotometric Determination of FeC12+. T o obtain the equilibrium constant (Pz) of the dichloro iron(II1) complex, we performed spectrophotometric determinations a t wavelengths of 430-460 nm. Under these conditions, the contribution of the free ferric ions to the total absorbance is negligible.
Thus, eq 11 permits the calculation of all kinetic and thermodynamic parameters from the measured relaxation time. For the fast decoupled formation of the inner-sphere dichloro complex fast
FeClin,aq2+ + Cia; + FeC12,aq+ 2'
or
To reduce the number of parameters we obtained the first equilibrium constant (PI = K,) by using systems with low
which appears at higher chloride concentrations, the ) be following relationship for the relaxation time ( T ~ can derived: 1 / =~k2([FeCli,2+], ~ + [Cl-1,)
+ k2
(15)
Aqueous Iron(II1) Chloride Complexes Formation 0.08
I
The Journal of Physical Chemistry, Vo/. 83, No. 13, 1979
TABLE 11: Values of
I
CCl~lTo,~00I0M p z 2 6 M INoCIOI) HCIO. 1 0 M 25OC
E’Fe3+for
1691
Different Wavelengths
h,nm
395
400
405
410
415
€IFe$+
0.79
0.72
0.69
0.60
0.46
I
0.06
TABLE 111: The Overall Equilibrium Constant p z and the Extinction Coefficients EFeC12+and f F e C p + in 1M HClO, and at Ionic Strength 2.6 M this work ref 1 9
a
7P
Y
\
0.04
c
2
h,
m -1
0.02
0‘
nm
430 440 450 460 I
0.05
1
I
1
0.15
0.10
0.20
0.25
C F e 3 + 1 ~#~ ~ , Figure 1. Plot of total chloride ion concentration divided by the difference in measured absorbances between chloride containing ( A ) and chloride free (A,) solutions vs. total ferric ion concentration. Data were obtained at different wavelengths from solutlons 0.010M in chloride ion, 1.0 M in HCIO,, of ionic strength 2.6 M (adjusted with NaCIO,), and at 25 OC.
TABLE I: The Overall Equilibrium Constant for FeCI2+ and the Extinction Coefficients Difference for Ferric Chloride Solutions 1 M in HClO, and of Ionic Strength 2.6 M EF,eC1z+- €FeCl*+
[Cl-Itot, M nm 10-2 415
E Fe3+
410 405 400 395
41 60 87 121 166
2 X
420 415 410 405 400
29 41 59 84 123
3X
420 415 410 405
27 39 57 83
(ref 1 9 ) 61 128 30
K, 7.8 7.3 7.2 7.8 7.4 7.4 7.9 7.6 7.5 7.1 8.5 8.6 8.2 7.8 av 7.6
~t 0.4
As will be shown later (eq 21 and 24), the relaxation amplitude can be utilized to calculate the reaction enthalpy (A In K ) . For this purpose the total voltage change, AV,, in the time-voltage curve has to be related to the change in absorbance, AA, as follows: AA=-
1 In 2.3031
(
At)
l+-
(16)
where 1 is the path length and Vothe voltage before applying the T-jump.
Results i. Spectrophotometric Data. FeC12+. The equilibrium constant for the monochloro ferric system, K1, was determined from absorbances at 395,400,405,410,415, and 420 nm in 1 M HC104 at an ionic strength of 2.6 M (adjusted with NaC104). Three series of solutions were investigated, each containing a constant total concentration of chloride ions ([Cl-I,, = 1 X 2X and 3 x M) and varying total concentrations of ferric ions ([Fe3+Ibt = 0.05-0.25 M). Figure 1shows the data for [Cl-1, = 0.010 M, in which the experimental points are given by circles and the calculated best fits (eq 3) are shown by the solid
4,
EFeCl.+
11.4 11.4 11.3 11.4
21 20 14 8.1
EFeCIZ+ fFeCl,+ €FeCl2+
18 8.5 4.0 1.9
30
12.6 5.8 2.3
lines. Table I gives the extinction coefficients as measured in this work and the corresponding literature values.lg Also included in the table are the values of the equilibrium constant, K1, calculated from eq 3, by using data as illustrated in Figure l. The somewhat higher values for K1 for systems containing 3 X M C1- may be due to the presence of small amounts of the dichloro complex. In order to obtain cFeCp+ it was necessary to determine the values of t’Fe3+, which were calculated with eq 5 and are shown in Table 11. Obviously, the values of t ’ ~ ~ 3are + small enough to allow the approximation tFeC12+ - t ’ ~ pN CFeCl2**
FeC12+. The equilibrium constant (pz)for the dichloro ferric complex, FeC12+,was obtained from absorbance measurements a t 430, 440, 450, and 460 nm of solutions containing a constant concentration of [Fe3+Ibt= 0.086 M and varying concentrations of C1- (0.030-0.5 M). All solutions were 1M in perchloric acid with a constant ionic strength of p = 2.6 M (NaC104). Table I11 gives the three parameters (tFeClZ+, C F ~ t,C and ~ p2) as calculated from eq 7 with = K1 from Table f. The values of cFeCp+ shown in Tables I and I11 obtained by different measurements are consistent. The excellent agreement of the p2 constant at four different wavelengths indicates the reliability of the t values determined in this work. ii. Relaxation Times. Relaxation measurements in solutions of variable chloride ion concentrations (0.03-2.0 M) containing 0.1 or 0.086 M Fe(C104)3in 1M HC104, at an ionic strength of 2.6 M (NaC104),showed three separate effects: (1)A very fast relaxation process (relaxation time T = 10-13 p s ) occurred in all solutions. Neither T nor the amplitude of the relaxation changed measurably with the chloride ion concentration. The signal consists of two components, a chemical and a physical one. The chemical effect can be attributed to the extremely rapid formation of an outer-sphere complex, Fe(H20)C12+,whereas the physical effect is due to the expansion of the solution and to changes in the extinction coefficients with temperature. These two effects are difficult to resolve with the present instrument. (2) A relaxation time of 100-200 ms was observed in some solutions depending upon the total chloride ion concentration. This relaxation time is within the range necessary for the formation of the monochloro inner-sphere complex (FeClin2+)as found by other investigators.6-11 (3) A t high chloride ion concentrations an additional relaxation effect appeared at a constant time of -540 ks. While T shows no significant dependence, the amplitude changes greatly with increasing chloride ion concentration. This relaxation process indicates the formation of more than one ferric-chloro complex. The values of 7 and the corresponding reaction rates are of the expected magnitude
1092
U. Strahm, R. C. Patel, and E. MatijeviE
The Journal of Physical Chemistry, Vol. 83, No. 73, 7979
following system of reaction steps on the basis of the relaxation time measurements:
(1)
Feaq3++ Cia;
=Fe(HzO)C1,,2+ very fast
(18)
slow
Fe(H20)Cla,2+eFeClin,aqP+ (111) (IV)
0.05
0.15
0.10
c - F ~ ~ * ICCI-I, ,+ (rnolell) Figure 2. Plot of reciprocal relaxation time vs. the sum of the equilibrium concentrations of ferric and chloride ions for solutions 1.O M in HCIO,, of ionic strength 2.6 M (adjusted with NaCIO,), and at 25 OC.
for disubstituted complexes. Because the three observed relaxation processes were on significantly different time scales, the measurements were treated as single relaxation steps. The relaxation time data obtained in the range of 100-200 ms for a series of solutions containing different total concentrations of ferric and chloride ions are plotted in Figure 2 (circles) in which the solid line represents the best fit computed according to eq 11. The outer-sphere, the inner-sphere, and the overall formation constants which gave this fit were KO,, = 1.1mol-l L, Kin= 5.9, and K1 = 7.6 mol-' L, respectively. The latter value is in excellent agreement with the independently determined K , from the static spectrophotometric data (Table I). The rate constant for the reaction Fe(H20)Cla,2+z% k .in
AA2 =
A In Ki,
rout1
(21)
1> -1
r 2 = ( [FeClin2+] 1 +-[Cl-] 1 + [FeC12+]
(22)
The total change in absorbance is AA,, = CA&Ati i
i = out, in, 2
(23)
in which expression A& is the overall concentration change for the step i. A& is related to the equilibrium concentrations of the species and to the change of the equilibrium constants by the system of equations shown in eq 24 which can be derived using the law of mass action and the principle of mass conservation for the chloride and the ferric ion. The terms A& can be expressed explicitly as a function of equilibrium constants and equilibrium ferric and chloride ion concentrations from eq 24, but the results contain numerous terms. In practice, it is convenient to use a more general numerical approach. Initially, the equilibrium concentrations in eq 24 were calculated using a set of independently determined equilibrium constants 1
[Fe(H,0)C12+] [Cl-]
[ Fe(H,O)Cl* '3
(20)
In K2
(cFeCI2+ -
(17)
1
In Kout
A In K,
+ FeC13,,q
where FeClh,,q2+
as calculated from eq 10-13 are kin = 21.8 mol-' L s-' and k-in = 3.7 S-'. As mentioned above, the relaxation time for the formation of ferric dichloride (eq 14) was constant (- 540 ps) under the chosen experimental conditions. A value for the backward rate constant k2(eq 15) of 185 was calculated from the average of seven different sets of experiments. From kZand the corresponding equilibrium constant K 2 = P2/K1 = 1.75 mol-' L-') we could obtain the value for the forward rate constant as k 2 = 324 mol-' L s-l. iii. Relaxation Amplitudes. In evaluation of the relaxation amplitudes, it is important to recognize that the measured absorbance change over a certain time scale may be affected by more than one reaction step. For example, a slow reaction preceded by a fast one could show a significantly different amplitude than the case in which the slow step occurred alone. Accepting the assumptions of Wendt and Strehlow? that outer-sphere complexes formed with more than one C1bound to ferric ion can be neglected, we propose the
A
FeC12,aq++ Cl,;
(19)
The formation of the outer-sphere complex is very fast and, therefore, it is not disturbed by the other reaction steps. The fast formation of the dichloro complex is decoupled from the very fast first step through the slow second step. The amplitude may only be influenced by step IV, if that formation rate constant is fast. The amplitude observed in the 100-200-ms time range is caused primarily by step I1 but it is also affected by concentration changes in steps I and 111, because these processes are fast as compared with the second step. Contributions from the step IV, if it occurred a t 200 ms or faster, would have to be considered. FeC13 may have a high extinction c~efficient;'~ thus, very small concentration changes could give rise to significant absorbance changes. The fourth step was not observed separately, and no evidence of a contribution to the amplitudes of the fast and/or slow processes was found. Relatively easy to describe are the absorbance changes12-14for the fast process I11 (if step IV can be neglected) and for the sum of the changes in absorbance taken over the entire time range. The fast process I11 can be treated as a single step
0.25
0.20
fast
FeClin,,qS++ Cia; e FeC12,aq+
rh-'
-
1 [ FeClhz+ I
r2-'
lr-
1
Aqueous Iron(II1) Chloride Complexes Formation
The Journal of Physical Chemistry, Vol. 83, No. 13, 1979
1693
TABLE IV: Species and Their Overall Formation Constants (25 'C, 1 M HClO, , p = 2.6 M ) species 1% P origin Fe(OH)z+ - 2.92 ref 3 - 5.7 ref 3 Fe( OH), - 3.22 ref 3 Fez(OH),, 0.045 eq 11-13 Fe( H,O)Clz+ FeClh2 0.81 eq 11-13 FeCl, 1.06 eq 7 +
and a species distribution program.20 From the Ati values obtained by multiplying the diagonalized inverse of the concentration matrix with the ( A In Ki) matrix, the total absorbance change could be computed according to eq 23 by using a curve-fitting program" in which the species distribution and matrix calculations were contained as subroutines. This procedure allows us to use as variables the equilibrium constants and the reaction enthalpies (A In Ki). In order to determine the Ati values needed in the amplitude treatment, we performed static absorbance measurements for the same concentration ranges of ferric and chloride ions by means of the optical system of the temperature jump apparatus a t 25 "C and 420 nm. The additional measurements of absorbancies were dictated by the wider band width used in the relaxation experiments to provide higher sensitivity in order to detect smaller concentration changes. The extinction coefficients were then calculated with the aid of very accurate equilibrium constants obtained from absorbance measurements with the Zeiss spectrophotometer at several wavelengths. The values so obtained are C F ~ S += 0.6, EFeC12+ = 12.1, and CFeC12+ = 59.5. The individual extinction coefficients for the outer- and inner-sphere complexes of iron(II1) monochloride cannot be obtained separately. The extinction coefficient of the is assumed to be practically outer-sphere complex cF~(H~O)C~Z+ identical with cFe3+, while the extinction coefficient of the inner-sphere complex can then be calculated from eFe(H o)c$+and the overall extinction coefficient for the iron$II) monochloride
4
Atin
=
AC2
=
CFe(HzO)C12+ - EFeS+ ss
CFeClzf
10
4
20
3 (log)
Figure 3. Plot of the absorbance change for the fast reaction (eq 19) (circles) and of the total absorbance change (squares) against the log of the total concentration of chloride ions for solutions either 0.086 or 0.10 M in Fe(C104)3,1.0 M in HCIO, of ionic strength p = 2.6 M (adjusted with NaCIO,), and at 25 OC.
and by taking into account the complexes and their independently determined overall formation constants (P), as given in Table IV. The importance of considering the first four complexes listed in this table lies only in their influence on the equilibrium concentrations of the species directly involved in the fast process I11 (eq 22). In addition, it is necessary to consider the possible effect of the complexes FeC13 and FeC14-. The computations have shown that the inclusion of the FeC14- ion could not improve the agreement between the calculated and observed data. On the other hand, the neutral species FeC13 is essential to getting a good fit for the amplitude data (Figure 3, circles). The computations yielded an overall equilibrium constant for the ferric trichloride complex, FeC13, log P3 = 1.8, and a value for the enthalpy change for the formation of FeCl,+, AH2 = 2.1 kcal mol-l. Finally, the total absorbance change was evaluated in terms of eq 23 and 24. In this case A In K2 was fixed in addition to the equilibrium constants given in Table IV and A In Kout,A In Kin,and log P3 were taken as variables. Figure 3 (squares) shows the experimental points for the total absorbance change and the curve that best fits the data. The following values were obtained: A In Kout= 0.061, A In Kin = 0.01, and log P3 = 1.9. The agreement in the log P3 values as calculated from the fast reaction step above and from the total absorbance change is excellent. The reaction enthalpies AHi,calculated from the van't Hoff equation, A In Ki = (AHi/RP)AT, for a 4-deg temperature jump are listed in Table V (a, b, c), together with the reaction entropies ASi, and the reaction free
0
E F ~ C ~~ ZC +F ~ ( H ~ O ) C = ~ Z 13.5 +
=
2
CCI-I/tFe3'
In the interpretation of the temperature jump measurements (eq 21 and 23) we used the following values: AEout
1
- eFeClin2+ = 45.4
The amplitude of the fast process I11 was treated by using a curve fitting and a species distribution program,
TABLE V : Thermodynamic Data for Single Reaction Steps and for Overall Formation Processes of Ferric Chloride Complexes
(a) Feaq3++ Claq-
Kout
Fe(H,0)CIaq2+
(b) Fe(H,O)ClW2'FT FeClh,aqzt (c) Feaq3*+ C1,-
~t
FeClh,aq2+
K,
(d) Feaq3+t Claq- -== FeClaq (e) FeCli,-,aq*t+
(f)Feag
-
K
Cia< 3 FeCl, ,aq+
+ 2Claq-
0,
kcal mol-'
A S , cal mol-' deg-'
1.1
-0.06
2.8
9.4
5.9
- 1.06
0.50
5.2
6.5
-1.12
3.2
15
7.6
- 1.22
4.1
18
1.8
-0.34
2.1
11.4
- 1.46
5.3
Kora
Kin
4in
kcal mol-'
A G,
process
FeCl, aqt
AH,
8.1 23
1094
The Journal of Physical Chemistry, Vol. 83, No. 73, 1979
U. Strahm, R. C. Patel, and E. Matljevie
TABLE VI: A Comparison of the Equilibrium Constants for Ferric-Chloro Complexes T," C P,M media methoda log Ki
ref
FeCl 25 25 25 25 25 25
2.6 2.6 2.6 2.6 2.6 2.6
(c) HC10, (I) HCIO, (I) HCIO, 1 M H', NaC10, 1 M H', NaC10, (I) HCIO,
SP cal SP sp, kin pulse rad cal
25 25 25
2.6 2.6 2.6
0.027 M H' 1 M H', C10,1 M HI, NaC10,
Fe( H,0)C12 kin pulse rad sp, kin
26.7 25 20 25
1 1.2 2 2.6
(I) H, NaClO, (c) HCIO, ( c ) HClO, 1 M H', NaC10,
FeC1, SP SP SP, pot. SP
26.7 20 25
1 2 2.6
(I) H, NaC10, (c) HClO, 1 M HI, NaCIO,
FeCI, SP sp, red sp, kin
0.62' 0.65' 0.77b 0.88 0.91' 0.92'
21 22 23 this work 11 8
0.3 * 0.18 0.50' 0.045
9 11 this work
0.11 0.18 0.30 0.17
19 24 25 this work
1.40 -0.06 t 0.78
19 25 this work
+
-
a sp, spectrophotometry; cal, calorimetry; kin, kinetic; pulse rad, pulse radiolysis; pot., potentiometry ; I, ionic strength; c, concentration. Interpolated,
energies AGi, as obtained from the equilibrium constants.
Discussion The purpose of this work was to describe as fully as possible the numerous equilibria and the species present in an acidic Fe3+-C1- system. A species distribution (Figure 4) was computed20for a constant iron(II1) concentration (8.6 X M) and a wide range of chloride ion concentrations. The equilibrium constants used for the Fe3+-C1- complexes were obtained from static spectrophotometric and from relaxation time data as described in previous sections. The hydrolysis of the Fe3+ ions was considered, by using the equilibrium KF~(oH)~+, and KFe2(0H)24+,as determined constants KF~OHZ+, earlier,4 but the hydrolyzed ions are present only in negligible quantities (all
