ADSORPTIO:~ AND BYDROLYSIS OF [Co(NH&I3+ ON CARBON
649
Effect of Surface Groups of Carbon on the Adsorption and Catalytic Base Hydrolysis of a Hexaamminecobalt(II1) Ion by A.kira Tomita* and Yasukatsu Tamai Chemical Research Institute of Non-Aqueous Solutions, Xohoku University, Katahira-cho, Sendai, Japan (Receikved March 8, 1970) Publication costs borne completely by The Journal of Physical Chemistry
Kinetic data are presented for the adsorption of a hexaamminecobalt(II1)ion from an aqueous solution onto several kinds of carbon and for the base hydrolysis catalyzed by carbon. The content of surface acidic groups of carbons, including a carbon black and an active carbon, was modified by oxidation or methylation. The pEI value of solution has a significant effect on the adsorption rate: an addition of NaOH increases the rate, and :BCl retards the adsorption. The observed rate constant increases with the amount of surface acidic groups and the amount of carbon. These results are consistent with the proposed mechanism which involves adsorption of complex ion on such acidic groups of carbon. Apparent activation energy for adsorption on carbon was estimated at 21 23 kcal/mol. The rate constant for base hydrolysis of the same complex ion was also found to have a close relation to the amount of surface acidic groups of carbon catalysts. It i s suggested that the active sites for hydrolysis may be such acidic groups which are also the active sites for the adsorption. N
Introduction Carbon catalyzers a number of different kinds of chemical reacti0ns.l I n the field of coordination chemistry, carbon is often used as a catalyst in the preparation nf Go(II1) complexes2 or in the racemization of optically active Co(II1) complexes.* The mechanistic details of this heterogeneous reaction or the mechanism of the adsorption process which must be the most important) step of the reaction are not understood ai prcbsent. The study of the adsorption and the reaction of metal complexes on carbon would give not only a guiding principle for preparing new coordination compounds but also some information about the reaction mechanism on metal catalysts supported on (carboncarriers. Several factors may influence the rate and the mechanism of reaction on carbon. I n the present paper, we shall confine our attention to the chemical nature of the carbon surfacr. The functional groups on carbon have been characterized by various organic and physical methods.’i Boehmb reported the presence of four different kinds of ,acidic surface oxides. The nature of carbon surface varies with the kind, the amount, and the distribution of these functional groups. It has been sho-xn that acidic groups on carbon play a very important role in adsorption of organic compounds6J from aqueous solution. I n the prec,ent study, we prepared carbons having different amounts of acidic groups by various methods of modification and investigated the relationship between the surface acidity and the adsorption rate of a [Co(NHz,l;,ls+ion onto carbon. The second object of this paper is to determine the dependence of the
catalytic activity of carbon on the surface acidity. The test reaction chosen for this purpose was the base hydrolysis of [Co(NH&I3+ into [ C O ( W H ~ ) ~ ( O H ) ] ~ + in the presence of carbon.S
Experimental Section Materials. Deionized distilled water and reagent grade chemicals were employed. [Co(NH3)6]CL was prepared as described in the literatureg and characterized analytically and spectrophotometrically. Two kinds of carbon with different surface area were employed in this work: (1) a carbon black “Seast 305 (HAF)” supplied by Tokai Electrode Mfg. Co., and (2) a commercial active carbon (AC) obtained from Kanto Chemical Co., Inc. Modijication of Carbon. To prepare carbon samples of various known contents of surface oxides for adsorption study, the surface of the carbon was oxidized with ”Os, (NH4)2S20s,and O2 by the method of Boehm.6 Partial methylation of surface carboxyl or (1) R. W. Coughlin, Ind. Eng. Chem. Prod. R e s . Develop., 8, 12 (1969). (2) F. Basolo and R. G. Pearson, “Mechanisms of Inorganic R e x tions,” Wiley, New York, N. Y., 1958, p 355. (3) B. E. Douglas, et al., J . Amer. Chem. Soc., 76, 1020 (1954); J . Inorg. Nucl. Chem., 24, 1355, 1365 (1962); 26, 601, 609 (1964). (4) J. B. Donnet, Carbon, 6, 161 (1968). (5) H. P. Boehm, Angew. Chem., 76, 742 (1964); 78, 617 (1966); Advan. Catal., 16, 179 (1966). (6) D. Graham, J . Phys. Chem., 59, 896 (1955).
(7) R. W. Coughlin, F. S . Ezra, and R . N. Tan, J . ColEoid Interface Sci., 28, 386 (1968). (8) J. Bjerrum, “*Metal Ammine Formation in Aqueous Solution,” Haase, Copenhagen, 1941, p 285. (9) J. Bjerrum and J. P. McReynolds, Inorg. Syn., 2, 216 (1946). The JOUTM~ of Physical Chemistry, Vol. 76,No.6,1971
AKIRATOMITA -4ND PASUKATSU TAMAI
650 Table I : Rat,e Conetstnt's for Adsorption and Hydrolysis and Properties of Carbon Surface area, Carbon
HAF HAF-n3 HAF-n2 HAF-n7 HAF-nl HAF-n6 IEAF-n4 WAF-a1 BAF-01 AC-d1 AC AC-01 hC-n4 AC-n3 AC-02 ____~-____-
Modifier
mz/p
81 82 78
376 HKOB 676 "01 18% HKOs
30% HYOs 427, HNOs 60% "OD (xH4)2S208 0 2
CJIlzN2 0 2
NaOH consumption, mequiv/100 g
78 71 94 170 1120 1230 1120
3070 HNOa 60% HNOa 0 2
1100
r----kads
2.2 5.7 7.9 7.0 10.1 12.5 14.7 19.4 40.1 14.9 32.7 38.7 105 111 118
X 108, min-1-----.
30'
50'
0.069
0.735 0.682 0.790 0.69 1.04 1.20 1.74 1.64 19.3 6.27 20.8 42.3 74.2 98 83.0
0.685 2.36 4.61 7.40 9.96 8.97
khYd x 103, min-1, 30°
3.29 4.41 5.40 7.48 8.07 9.34 14.0 34.9 160
hydroxyl groups was carried out using the technique of Studebakey. et a1,'O The amount of surface acidic groups decreased as shown in Table I. Suifa'ace Area and Acidity. Nitrogen adsorption measurement? were made at liquid nitrogen temperature, and the mrface area was calculated by the method of Joyner, et The acidic groups were determined by shaking with excess standard base for an equilibration pernod of 24 h r and back-titrating with standard HCI. Adsorptivn Rate. The weighed sample of the complex was dJssolved in the required volume of solvent, and the solution was allowed to reach a constant temperature in a water bath before the weighed sample of carbon was added. A number of amber tubes (40 ml) were mounted on a shaker in a water bath incubator and were shak2n at a rate of 150 times/min. After different i n t e n d s of time they were removed rapidly and cooled in order to inhibit the further progress of the adsorption. The carbon was removed by centrifugation or filtration, and the content of [Co(NH3)6I3+ ion vias analyzed by measuring the optical absorbance at 474 mp. 7'"r rate was followed by the absorbance of small aliquots of the supernatant solution. Spectral measuremeiit was made using 1-cm silica cells in a Hitachi EPU-2 or a Gary 14 spectrophotometer. Nyclrolyas Rate. As the hydrolysis occurred only in the presence of ammonia, aqueous ammonia (1.5 N ) was used as 3. solvent unless otherwise stated. The spectrophotonirtrilc method was used to follow the rate of base hyclrolysis. Neasurements were made at 474 m.u where a relatively large absorption difference occurred 'befm-een [ C O ( K H & ] ~ +and [Co(NH3)6. (OH)I"+,
of carbon. Pretreatment of these samples is designated by an appended "n," "a," and '(0'' for oxidation by "03, (NH&S20S,and Q2, or "d" for methylation by CH2N2. The quantities of NaHC03 and Na2C03 consumed were also determined. They are not listed in this table because they were found to have a close relation to the total acidity (NaOE consumption) as reported by B ~ e h r n . ~ Upon addition of carbon to the solution of [Co(NH3)~]Cla,the yellow color due to the [ C O ( N H ~ ) ~ ] ~ + ion disappeared gradually and the solution became colorless. The fact that no other bands were observed under these conditions suggests all cobaltic ions were completely adsorbed onto carbon. The kinetic data were plotted for a first-order process, log CO/(COC) vs. t where Go (mM) js the initial concentration of the complex ion and C (mM) is the amount of the complex ion per unit volume that has disappeared in time 1. Except for the very initial region, a straight line was obtained. The concentration at the later region was represented by the equation
Results The B.E.T. surface area and the quantities of NaOH consumed are summarized in Table I for various types
(10) M. L. Studebaker, E. W. D. Nuffman, A. C . Wolfe, and L. G . Nabors, Ind. 1Sng. Chem., 48, 162 (1956). (11) L. G. Joyner, E. B. Weinberger, and @. IT. Montgomery, J. Amer. Chem. Soc., 67, 2182 (1945).
The Journa2 o j Piiysical iChemistry, Vol. 76, N o . 6,1971
c = co(l-
(1)
ae-kads')
Table 11: Activation Parameters E,,
LOEA ,
Carbon
kcal/mol
min-1
HAF AC-dl AC AC-01 AC-02
23 21 21 22 22
7.4 12.2 12.4 13.2 13.6
ADSORPTION A N D H:YDROLYSISOF [Co(NH&,I3+ON CARBON
Yi
651 b
-
Q
0.8 0.6
0.4 0.2 1000
:300
600
900
1200
Time (min 1 Figure 1. Adsorption of [Co(NHa),]3 + on carbons: initial concentration, 116mM x 10 ml; temperature, 30.0': carbons (0.01 g) are: a, AC-02; b, AC-01; c, AC; and d, AC-d1.
2000
Time ( min 1 Figure 3. Effect of HCl on the rate of adsorption on AC. The ratios, HCl/Co, are: a, 0.00; b, 0.07; c, 0.14; d, 0.35; and e, 0.70.
of the rate plot become larger. The observed rate constant, k s d s , may be represented as kads
= kl
+
[OH]
(2)
where kl = 2.30 X min-l and kz = 0.172 1 mol-1 min-l at 30.0". In order to study the mechanism of adsorption, the stoichiometry of adsorbed species was determined by analyzing NHa and C1- remaining in solution (see Table 111). The fact that NH3 is liberated during the = 20 Table 111: NEs and Cl- in Solution"
50
100
150
Carbon (mg) Figure 2. Effect of the amount of carbon on the rate of adsorption: solution, 16 mkf X 10 ml; temperature, 50.0'; carbon: a,, AC-o2; and b, AC.
where a is a coiisitant. Typical plots for AC series are shown in Figurle 1, and an apparent rate constant kads was determined from the slope of the linear part. Table I1 lists activation parameters obtained by the measurements at 30, 50, and 60". Figure 2 shows the relationship between the observed rate constant ]cads and the amount of carbon. Considering that the plots are linear and pass the original point, the raie constant can be expressed as a linear function of the amount of carbon. The effect c f pH on the rate is very significant. Excess acids inhibit the adsorption as judged from the results in Figure 3. However, upon addition of appropriate amounts of HCI (curve d in Figure 3), rate plots do not obey the above equation. The adsorption is almost inhibited at an early stage, and the rate increases gradualiy as the extent of the adsorption progresses. When the amount of NaOH added increases, both intercept (a in eq 1) and slope (/cads)
NHa in
Act mg
100 100 100 100 10 10 10 10
Time, hr
CO&dal
0.5 2.0 4.0 24.0 8.0 24.0 72.0
26.2 50.6 78.2 160.0 21.1 37.7 67.0 118.@
b
pmol
c1-/
soh, rmol
NHs/C%d.
COinitie.1
60.7 130.2 207 398 51.7 94.3 173 318
2.3 2.6 2.6 2.5 2.5 2.5 2.6 2.7
3.0 3.0 3.0 3.0 3.0
a Initial concentration of [CQ(NHS)B]CI~; 16 mill X 10 ml, temperature, 30.0'. Temperature was raised to 60". Number of adsorbed layers i~ approximately two.
adsorption was evidenced by sweeping NH3 out of the solution with a stream of air and into a solution containing Nessler reagent. The pH of the solution rose considerably, but the linearity of the rate plots was held under usual conditions. Exceptional is the case shown in Figure 3. The amount of NH3 shown in Table I11 was obtained by titrating with NC1. The amount of cobalt ion adsorbed was calculated from the disappearance of the band a t 474 mp. The ratio, NHa/Co, is always higher than two and can be regarded as fairly constant. The above fact is in contrast with the case of the adsorption of [ C U ( N H ~ ) ~on ] ~ silica + The Journal of Physical Chemistry, Vol. 76,NO. 6,1971
AKIEATOMITA AND YASTJKATSU TAMAI
652
mated from the absorbance at 474 and 503 mp by assuming the absence of any other species which have a significant contribution to the absorbance in this region. The ratio khyd/kobs varies from zero to unity with experimental conditions, especially with the concentration of NH3 and the amount of carbon. Figure 4 and Table TV show that khyd/k,bs is nearly unity in
Table IV : Competition between Adsorption and Hydrolysis AC, mg
Wavelength
( mp
1
Figure 4. Spectral cliaiige during base hydrolysis catalyzed by 16 mi44 in 1.5 N NHs aqueous AC (0.1 9 ) ; [C0(17iHa)~]~+; solution, 10 nil; lempei*ature, 30.0"; times from mixing with catalyst are: a, 0 ; h, 2'; c, 4; d, 8 ; e, 14; f, 35, 168, and 1200 miu.
gel reported by SmIth, et a1.,12who observed that the NH3/Cu ratio in adsorbed species is higher than that of primary complex species initially in solution. Chloride ion was determined volumetrically by Mohr's method. 41 was noteworthy that no C1- ions in [Co(NN3)s]C13 W C T ~adsorbed on carbon but remained in solution. Figure 4 shows the spectral change during the catalytic base hydrolysis of the [Co(KH3)6l3+ion into the [CO(NH~)~(OH) l2 + ion. The absorption maxima of these cornplexcs are at 474 and 503 mp, and it has been bacce at these wavelengths changes iitration. The change of absorbance at 474 mp was pIotted in the pseudo-first-order expression tigainst lime, and the observed rate constant, k o b s , W P S obtained. Generally speaking, the decrease in the absorption at 4'74 mp in this system is ue to convei*sion t o the hydroxo complex and/or irreversible adaorption onto the carbon. The method of judgement for these alternatives is very simple : the absorbaiiec at 503 mp decreases at the same rate as thaf at 474 1np if only irreversible adsorption occurs, whereas if only hydi-olysis occurs, the sum of concentrations calculated fiwm these two bands must remain constant. S'hc apparent rate constant must be the sum of constants for adsorption, kads, and for hydrolysis, khyli, becautse /Gobs and k a d s are the firsborder rate constants with respect l o the concentration of [CO(IVH~) 13 -i-. kots
== k i d s $. khyd
1.5 0.15 0.060
16 10 16 16 10
0.00
0.15 0.00
16
7488,'
min-1
khya/kobsb
160 4.22 4.53 2.36 0.58 0.10
0.38 0.02 0.0 -0.0 0.0
min
6 60 300
1.0
d
-2000 d
'
[ C O ( N N ~ ) ~10 ] ~mI. +, Time required k h y d = hobs - IC,?,. for the peak shift to 488 mp. No shift was observed.
1.5 N EH3solution. This ratio decreases with decreasing concentration of NH3 and the amount of carbon. I n order to avoid this complexity, the concentration of NHa was made 1.5 N for the comparison of the catalytic activity of carbons. It, is assumed that the value khyd can be used as an indication of the catalytic activity of carbon. These rate constants are summarized in Table 1 together with other properties of carbons.
Discussion It is evident from Table I that the adsorption rate constants are closely related to surface acidities rather than surface areas. I n view of the correlations with the total acidity on the carbon surface (Figures 5 and 6), it seems likely that the adsorption may be initiated, for example, in the following manner.
+
~[CO(NH&]~+--d
i2NH,
(4)
(3)
The amount of disappeared cobalt ions, which is necessary fop. the calculation of /Gads, was determined from the concentrations of [Co(NH&I3+ and [Co(NH3),(OH)12+ ions in solution. These concentrations could be estiThe Journal of .Phwical Chemistry, 1'01. 76,No. 6,19ri
100 100 100 100 10 10
hobs X 108,
C O ~ + , ~"a, miM N
(12) G. W. Smith and H. W. Jacobson, J , Phus. Chem., 60, 1008 (1956).
ADSORPTIOE A N D HYDROLYSIS OF [ C O ( ~ ” ~ ON ) ~ CARBON ]~+
Total acidity ( rneq/Ioog 1 Figure 5 . Ilepei dence of observed rate constants for adsorption o n total acidity of AC series: temperatures are: a, 64.9. b, 50.0; 2nd e, 30.0”.
i/’
Y
I
E
1-
E
[Co(WH3)6I3+, would interact with these negatively charged sites more readily than the undissociated neutral sites. (2) Decrease of these negatively charged sites, on the contrary, retarded the adsorption. These groups may be blocked by methylation with CH2K2,and the absorption onto AC-dl was much slower than that onto AC as was expected. Protonation of these groups also retards the adsorption as illustrated in plot e and an induction period as in plot d in Figure 3. The induction period was followed by an increase of rate due to the neutralization of HC1 with the librated NHY. (3) Contrary to the present case, the reaction rate of complex anions, such as [Co edtaI2- or [Co(]1\’O&l3--, varied inversely as the surface activity of carbon. A detailed account of these reactions will be described in subsequent communications. The absorption of the complex ion onto carbon (eq 4) facilitates the subsequent reactions. More NH3 molecules were substituted by JTater and/or hydroxide ion according to eq 5, unless NH3 was added to the solution in excess. Table I11 indicates that the amount of KH3 in solution is 2.5 times the amount of complex ion adsorbed. Considering that a part of NE3 molecules does not dissolve in solution but adsorbs on carbon, at least three molecules of NH, seem to be replaced ] ~ upon + adsorption. Consefrom one [ C O ( X H ~ ) ~ion quently, the value (m n ) in eq 5 must be 2 or above.
+
b
u
653
+
mH,O
+
nOH-
-
%
x
IO
15
‘Totall acidity ( meq/ IOOg) Figure 6. Deperdence of observed rate constants on total acidity of HAF wries: a, adsorption at 50.0’; b, hydrolysis et 30.0”.
If such groups on a (carbon surface are the active sites for adsorption, it is ;supposed that the electrostatic nature of’ the surface and the solute would be one of the most importanr factors. Evidences in favor of this expectation were obtained from the following experiments. (1) The rate of adsorption increased upon addition of XaOH. Ir. the higher pH region, -COOH and -OH group:s on the carbon surface dissociate into -COO- and -0 - anionic groups. Tho complex cation,
When the amount of a complex ion was sufficiently large compared with carbon, the multilayer formation was observed after a long time for zdsorption. One example is shown at the bottom of Table 111. A possible explanation for this phenomenon may include the formation of a polynuclear cobalt complex involving two hydroxide ion bridges. h4any examples for this type of polynuclear complexes have been reported.I3 Table I11 demonstrates that NH3 moleeules dissociate from the complex ion in the second layer, too. This fact can be reasonably interpreted by the bridge formation. This model is, however, highly speculative, and further work Tvould be required to explain fully this multilayer formation. A close relationship between the rate of hydrolysis and the total acidity was also observed and is shown (13) “Gmelins Handbuch der Anorganischen Chemie,” Kobalt, 58, B-2, Verlag Chemie, GmbH, Weinheim, 1964, pp 621-633. The Journal of”Physical Chemistrg, V a l . 7 5 , X o . 5 , 1971
MASAOYAFUSOAND MICHAELE, GREEN
654 active in Figure sites6for (plot hydrolysis bl). It are may also bethe concluded acidic functional that the groups on the carbon surface, and the initiation step for hydrolysis is the same as that described in eq 4. Further substitution of NHs like eq 5 might be depressed in bhc present case, because the concentrated aqueous ammonia was used as a solvent. The hydroxide ion ahtacked the Co-0 bond rather than the Co-N bond, and the product, [Co(NH&(OH) 12+, was obtained IBS shown in the following equation.
~ + W W d 5 1 2 +
+
OH-
-t
6G(~H&(OH)]2* ( 6 )
Noise Spectra Associated with Hydrochloric Acid Transport through Some
Cation-Exchange Membranes1 by Masao Yafuso and Michael E. Green* The City College of The City University of New York, New York, New Yorlc 10031 (Received August 7, 1970) Publication costs assisted by The City College of the City Universitu of New York
The noise spectra and some other electrical phenomena associated with transport of hydrogen ion across a cation-exchange membrane have been studied a t currents great enough to cause the formation of a depletion layer. Under these conditions, one obtains noise spectra which show a frequency dependence of approximately u - ~a t low frequencies and a-4 a t high frequencies. Combined with data on voltage drops as a function of time and distance from a membrane, we find the principal noise source to be within 40 pm of the membrane surface. The frequency a t which the change in slope occurs was measured as a function of current, temperature, and concentration. The total current can be split into a diffusion contribution and a contribution from 8’ and OH- ions from dissociation of water. The corresponding electric field for dissociatioii is calculated from the Onsager equation. The frequency a t which the change of slope occurs, taken a t constant field, 1s B function of temperature, with an activation energy of 16.4 & 0.9 kcal/mol, and the dissociation flux is proportional to this frequency. Possible sources for the noise are considered.
Introduction I n a previous paper2 we reported some results on the noise spectra associated with the transport of various ions t h sough several cation- and anion-exchange membranes. ‘Et was not possible a t that time to give a mechanism for the generation of the noise, except to note that it was associated with the formation of the depletion 1a:yer (the layer formed in front of the membrane when the current exceeds the maximum current density which the ions in solution can maintair~).~ The critical current density has been measured by many authors, and from this the thickness of the diffusion layer h,as been estimated and descriptions of the associated phenomena given in terms of diffusional fluxes.* However, a microscopic model for these processes does not appear to be available a t the present The Journal of Physical Chemistry, Vol. Y 6 , No. 6,1971
time. Gregor and Miller5 did attempt t o estimate the field at which water dissociation occurs from a kinetic argument, but their model seems, at least, incomplete. By studying noise spectra, one can get kinetic data not limited to the slowest step of the overall process, and this has been applied in various systems of semi(1) Based upon a dissertation submitted to The City University of New York by M. Y. in partial fulfillment of the requirements for the degree of Doctor of Philosophy. (2) (a) M. E. Green and M. Yafuso, J . Phys. Chem., 72,4072 (1968) ; (b) M. E. Green and M.Yafuso, ibid., 73, 1626 (1969). (3) F. Helferich, “Ion Exchange,” McGraw-Hill, New York, PI’. Y., 1962, p 399. (4) (a) A. M. Peers, Discuss. Faradau SOC.,21, 124 (1956); (b) K, S.Spiegler, U. S.Office Saline Water Res. Develop. Prog. Rep. 1968. (5) 1%. Gregor and I. F. Miller, J. Amer. Chem. Soc., 86, 5689 (1964).
