- Email: [email protected]
Transfer rate of some tervalent cations in the biphasic system HClO4, water-dinonylnaphthalenesulfonic acid, toluene—I
J inor~ nucl. ('hem., 197h, Vol 3N pp. 1687-1693. Pergamon Press.
Printed in Great Britain
TRANSFER RATE OF SOME TERVALENT CATIONS IN THE BIPHASIC SYSTEM HC104, WATERDINONYLNAPHTHALENESULFONIC ACID, TOLUENE--I TRANSFER
RATE OF IRON(III) DETERMINED CHEMICAL
BY
REACTIONS
P. R. DANESI, K CHIARIZ1A and A. SALTELL[ Laboratorio Chimica Fisica, C.S.N. Casaccia, C.N.E.N., Rome, Italy (First received 26 June 1975: in ret:isedform 15 October 1975) Abstract--The transfer rate of iron(Ill) ions between aqueous perchlor~c acid solutions and toluene solutions of dinonylnaphthalenesulfonic acid has been studied by using a constant interracial area cell. The initial forward and reverse transfer rates have been evaluated as function of the interfacial area, volume and chemical composition of the system. The results have been explained in terms of interfacial reactions which rate constants have been evaluated. An equation capable to describe the dependence of the transfer rate from both the composition of the system and the time has been derived. INTRODUCTION
THE mass transfer rate of metal cations in liquid-liquid solvent extraction systems is in the most general case a function of both the kinetics of the chemical reactions taking place in the system and the rates of diffusion of the involved species in the two liquid phases. The simultaneous presence of both chemical reactions and diffusion processes in liquid-liquid mass transfer has been recently reviewed[l] and several mathematical diffusion models with heterogeneous chemical reactions have been elaborated[l,2]. However several cases seem to exist where the rate of mass transfer of cations in liquid-liquid solvent extraction systems is mainly controlled by interfacial processes [3-7]. In this instance the mathematical model of the system greatly simplifies and providing both the interfacial and thermodynamic properties of the system are well known a kinetic scheme can also be derived. In the present paper the mass transfer of Fe(lli) between aqueous perchloric acid solutions and toluene solutions of the liquid cation exchanger dinonylnaphthalenesulfonic acid, HD, has been studied. It was in fact presumed that with such a system, previously well characterized by us in its interfacial[8] and thermodynamic [8, 9] properties, it would have been possible to receive some further indication on the role of interfacial chemical reactions on the overall mass transfer rate. EXPERIMENTAL Materials Iron powder 97% pure, HCIO4 70% RPE-ACS and toluene RPE-ACS were Carlo Erba products. An Fe(lII) perchlorate solution in HCI04 has been prepared by dissolving the iron powder in hot perchloric acid and boiling for several hours in order to eliminate Fe(II). To test that the solution contained no Fe(lI) the reaction with orthophenantroline has been used. The Fe(III) content has been determined by titration with EDTA using sulfosalicilic acid as indicator. The hydrogen ion concentration has been determined by a standard procedure. A stock solution of dinonylnaphthalenesulfonic acid, HD, supplied by Columbia Organic Chemicals Co. Inc., and purified as previously reported[10], was prepared by dissolving HD in toluene and by determining its concentration by potentiometric titration[4]. A stock solution of Fe(Ill) in the organic phase has been prepared by shaking equal volumes of the aqueous Fe(III) perchlorate stock solution with a concentrated solution of HD in toluene. Its concentration has been evaluated by the distribution coefficient. All the solutions used in the experiments have been prepared by mixing the stock solutions either with bidistilled water or HCIO4
standard solutions or toluene. All other reagents were Carlo Erba analytical grade purity products. The calomel saturated NaC1 reference electrode and the agar-agar NaCI bridge have been prepared according to standard polarographic procedures. Technique A constant interfacial area cell (Fig. 1) has been employed in most of the experiments. This cell permits both forward and back kinetic experiments of the Fe(II1) extraction, keeping the interface quiescent, with a continuous monitoring of the concentration of Fe(lII) in the aqueous phase as function of time. To the purpose the cell contains a platinum rotating electrode Metrohom SM/5 and a calomel NaC[ saturated electrode connected via an agar-agar NaCI bridge. These two electrodes are used in an amperometric system to detect iron. To the purpose a Metrohom Potentiograph model E 436 was used. Preliminary experiments have shown that applying a polarization voltage equal to zero a very good linear proportionality exists between the current and the Fe(lll) concentration in the aqueous phase. Moreover the current variations caused by Fe(IIl) concentrations variations are practically instantaneous. Therefore it was possible to folliow the concentration variations of Fe(III) with time by simply recording the current vs time curves and by properly calibrating the system each time using Fe(III) solutions of known concentration. The forward extraction kinetics of Fe(IIl) were therefore performed in the following way: a known volume (generally 40 cm ') of HCIO~ of known concentration was placed in the cell and after checking the instrumental zero the circuit was closed starting at the same time the stirring of the cell. A constant residual current was generally established within 2 hr. Afterwards progressive amounts of the stock solution of Fe(III) were added obtaining in this way a current vs concentration calibration curve. A volume of HD solution in toluene equal to that of the aqueous phase was then added by stopping, during the time of the addition of the organic phase, both the rotation of the platinum electrode and the stirring. The starting of the kinetics then coincided with the new starting of the stirring and of the rotation of the Pt electrode. The current decrease with time was then directly correlated to the concentration vs time curve since the recorder chart speed was known. The organic phase was each time preequilibrated with an aqueous HCIO4 solution having the same concentration as that in the cell. Figure 2 reports a typical forward extraction kinetics experiment. The backward extraction kinetics (organic Fe(IIl)~aqueous Fe(IIl)) were performed as follows: after addition in the cell of the aqueous perchloric acid and attainement of the residual cmrent an organic phase consisting of a toluene solution of HD and Fe(II1) of known concentration was added to the aqueous phase solution following the same procedure as for the forward kinetics. After the back extraction had proceeded for some time, the extraction of iron was stopped by removing the organic phase by careful suction followed by repeated washing with pure toluene to remove
1687
P. R. DANESI and R. CHIARIZIA
1688
Stirrermotorand I gearOox II
----I ~"
Toth..... ......ording
II
|lOutlet
~I | - I I
J Agar-Agar ~O~s.S.
J
1"1I J [ ]
"~
|
~
TOlhe recording amp..... for (+)
(LI( "m III I
~~ -1 -- tiJl '~1] AqueousJ
-Th..... IOI ~,~ inlet
Glass fritt Fig. 1. Apparatus used to study the transfer rate of Fe(III) by continuous monitoring of the Fe 3+ aqueous concentration through amperometry. Interface radius=2.45cm; interface geometrical area=18.9cm2; stirrer blades = 20 x 5 x 2 ram; distance between stirrer blades = 20 mm; volume of aqueous and organic phases ~40 cm3; Pt electrode rotation speed = 750 rpm; stirrer rotation = 1I0 rpm,
I10 I00 90 80 70
~ 6o i_ ~ 50 C3
L 40
20 I0 O
OI 234567891011P2 Time, min
Fig. 2. Typical forward kinetics; [Fe3÷] = 6.87 x 10-" M, C'm, = 0.046 M, [HCIO4] = 0.451 M; Volume of aqueous and organic phases = 43 cm3; ® residual current; ®, ®, ®, ®, ®, ® currents corresponding to progressive volume additions of the aqueous Fe ÷3 stock solution; ® stopping of the stirrer and Pt rotating electrode and addition of the organic phase; ® starting of the forward extraction kinetics; # current v s time plot. traces of HD. At this point the calibration curve, current vs concentration, was started by adding to the aqueous phase known amounts of the Fe(III) stock solution. Figure 3 reports a typical back extraction kinetics experiment. The temperature of the cell was kept at 25.00-0.05°C by a circulating ultrathermostat. The stirring speed was kept constant in all experiments at 110 rpm. Due to the slow extraction kinetics the initial transfer rates were obtained by the slope of the straight line which was well representing the first five minutes of the kinetic experiments. The forward kinetic experiments performed to obtain information on the influence on the rate of the interfacial area have been obtained by using a cell having the same design as that of Fig. 1 but a much larger diameter. The interracial area was then reduced by means of teflon interface limiting rings.
solutions and HD toluene solutions for the following experimental conditions: temperature = 25°C, [Fe+3] = 9.85 x 10-4 M, [HC10,] =0.53-2.00M, C,D (total organic concentration of HD) = 0.018-0.23 M, have been obtained by shaking for 60rain (enough to attain equilibrium) 10cm3 of each phase in a mechanical shaker. After centrifugation and phase separation the equilibrium concentration of Fe 3+, reduced to Fe 2÷ with hydroxylamine chloridrate, was determined spectrophotometrically at 510nm by adding ortophenantroline and 2M sodium acetate buffer and using a Beckmann DK2A spectrophotometer. Figure 4 shows the log D vs log (~nDand log D vs log [HC104] plots. Both are well fitted by straight lines of slopes +1 and -3 respectively, thus indicating that the extraction of iron(III) is well represented by the following stoichiometry
Equilibrium distribution studies o.f Fe(III) The distribution ratio, D of Fe(IlI) between aqueous HCIO4
Fe3++ ( H D ) ~ F e H , , - 3 D m + 3H +
(1)
Transfer rate of iron(III) in the biphasic system HCIO~
1689
2O !
90
8Ovo
~
~®
60
-~ 50
J<~7b
2()
i
o
2
4
,
i
6
i
i
8
i
i
,
i
i
ko ~2 i~
Time. rain Fig. 3. Typical reverse kinetics; [Fe] = 1.0x 10 ~M, CH~= 0.025 M, IHCI0~] ~ l.ll M: volume of aqueous and organic phases = 40 cm3; @ residual current; ~ stopping of the stirrer and Pt rotating electrode and addition of the organic phase; ~, starting of the reverse extraction kinetics; ® current vs, time plot: ® removal of the organic phase. ~, @. ®, ®, Q, ~ currents corresponding to progressive volume additions of the aqueous Fe ~4 stock solut]o . '
n
\
\ ~
°~°
!O
¢ eo '
./
o"
%
/ai
\ 2
-I
@
4
IogCHo, M
/
-I
J-I i, I
O log [ H C I 0 4 ] , M
Fig. 4. (a) log D (distribution ratio of Fe(lI[)) vs log CHI>plot; (b) log D vs log [HCIO,] plot: (a) [HCI(L} : 1.02 M, {Fe '+] = 9.85 x 10 ~ M, T = 25°C; (b) (~,i~ = 0.046 M, [Fe 3*] = 9.85 x 10 4 M, T = 25°C; The slopes of the straight lines are +1 for (a) and -3 for (bl. where m ~ 7 is the degree of polymerization of HD [8] and the bar indicates organic species. These results well agree with those of references 11-13. RESULTS AND DISCUSSION
Dependence of the initial forward mass transfer rate on the volume of the phases, and interracial area In order to examine the dependence of the mass transfer kinetics from bulk chemical reactions, initial mass transfer rates were obtained as function of the volume of the phases and of the interfacial area. The forward mass transfer rates at time zero divided by the initial iron(III) concentration, V' = V/[Fe ~+] (with V = -d[Fe3+]/dt) are plotted in Fig. 5 as function of the volume of the aqueous phase (an equal volume of organic
phase was used) at constant interfacial area. A straight line passing through the origin is obtained thus indicating that reactions occurring in the bulk of the phases are not rate determining. Figure 6 shows a plot of V' vs A~ (geometrically calculated interfacial area) at constant volume. Also in this case a straight line is obtained. However the line does not go through the origin indicating that the real interfacial area and the one calculated through the cell diameter are not coinciding. This fact can be explained by considering that the platinum electrode, rotating at 750 rpm, causes a vortex of fixed area which contributes with an additional constant term to the geometrically calculated interfaciat area. However the data of Fig. 6 indicate that the rate determining processes are taking place in proximity of the
1690
P . R . DANESI and R. CHiARIZ~A
1.5
15
NO--L0
0.5
0,5
0
0 0
I
3 I/V,
20
cc-a I02
i
i
40
60
A I , crn 2
Fig. 5. V' vs 1IV (volume) plot; V is the volume of the aqueous and organic phases; [Fe ~*]=6.87x 10-~M, C ~ = 0 . 0 2 0 M , [HCI04] = 0.451 M; interface area = 18.9 cm 2.
Fig. 6. V' vs A~ (geometrical interfacial area) plot; [Fe 3+] = 7.30 x 10 4 M, CHD = 0.057 M, [HCI04] = 0.451 M; aqueous and organic volumes = 108 cm ~.
interface (either interracial reactions or diffusion processes in the interracial liquid films).
yF,tao,)3; yF,tao,)3 has been estimated through the Davies equation[16] assuming that the ionic strength is only determined by the perchloric acid. In order to explain the experimental results of Fig. 7 a model based on rate determining interfacial chemical reactions has been elaborated. These reactions can be formally treated as adsorption-desorption processes as in Ref. [3]. On the basis of the known interracial properties of the HDtoluene-HC104 system[8] it has been assumed that an interracial complex between iron ions and the interracial monolayer of HD molecules is formed. This interracial complex can in turn react either with the aqueous or the organic phase. The nature of this interracial complex must be of course the same independently of which phase the iron is coming from. Further the assumption is made that the steady state is obtained at the interface instantaneously (t = 0) i.e. when no net mass transfer has yet taken
Forward initial mass transfer rates Figure 7 a, b, c (full points) shows the experimental forward initial mass transfer rates (t = 0) divided by the Fe(III) activity, V* = V/aFe3+, plotted as function of the following variables: ale3+ (6',D, [HC104] constant, Fig. 7a), an+ ([Fe3+], CHI~constant, Fig. 7b), t~im ([Fe3+], [HC104] constant, Fig. 7c). The mean activity of the H ÷ in the aqueous perchloric acid (all+) has been calculated as the product of the mean molar activity coefficient, yuc~o,, and the molarity of the solution, yHao, has been obtained from literature data [14, 15]. The mean molar activity coefficient of Fe(III) in the aqueous phase (aFP) has been computed as the product of the iron molar concentration and the mean molar activity coefficient of Fe(III) perchlorate,
O•O \o
4.0 ~_O3.0
(a) (b)
~O 5,(3
%\
.;
~
2.0
4.° -- • 0 0
• •
o.O
°o'o~
o%---
3.G
1.0
o!5 GH+ ,:o o/°"O/°--
40 N 3,0 O
2.0
/
/
2'0
• 4[0
6'.0
810
aFe+ •,[0 4 (c)
1.0 0
0.05
O. I0
0 15
CHD, M
Fig. 7. V* vs concentrations (or activity) plots; (a) V* vs aFe 3+ ( C B D = 0.046 M, [HC10,] = 0.451 M), (b) V* vs a,÷(CHr, = 0.046 M, [Fe 3÷] = 6.87 x 10 "M), (c) V* vs C'HD([HCIO,] = 0,451 M, [Fe 3+] = 6.87 x 10 4 M); interface area= 18.9cm 2, aqueous and organic volumes=43cm3; full points=experimental values; continuous lines calculated through eqn (10) with K, = 4.80 x 10-2 and KR = 1.16 x 10-1.
1691
Transfer rate of iron(Ill) in the biphasic system HCIO, place. The situation holding at t = 0 is schematically reported in Fig. 8a, when Fe(III) is initially present only in the aqueous phase. In order to derive an equation capable to interpret the experimental initial mass transfer rates the following reactions have been assumed: 1. Formation of an interfacial complex between iron ions and the HD molecules which are absorbed at the interface Fe '' + 2(HD),
)(FeD2+),+2H *
(2)
where the subscript i indicates interfacial species. The rate of this process will be
determined [8] that in our experimental conditions [HD], is constant. 2. Dissociation of the interracial complex, (FeD,.')~, to form again iron ions in the aqueous phase and interracial HD (FeD2~)~ +2H +
)Fe '' + 2(HDL
The rate of this process, assuming ideal behaviour of inteffacial species, will be proportional to the concentration of the interfacial complex and to the square of the aqueous hydrogen ion activity
Vd,, = K~[FeD~'I~a~. V, = K ~av.,.[HD]f = K,aw,*
V:
Interfacial complex
org.
Fe ~
(FeD.+)~ +(HD),., ~ F e H . , ,D,,,+2(HD), + H .
V.L. = K~[FeD.~-],[(HD),,,] = K~[FeD:+],d,,.
Interface
Vo (b) Interracial
aq
3. Reaction of the interracial complex with the HD molecules present in the organic phase to form the final product of the overall mass transfer reaction. Since HD is very strongly polymerized in our experimental conditions [8] this reaction can be written as follows (6)
The rate of this process, assuming that also organic species behave ideally, will be proportional to the concentration of the interracial complex and of the polymerized HD
_
aq
(5)
(3)
where K',[HDI~: = K,, since it has been experimentally
(o)
(4)
complex
where K:= K'/m since C.o=[HD]+ m[(H-I)).,I= m [(HD),.]. In order to express the overall rate of mass transfer of iron into the organic phase, Vj., as function of the experimentally studied variables ave>, a . , C.D and the specific rates K,, K.., K3, the concentration of the interfacial complex has to be derived by using the: flux equation at the interface
org
V, -V,l, - V,~,,= 0
Fe(lll)
(7)
(8)
It follows then
Interface
KIOFe3" [FeD:~]~ = K~a~ + K2C.,
(9)
By substituting eqn (9) in eqn (7) we get (c)
I[/°nterfacial
-- ~
aq.
org.
Fe*~
Fe(lll)
complex
I nterface
Fig. 8. (al Scheme of the forward (aqueous ~ organic) mass transfer of Fe(Ill) at t - O; Vo =rate of Fe(III) interfacial adsorption to give the interracial complex; Vd. =rate of the opposing desorption; V.,, = rate of the forward desorption. (b) Scheme of the reverse (organic~aqueous) mass transfer of Fe(lll) at t = O; V,, = rate of Fe(III) inteffacial adsorption to give the interfacial complex; Vo,, = rate of the opposing desorption; V..,- rate of the forward desorption. (c) Scheme of the mass transfer (either forward or reverse) of Fe(III) at t~O; V,, and V. = rates of interfacial adsorption of Fe(III); V.,, and Vdo = rates of desorption respectively towards the aqueous and organic phase.
/ I N ( ~o} 8 No 9--H
K I (~HI) V* = KRa~-+
CHD
I IO)
with KR = K3/K2. Equation (10) represents well the behaviour of the experimental data of Fig. 7. The values of K, and KR, calculated by curve fitting, are the following: K, = (4.80+0.10)x 10 2min ~; KR = (1.16-+0.10);< 10 moles/l) ~. The solid lines reported in Fig. 7 have been calculated by eqn (10) using the K, and K~ values reported above. The agreement between calculated curves, and experimental points is satisfactorily good.
Reverse initial-mass transfer rate Figure 9 a,b,c, (full points), shows the experimental reverse initial mass transfer rates (t = 0), V* = WlF-el, (ideal behaviour of iron containing organic species is also assumed), plotted as function of the following variables: [Fe] (initial Fe(III) concentration in t h e organic phase), ((~,D, [HC104] constant, Fig. 9a), aw ([Fe], Cm, constant,
1692
P.R. DANESIand R. CHIARIZIA (a)
(b) 8
8
0
U7 >6
7
O--o--o
OmO
~
5 4
0'05
0
OllO C~,
M
i
i
i
I0
15
20
•
i
25 [t~e] I0 4 , M
8 •
•
6i
)
%
/ /o
74
2
05
I0
1"5
'20
c7H+
Fig. 9. ¥* vs concentration (or activity) plots; (a) V* vs [Fe] (CHD=0.025M, [HCIO4]= 1.11M), (b) ¥* vs (~im([HC104]= 1.18M, [Fe] = 7.01 × 10-"M), (c) V* vs aH+(C'H,= 0.0125M, [Fe] = 1.161× I0-3 M); interface area= 18.9cm2, aqueous and organic-volumes=40cm3; full points=experimental values; continuous lines calculated through eq. (13)with/(~ = 8.57× l0-3 and KR = 1.16× 10-1. m
Fig. 9b), CHo ([Fe], [HCIO4] constant, Fig. 9c). To derive the equation for the initial reverse mass transfer rate we have to consider that the chemical'reactions given by the interracial complex (FeD2+)~ must be the same as those previously reported. The situation holding at t = 0 is schematically shown in Fig. 8(b) when Fe(III) is initially present only in the organic phase. We have in this case: 1. Formation of the inteffacial complex, (FeD2+)~,from the iron containing complex formed between HD and iron present in the organic phase. This reaction can be written as the result of two series reactions, the first one (slow) having a rate proportional to the organic iron concentration, such as )
FeHm 3D,,,
+
(FED2)~ +(Hm 3Din-2)
General equation for the mass transfer rate Once the parameters Kl, /~1 and KR have been evaluated and the kinetic scheme identified by studying the initial transfer rates (t = 0), a general equation can be derived to represent the dependence of the mass transfer rate from the composition of the system and time. The situation holding at t # 0 is schematically reported in Fig. 8c. Also in this case the concentration of the interfacial complex can be evaluated from the interfacial steady state condition
Vo+Vo=V~o+V.o and eqns (3), (5), (7), (12), i.e.
(11) [FED2+], _ K~aFo3++/(,[Fe] K3a~+ + K:CHD
and a second one (fast) consisting in the readjustement of the freed polymerized HD. The rate of this process will be Vo = K,[Fe]
(12)
where [Fe] = [FeHm-3D~] = initial organic iron concentration. 2. The reactions of the interracial complex with the polymerized extractant to form again the iron-HD organic complex, and with hydrogen ions to give Fe 3+ ions in the aqueous phase, will be the same as reactions (6) and (4). Following the same procedure as in the forward initial mass transfer rate to derive the concentration of the interracial complex, we obtain /~la
V* =
(11)
2+
(13)
(15)
where avo3+and [Fe], arv and CHO are the values at time t # 0. If [W] and CHD are much I j g e r then the total iron concentration [Fe],o, = [Fe 3+]+ [Fe], they can be considered as time independent. The overall mass transfer of iron from the water to the organic phase will then be given by V = V, - Va,.
(16)
After some rearrangement it follows V = [Fe3+]a - / 3 _ y
d[Fe 3+] dt
[17]
wher~ -
+
-
2
a = yFo(c~o,~3K1CnD K1KRaH+,
Equation (13) represents well the behaviour of the data of Fig. 9. In this case only one parameter, /(~, has to be evaluated by curve fitting. Its value was found to be /(~ = (8.57 - 0.50) x 10-3 min-1. The solid lines reported in Fig. 9 have been calculated by eqn (13) using the K~ and KR values reported above. Also in this case the agreement between calculated curves and experimental points is good.
/3 = [Fe]tot/(1KRa ~+, y = KRa z8÷+ CHD. The integrated form of eqn (17) is In {([Fe],o,a -/3)/([Fe 3+]a - fl)} = (a / y) t.
(18)
A test of the validity_of eqn (18) has been performed in conditions where Cm~ and aK+,>[Fe]tot. In Fig. 10
1693
Transfer rate of iron(Ill) in the biphasic system HC10~
'5 'O
of HD molecules. The experimental results have enabled us to elaborate a definite scheme of rate determiining reactions and to calculate the elementary rate constants. By means of these parameters a general equation representing the dependence of the mass transfer rate from both the chemical composition of the system and time has been obtained.
/O/O
4 f)
•/
•/
•/ "2. 5 ')
•/•/
•/
REFERENCE
a~ / !?
o/o/
°/° o/
o/ O/
+~
,) ~/
O/
C
50 Tree,
IO0
mlR
Fig. 10. Log{(c~[Fe~*],,,-/3)/(a[Fe~+]-/3)} vs. time plot for a kinetic experiment performed in the following conditions: initial IFe 3+] = 8.60 x 10 ,4 M,, OHr,= 0.066M, [HCIO4]= 0.451M, calculated slope of the straight line=c~/2.303~,= 6.4 × 10 ~ min ': full points = experimental values. log {([Fe],o,C~ -/3)/([Fe~']o~ -/3)} is plotted vs t. A straight line is obtained and the agreement between its slope value and the calculated a/2.303y coefficient is satisfactory. Moreover a comparison between the equilibrium D value and the one calculated from eqn (18) is possible. In fact for t = ~, [Fe+3]~q =/3/o~ = 0.91 x 10 -4. It follows a calculated D value of 8.4. The experimental D value in the same conditions, obtained from the data of Fig. 4 is about 10. The agreement is reasonably good. CONCLUSION In our experimental conditions the transfer rate of Fe(III) can be explained in terms of chemical reactions, occurring at the interface, between the iron present in the bulk solutions and the interfacially adsorbed monolayer
1. C. Hanson, M. A. Hughes and J. G. Marsland, Proc. Int. Solvent Extraction Conference 1974, Vol. 3, p. 2401. Society of Chemical Industry, London 11974). 2. G. A. Yagodin, V. V. Tarasov and N. F. Kizim, Proc. Int. Solvent Extraction Conference 1974, Vol. 3, p. 2541. Society of Chemical Industry, London 11974). 3. F. Baumg/irtner and L. Finsterwalder, J. Ph)~. Chem. 7,11, 108 (1970). 4. J. W. Roddy and C. F. Coleman, Solvent Extraction Rec. 1 ( l ) 63-91 11971). 5. D. S. Flett, D. N. Okuhara and D. R. Spink, J. lnore. NucL Chem. 35, 2471 11973). 6. D. S. Flett, J. A. Hartlage, D. R. Spink and D. N. Okuhara, J. lnorg. Nud. Chem. in press. 7. D. F. Flett. M. Cox and J. D. Heels, J. lnor~,,. Nucl Chem. in press. 8. P. R. Danesi, G. D'Alessandro, R. Chiarizia and B. Scuppa, .I lngrg. Nucl. Chem. in press. 9. R. Chiarizia, P. R. Danesi, M. A. Raieh and (i. Scibona, .I. lnorg. Nucl. Chem. 37, 1495 11975). 10. P. R. Danesi, R. Chiarizia and G. Scibona, J. lnor,,,. Nncl. Chem. 35, 3926 11973). 11. J. M. White, P. Tang and N. C. Li, J. lnor~,,. NucL Chem. 14. 255 11960). 12. D. F. C. Morris and H. R. Wilson, J. lnor,e. NucL Chem. 31, 1532 (1969). 13. D. F. C. Morris and P. J. Sturgess, Ele(trochim. A(ta 14, 629 11969). 14. Yung-Chi Wu and W. J. Hamer, NBS Rep. 10002, part XIII (1969). t5. Y. Marcus and A. S. Kertes, Ion Exchange and Sohent Extraction of Metal Complexes, p. 929. Wiley-lnlerscience. London 11969). 16. C. W. Davies, J. Chem. Soc. 2093 11938).
Printed in Great Britain
TRANSFER RATE OF SOME TERVALENT CATIONS IN THE BIPHASIC SYSTEM HC104, WATERDINONYLNAPHTHALENESULFONIC ACID, TOLUENE--I TRANSFER
RATE OF IRON(III) DETERMINED CHEMICAL
BY
REACTIONS
P. R. DANESI, K CHIARIZ1A and A. SALTELL[ Laboratorio Chimica Fisica, C.S.N. Casaccia, C.N.E.N., Rome, Italy (First received 26 June 1975: in ret:isedform 15 October 1975) Abstract--The transfer rate of iron(Ill) ions between aqueous perchlor~c acid solutions and toluene solutions of dinonylnaphthalenesulfonic acid has been studied by using a constant interracial area cell. The initial forward and reverse transfer rates have been evaluated as function of the interfacial area, volume and chemical composition of the system. The results have been explained in terms of interfacial reactions which rate constants have been evaluated. An equation capable to describe the dependence of the transfer rate from both the composition of the system and the time has been derived. INTRODUCTION
THE mass transfer rate of metal cations in liquid-liquid solvent extraction systems is in the most general case a function of both the kinetics of the chemical reactions taking place in the system and the rates of diffusion of the involved species in the two liquid phases. The simultaneous presence of both chemical reactions and diffusion processes in liquid-liquid mass transfer has been recently reviewed[l] and several mathematical diffusion models with heterogeneous chemical reactions have been elaborated[l,2]. However several cases seem to exist where the rate of mass transfer of cations in liquid-liquid solvent extraction systems is mainly controlled by interfacial processes [3-7]. In this instance the mathematical model of the system greatly simplifies and providing both the interfacial and thermodynamic properties of the system are well known a kinetic scheme can also be derived. In the present paper the mass transfer of Fe(lli) between aqueous perchloric acid solutions and toluene solutions of the liquid cation exchanger dinonylnaphthalenesulfonic acid, HD, has been studied. It was in fact presumed that with such a system, previously well characterized by us in its interfacial[8] and thermodynamic [8, 9] properties, it would have been possible to receive some further indication on the role of interfacial chemical reactions on the overall mass transfer rate. EXPERIMENTAL Materials Iron powder 97% pure, HCIO4 70% RPE-ACS and toluene RPE-ACS were Carlo Erba products. An Fe(lII) perchlorate solution in HCI04 has been prepared by dissolving the iron powder in hot perchloric acid and boiling for several hours in order to eliminate Fe(II). To test that the solution contained no Fe(lI) the reaction with orthophenantroline has been used. The Fe(III) content has been determined by titration with EDTA using sulfosalicilic acid as indicator. The hydrogen ion concentration has been determined by a standard procedure. A stock solution of dinonylnaphthalenesulfonic acid, HD, supplied by Columbia Organic Chemicals Co. Inc., and purified as previously reported[10], was prepared by dissolving HD in toluene and by determining its concentration by potentiometric titration[4]. A stock solution of Fe(Ill) in the organic phase has been prepared by shaking equal volumes of the aqueous Fe(III) perchlorate stock solution with a concentrated solution of HD in toluene. Its concentration has been evaluated by the distribution coefficient. All the solutions used in the experiments have been prepared by mixing the stock solutions either with bidistilled water or HCIO4
standard solutions or toluene. All other reagents were Carlo Erba analytical grade purity products. The calomel saturated NaC1 reference electrode and the agar-agar NaCI bridge have been prepared according to standard polarographic procedures. Technique A constant interfacial area cell (Fig. 1) has been employed in most of the experiments. This cell permits both forward and back kinetic experiments of the Fe(II1) extraction, keeping the interface quiescent, with a continuous monitoring of the concentration of Fe(lII) in the aqueous phase as function of time. To the purpose the cell contains a platinum rotating electrode Metrohom SM/5 and a calomel NaC[ saturated electrode connected via an agar-agar NaCI bridge. These two electrodes are used in an amperometric system to detect iron. To the purpose a Metrohom Potentiograph model E 436 was used. Preliminary experiments have shown that applying a polarization voltage equal to zero a very good linear proportionality exists between the current and the Fe(lll) concentration in the aqueous phase. Moreover the current variations caused by Fe(IIl) concentrations variations are practically instantaneous. Therefore it was possible to folliow the concentration variations of Fe(III) with time by simply recording the current vs time curves and by properly calibrating the system each time using Fe(III) solutions of known concentration. The forward extraction kinetics of Fe(IIl) were therefore performed in the following way: a known volume (generally 40 cm ') of HCIO~ of known concentration was placed in the cell and after checking the instrumental zero the circuit was closed starting at the same time the stirring of the cell. A constant residual current was generally established within 2 hr. Afterwards progressive amounts of the stock solution of Fe(III) were added obtaining in this way a current vs concentration calibration curve. A volume of HD solution in toluene equal to that of the aqueous phase was then added by stopping, during the time of the addition of the organic phase, both the rotation of the platinum electrode and the stirring. The starting of the kinetics then coincided with the new starting of the stirring and of the rotation of the Pt electrode. The current decrease with time was then directly correlated to the concentration vs time curve since the recorder chart speed was known. The organic phase was each time preequilibrated with an aqueous HCIO4 solution having the same concentration as that in the cell. Figure 2 reports a typical forward extraction kinetics experiment. The backward extraction kinetics (organic Fe(IIl)~aqueous Fe(IIl)) were performed as follows: after addition in the cell of the aqueous perchloric acid and attainement of the residual cmrent an organic phase consisting of a toluene solution of HD and Fe(II1) of known concentration was added to the aqueous phase solution following the same procedure as for the forward kinetics. After the back extraction had proceeded for some time, the extraction of iron was stopped by removing the organic phase by careful suction followed by repeated washing with pure toluene to remove
1687
P. R. DANESI and R. CHIARIZIA
1688
Stirrermotorand I gearOox II
----I ~"
Toth..... ......ording
II
|lOutlet
~I | - I I
J Agar-Agar ~O~s.S.
J
1"1I J [ ]
"~
|
~
TOlhe recording amp..... for (+)
(LI( "m III I
~~ -1 -- tiJl '~1] AqueousJ
-Th..... IOI ~,~ inlet
Glass fritt Fig. 1. Apparatus used to study the transfer rate of Fe(III) by continuous monitoring of the Fe 3+ aqueous concentration through amperometry. Interface radius=2.45cm; interface geometrical area=18.9cm2; stirrer blades = 20 x 5 x 2 ram; distance between stirrer blades = 20 mm; volume of aqueous and organic phases ~40 cm3; Pt electrode rotation speed = 750 rpm; stirrer rotation = 1I0 rpm,
I10 I00 90 80 70
~ 6o i_ ~ 50 C3
L 40
20 I0 O
OI 234567891011P2 Time, min
Fig. 2. Typical forward kinetics; [Fe3÷] = 6.87 x 10-" M, C'm, = 0.046 M, [HCIO4] = 0.451 M; Volume of aqueous and organic phases = 43 cm3; ® residual current; ®, ®, ®, ®, ®, ® currents corresponding to progressive volume additions of the aqueous Fe ÷3 stock solution; ® stopping of the stirrer and Pt rotating electrode and addition of the organic phase; ® starting of the forward extraction kinetics; # current v s time plot. traces of HD. At this point the calibration curve, current vs concentration, was started by adding to the aqueous phase known amounts of the Fe(III) stock solution. Figure 3 reports a typical back extraction kinetics experiment. The temperature of the cell was kept at 25.00-0.05°C by a circulating ultrathermostat. The stirring speed was kept constant in all experiments at 110 rpm. Due to the slow extraction kinetics the initial transfer rates were obtained by the slope of the straight line which was well representing the first five minutes of the kinetic experiments. The forward kinetic experiments performed to obtain information on the influence on the rate of the interfacial area have been obtained by using a cell having the same design as that of Fig. 1 but a much larger diameter. The interracial area was then reduced by means of teflon interface limiting rings.
solutions and HD toluene solutions for the following experimental conditions: temperature = 25°C, [Fe+3] = 9.85 x 10-4 M, [HC10,] =0.53-2.00M, C,D (total organic concentration of HD) = 0.018-0.23 M, have been obtained by shaking for 60rain (enough to attain equilibrium) 10cm3 of each phase in a mechanical shaker. After centrifugation and phase separation the equilibrium concentration of Fe 3+, reduced to Fe 2÷ with hydroxylamine chloridrate, was determined spectrophotometrically at 510nm by adding ortophenantroline and 2M sodium acetate buffer and using a Beckmann DK2A spectrophotometer. Figure 4 shows the log D vs log (~nDand log D vs log [HC104] plots. Both are well fitted by straight lines of slopes +1 and -3 respectively, thus indicating that the extraction of iron(III) is well represented by the following stoichiometry
Equilibrium distribution studies o.f Fe(III) The distribution ratio, D of Fe(IlI) between aqueous HCIO4
Fe3++ ( H D ) ~ F e H , , - 3 D m + 3H +
(1)
Transfer rate of iron(III) in the biphasic system HCIO~
1689
2O !
90
8Ovo
~
~®
60
-~ 50
J<~7b
2()
i
o
2
4
,
i
6
i
i
8
i
i
,
i
i
ko ~2 i~
Time. rain Fig. 3. Typical reverse kinetics; [Fe] = 1.0x 10 ~M, CH~= 0.025 M, IHCI0~] ~ l.ll M: volume of aqueous and organic phases = 40 cm3; @ residual current; ~ stopping of the stirrer and Pt rotating electrode and addition of the organic phase; ~, starting of the reverse extraction kinetics; ® current vs, time plot: ® removal of the organic phase. ~, @. ®, ®, Q, ~ currents corresponding to progressive volume additions of the aqueous Fe ~4 stock solut]o . '
n
\
\ ~
°~°
!O
¢ eo '
./
o"
%
/ai
\ 2
-I
@
4
IogCHo, M
/
-I
J-I i, I
O log [ H C I 0 4 ] , M
Fig. 4. (a) log D (distribution ratio of Fe(lI[)) vs log CHI>plot; (b) log D vs log [HCIO,] plot: (a) [HCI(L} : 1.02 M, {Fe '+] = 9.85 x 10 ~ M, T = 25°C; (b) (~,i~ = 0.046 M, [Fe 3*] = 9.85 x 10 4 M, T = 25°C; The slopes of the straight lines are +1 for (a) and -3 for (bl. where m ~ 7 is the degree of polymerization of HD [8] and the bar indicates organic species. These results well agree with those of references 11-13. RESULTS AND DISCUSSION
Dependence of the initial forward mass transfer rate on the volume of the phases, and interracial area In order to examine the dependence of the mass transfer kinetics from bulk chemical reactions, initial mass transfer rates were obtained as function of the volume of the phases and of the interfacial area. The forward mass transfer rates at time zero divided by the initial iron(III) concentration, V' = V/[Fe ~+] (with V = -d[Fe3+]/dt) are plotted in Fig. 5 as function of the volume of the aqueous phase (an equal volume of organic
phase was used) at constant interfacial area. A straight line passing through the origin is obtained thus indicating that reactions occurring in the bulk of the phases are not rate determining. Figure 6 shows a plot of V' vs A~ (geometrically calculated interfacial area) at constant volume. Also in this case a straight line is obtained. However the line does not go through the origin indicating that the real interfacial area and the one calculated through the cell diameter are not coinciding. This fact can be explained by considering that the platinum electrode, rotating at 750 rpm, causes a vortex of fixed area which contributes with an additional constant term to the geometrically calculated interfaciat area. However the data of Fig. 6 indicate that the rate determining processes are taking place in proximity of the
1690
P . R . DANESI and R. CHiARIZ~A
1.5
15
NO--L0
0.5
0,5
0
0 0
I
3 I/V,
20
cc-a I02
i
i
40
60
A I , crn 2
Fig. 5. V' vs 1IV (volume) plot; V is the volume of the aqueous and organic phases; [Fe ~*]=6.87x 10-~M, C ~ = 0 . 0 2 0 M , [HCI04] = 0.451 M; interface area = 18.9 cm 2.
Fig. 6. V' vs A~ (geometrical interfacial area) plot; [Fe 3+] = 7.30 x 10 4 M, CHD = 0.057 M, [HCI04] = 0.451 M; aqueous and organic volumes = 108 cm ~.
interface (either interracial reactions or diffusion processes in the interracial liquid films).
yF,tao,)3; yF,tao,)3 has been estimated through the Davies equation[16] assuming that the ionic strength is only determined by the perchloric acid. In order to explain the experimental results of Fig. 7 a model based on rate determining interfacial chemical reactions has been elaborated. These reactions can be formally treated as adsorption-desorption processes as in Ref. [3]. On the basis of the known interracial properties of the HDtoluene-HC104 system[8] it has been assumed that an interracial complex between iron ions and the interracial monolayer of HD molecules is formed. This interracial complex can in turn react either with the aqueous or the organic phase. The nature of this interracial complex must be of course the same independently of which phase the iron is coming from. Further the assumption is made that the steady state is obtained at the interface instantaneously (t = 0) i.e. when no net mass transfer has yet taken
Forward initial mass transfer rates Figure 7 a, b, c (full points) shows the experimental forward initial mass transfer rates (t = 0) divided by the Fe(III) activity, V* = V/aFe3+, plotted as function of the following variables: ale3+ (6',D, [HC104] constant, Fig. 7a), an+ ([Fe3+], CHI~constant, Fig. 7b), t~im ([Fe3+], [HC104] constant, Fig. 7c). The mean activity of the H ÷ in the aqueous perchloric acid (all+) has been calculated as the product of the mean molar activity coefficient, yuc~o,, and the molarity of the solution, yHao, has been obtained from literature data [14, 15]. The mean molar activity coefficient of Fe(III) in the aqueous phase (aFP) has been computed as the product of the iron molar concentration and the mean molar activity coefficient of Fe(III) perchlorate,
O•O \o
4.0 ~_O3.0
(a) (b)
~O 5,(3
%\
.;
~
2.0
4.° -- • 0 0
• •
o.O
°o'o~
o%---
3.G
1.0
o!5 GH+ ,:o o/°"O/°--
40 N 3,0 O
2.0
/
/
2'0
• 4[0
6'.0
810
aFe+ •,[0 4 (c)
1.0 0
0.05
O. I0
0 15
CHD, M
Fig. 7. V* vs concentrations (or activity) plots; (a) V* vs aFe 3+ ( C B D = 0.046 M, [HC10,] = 0.451 M), (b) V* vs a,÷(CHr, = 0.046 M, [Fe 3÷] = 6.87 x 10 "M), (c) V* vs C'HD([HCIO,] = 0,451 M, [Fe 3+] = 6.87 x 10 4 M); interface area= 18.9cm 2, aqueous and organic volumes=43cm3; full points=experimental values; continuous lines calculated through eqn (10) with K, = 4.80 x 10-2 and KR = 1.16 x 10-1.
1691
Transfer rate of iron(Ill) in the biphasic system HCIO, place. The situation holding at t = 0 is schematically reported in Fig. 8a, when Fe(III) is initially present only in the aqueous phase. In order to derive an equation capable to interpret the experimental initial mass transfer rates the following reactions have been assumed: 1. Formation of an interfacial complex between iron ions and the HD molecules which are absorbed at the interface Fe '' + 2(HD),
)(FeD2+),+2H *
(2)
where the subscript i indicates interfacial species. The rate of this process will be
determined [8] that in our experimental conditions [HD], is constant. 2. Dissociation of the interracial complex, (FeD,.')~, to form again iron ions in the aqueous phase and interracial HD (FeD2~)~ +2H +
)Fe '' + 2(HDL
The rate of this process, assuming ideal behaviour of inteffacial species, will be proportional to the concentration of the interfacial complex and to the square of the aqueous hydrogen ion activity
Vd,, = K~[FeD~'I~a~. V, = K ~av.,.[HD]f = K,aw,*
V:
Interfacial complex
org.
Fe ~
(FeD.+)~ +(HD),., ~ F e H . , ,D,,,+2(HD), + H .
V.L. = K~[FeD.~-],[(HD),,,] = K~[FeD:+],d,,.
Interface
Vo (b) Interracial
aq
3. Reaction of the interracial complex with the HD molecules present in the organic phase to form the final product of the overall mass transfer reaction. Since HD is very strongly polymerized in our experimental conditions [8] this reaction can be written as follows (6)
The rate of this process, assuming that also organic species behave ideally, will be proportional to the concentration of the interracial complex and of the polymerized HD
_
aq
(5)
(3)
where K',[HDI~: = K,, since it has been experimentally
(o)
(4)
complex
where K:= K'/m since C.o=[HD]+ m[(H-I)).,I= m [(HD),.]. In order to express the overall rate of mass transfer of iron into the organic phase, Vj., as function of the experimentally studied variables ave>, a . , C.D and the specific rates K,, K.., K3, the concentration of the interfacial complex has to be derived by using the: flux equation at the interface
org
V, -V,l, - V,~,,= 0
Fe(lll)
(7)
(8)
It follows then
Interface
KIOFe3" [FeD:~]~ = K~a~ + K2C.,
(9)
By substituting eqn (9) in eqn (7) we get (c)
I[/°nterfacial
-- ~
aq.
org.
Fe*~
Fe(lll)
complex
I nterface
Fig. 8. (al Scheme of the forward (aqueous ~ organic) mass transfer of Fe(Ill) at t - O; Vo =rate of Fe(III) interfacial adsorption to give the interracial complex; Vd. =rate of the opposing desorption; V.,, = rate of the forward desorption. (b) Scheme of the reverse (organic~aqueous) mass transfer of Fe(lll) at t = O; V,, = rate of Fe(III) inteffacial adsorption to give the interfacial complex; Vo,, = rate of the opposing desorption; V..,- rate of the forward desorption. (c) Scheme of the mass transfer (either forward or reverse) of Fe(III) at t~O; V,, and V. = rates of interfacial adsorption of Fe(III); V.,, and Vdo = rates of desorption respectively towards the aqueous and organic phase.
/ I N ( ~o} 8 No 9--H
K I (~HI) V* = KRa~-+
CHD
I IO)
with KR = K3/K2. Equation (10) represents well the behaviour of the experimental data of Fig. 7. The values of K, and KR, calculated by curve fitting, are the following: K, = (4.80+0.10)x 10 2min ~; KR = (1.16-+0.10);< 10 moles/l) ~. The solid lines reported in Fig. 7 have been calculated by eqn (10) using the K, and K~ values reported above. The agreement between calculated curves, and experimental points is satisfactorily good.
Reverse initial-mass transfer rate Figure 9 a,b,c, (full points), shows the experimental reverse initial mass transfer rates (t = 0), V* = WlF-el, (ideal behaviour of iron containing organic species is also assumed), plotted as function of the following variables: [Fe] (initial Fe(III) concentration in t h e organic phase), ((~,D, [HC104] constant, Fig. 9a), aw ([Fe], Cm, constant,
1692
P.R. DANESIand R. CHIARIZIA (a)
(b) 8
8
0
U7 >6
7
O--o--o
OmO
~
5 4
0'05
0
OllO C~,
M
i
i
i
I0
15
20
•
i
25 [t~e] I0 4 , M
8 •
•
6i
)
%
/ /o
74
2
05
I0
1"5
'20
c7H+
Fig. 9. ¥* vs concentration (or activity) plots; (a) V* vs [Fe] (CHD=0.025M, [HCIO4]= 1.11M), (b) ¥* vs (~im([HC104]= 1.18M, [Fe] = 7.01 × 10-"M), (c) V* vs aH+(C'H,= 0.0125M, [Fe] = 1.161× I0-3 M); interface area= 18.9cm2, aqueous and organic-volumes=40cm3; full points=experimental values; continuous lines calculated through eq. (13)with/(~ = 8.57× l0-3 and KR = 1.16× 10-1. m
Fig. 9b), CHo ([Fe], [HCIO4] constant, Fig. 9c). To derive the equation for the initial reverse mass transfer rate we have to consider that the chemical'reactions given by the interracial complex (FeD2+)~ must be the same as those previously reported. The situation holding at t = 0 is schematically shown in Fig. 8(b) when Fe(III) is initially present only in the organic phase. We have in this case: 1. Formation of the inteffacial complex, (FeD2+)~,from the iron containing complex formed between HD and iron present in the organic phase. This reaction can be written as the result of two series reactions, the first one (slow) having a rate proportional to the organic iron concentration, such as )
FeHm 3D,,,
+
(FED2)~ +(Hm 3Din-2)
General equation for the mass transfer rate Once the parameters Kl, /~1 and KR have been evaluated and the kinetic scheme identified by studying the initial transfer rates (t = 0), a general equation can be derived to represent the dependence of the mass transfer rate from the composition of the system and time. The situation holding at t # 0 is schematically reported in Fig. 8c. Also in this case the concentration of the interfacial complex can be evaluated from the interfacial steady state condition
Vo+Vo=V~o+V.o and eqns (3), (5), (7), (12), i.e.
(11) [FED2+], _ K~aFo3++/(,[Fe] K3a~+ + K:CHD
and a second one (fast) consisting in the readjustement of the freed polymerized HD. The rate of this process will be Vo = K,[Fe]
(12)
where [Fe] = [FeHm-3D~] = initial organic iron concentration. 2. The reactions of the interracial complex with the polymerized extractant to form again the iron-HD organic complex, and with hydrogen ions to give Fe 3+ ions in the aqueous phase, will be the same as reactions (6) and (4). Following the same procedure as in the forward initial mass transfer rate to derive the concentration of the interracial complex, we obtain /~la
V* =
(11)
2+
(13)
(15)
where avo3+and [Fe], arv and CHO are the values at time t # 0. If [W] and CHD are much I j g e r then the total iron concentration [Fe],o, = [Fe 3+]+ [Fe], they can be considered as time independent. The overall mass transfer of iron from the water to the organic phase will then be given by V = V, - Va,.
(16)
After some rearrangement it follows V = [Fe3+]a - / 3 _ y
d[Fe 3+] dt
[17]
wher~ -
+
-
2
a = yFo(c~o,~3K1CnD K1KRaH+,
Equation (13) represents well the behaviour of the data of Fig. 9. In this case only one parameter, /(~, has to be evaluated by curve fitting. Its value was found to be /(~ = (8.57 - 0.50) x 10-3 min-1. The solid lines reported in Fig. 9 have been calculated by eqn (13) using the K~ and KR values reported above. Also in this case the agreement between calculated curves and experimental points is good.
/3 = [Fe]tot/(1KRa ~+, y = KRa z8÷+ CHD. The integrated form of eqn (17) is In {([Fe],o,a -/3)/([Fe 3+]a - fl)} = (a / y) t.
(18)
A test of the validity_of eqn (18) has been performed in conditions where Cm~ and aK+,>[Fe]tot. In Fig. 10
1693
Transfer rate of iron(Ill) in the biphasic system HC10~
'5 'O
of HD molecules. The experimental results have enabled us to elaborate a definite scheme of rate determiining reactions and to calculate the elementary rate constants. By means of these parameters a general equation representing the dependence of the mass transfer rate from both the chemical composition of the system and time has been obtained.
/O/O
4 f)
•/
•/
•/ "2. 5 ')
•/•/
•/
REFERENCE
a~ / !?
o/o/
°/° o/
o/ O/
+~
,) ~/
O/
C
50 Tree,
IO0
mlR
Fig. 10. Log{(c~[Fe~*],,,-/3)/(a[Fe~+]-/3)} vs. time plot for a kinetic experiment performed in the following conditions: initial IFe 3+] = 8.60 x 10 ,4 M,, OHr,= 0.066M, [HCIO4]= 0.451M, calculated slope of the straight line=c~/2.303~,= 6.4 × 10 ~ min ': full points = experimental values. log {([Fe],o,C~ -/3)/([Fe~']o~ -/3)} is plotted vs t. A straight line is obtained and the agreement between its slope value and the calculated a/2.303y coefficient is satisfactory. Moreover a comparison between the equilibrium D value and the one calculated from eqn (18) is possible. In fact for t = ~, [Fe+3]~q =/3/o~ = 0.91 x 10 -4. It follows a calculated D value of 8.4. The experimental D value in the same conditions, obtained from the data of Fig. 4 is about 10. The agreement is reasonably good. CONCLUSION In our experimental conditions the transfer rate of Fe(III) can be explained in terms of chemical reactions, occurring at the interface, between the iron present in the bulk solutions and the interfacially adsorbed monolayer
1. C. Hanson, M. A. Hughes and J. G. Marsland, Proc. Int. Solvent Extraction Conference 1974, Vol. 3, p. 2401. Society of Chemical Industry, London 11974). 2. G. A. Yagodin, V. V. Tarasov and N. F. Kizim, Proc. Int. Solvent Extraction Conference 1974, Vol. 3, p. 2541. Society of Chemical Industry, London 11974). 3. F. Baumg/irtner and L. Finsterwalder, J. Ph)~. Chem. 7,11, 108 (1970). 4. J. W. Roddy and C. F. Coleman, Solvent Extraction Rec. 1 ( l ) 63-91 11971). 5. D. S. Flett, D. N. Okuhara and D. R. Spink, J. lnore. NucL Chem. 35, 2471 11973). 6. D. S. Flett, J. A. Hartlage, D. R. Spink and D. N. Okuhara, J. lnorg. Nud. Chem. in press. 7. D. F. Flett. M. Cox and J. D. Heels, J. lnor~,,. Nucl Chem. in press. 8. P. R. Danesi, G. D'Alessandro, R. Chiarizia and B. Scuppa, .I lngrg. Nucl. Chem. in press. 9. R. Chiarizia, P. R. Danesi, M. A. Raieh and (i. Scibona, .I. lnorg. Nucl. Chem. 37, 1495 11975). 10. P. R. Danesi, R. Chiarizia and G. Scibona, J. lnor,,,. Nncl. Chem. 35, 3926 11973). 11. J. M. White, P. Tang and N. C. Li, J. lnor~,,. NucL Chem. 14. 255 11960). 12. D. F. C. Morris and H. R. Wilson, J. lnor,e. NucL Chem. 31, 1532 (1969). 13. D. F. C. Morris and P. J. Sturgess, Ele(trochim. A(ta 14, 629 11969). 14. Yung-Chi Wu and W. J. Hamer, NBS Rep. 10002, part XIII (1969). t5. Y. Marcus and A. S. Kertes, Ion Exchange and Sohent Extraction of Metal Complexes, p. 929. Wiley-lnlerscience. London 11969). 16. C. W. Davies, J. Chem. Soc. 2093 11938).