untitled Kinetics assisted Amel Mous Amal Abd‐E Chemistry Depart *Corresponding a Tel.: +2.03.010605 ARTICLE INFO Received: 04 Janu Received in revis Accepted: 09 Mar Online: 30 June 2 KEYWORDS Ion‐pairing Ethane‐1,2‐diol Solvent effect Thermodynamics Kinetics of aquati Bromopentammi 1. Introductio Despite th significant in kinetic nature media. Jones et coefficient, ko univalent liga while that of th Amira et a the rate of a perchlorate an in only one pe of some dicarb The presen the aquation presence of io and in binary 50%) at differ highlighting a and extra ther Besides the Li relations) has ion‐pair aquat interest due t and ion solvat and practical i 2. Experimen Succinic ac diol (B.D.H. and mech by ion‐pa stafa Ismail* El‐Hamid Za tment, Faculty of Sc uthor at: Chemistry 531801; fax: +2.03. ORMATION uary 2012 ed form: 07 March rch 2012 012 s ion ine cobalt(III) on he observation solution, there e of these ion al. [1] found bs with the sti and is linear fo he divalent dica al. [2‐8] have d quation of chl nd chloropenta rcent (10 or 20 boxylate solutio nt work is focu of bromope on‐pairing suc y mixtures of w ent temperatur any new corre rmodynamic an near Free Ener been examined tion reaction. to their reactiv tion effects wh importance. ntal cid (B.D.H. Anal Analar) was Eu ISSN 2153‐ Europ J hanism of airing suc *, Seleim Mo aghloul and cience, Alexandria U y Department, Facu 5932488. E‐mail a h 2012 that ion‐assoc e are relatively ‐paired entitie that the corre iochiometric co or the aquatio arboxylate ligan escribed the io loro, bromopen amminechromi 0%) mixed aque ons. sed on the kine ntamminecoba cinate anion in water and eth res. This sort of lations based nalysis of the ob rgy Relationship d to show the ro Succinate solut vity spectrum hich leads to im lar) was recrys redistilled, an uropean Journal Europe 2249 (Print) / IS DOI:10.5155 pean Jo Journal home f aquation cinate an ohamed Sele Mohamed F University, Ibrahim ulty of Science, Alex ddress: amelmostaf ABSTRACT The aquation k presence of diff ethane‐1,2‐diol succinate ion c different exper calculated and the obtained re interactions. An the ion‐pairing is proposed. ciation is kinet few studies on s in mixed so elation of the oncentration o on of [Cr(NH3) nds is non linea on‐pairing effec ntamminecobal um(III) perchl eous organic so etic solvent effe alt(III) ion, in n aqueous me hane‐1,2‐diol (u f study will assi on thermodyn btained kinetic p (and other sim ole of solvent o tions are of sp toward ion‐pa mportant theore tallized. Ethane nd mixtures of Chemistry 3 ( ean Journal of Ch SSN 2153‐2257 5/eurjchem.3.2. ournal o epage: www.e n of brom ion in eth eim, Fawzy Amira mia, Alexandria, 213 xandria University, afa@yahoo.com (A.M kinetics of brom ferent concentra l (up to 50%, w concentration (L rimental conditi discussed in ter esults have been n empirical corr succinate ligand ically n the olvent rate of the 5Cl]+2 ar. cts on lt(III) orate olvent cts of n the dium up to st for namic data. milar on the pecial airing etical e‐1,2‐ were prep com conv (B.D 2.1. S and conn /KN liber pote due rate metr 3. Re A amm (80% cont w/w squa data whe V∞ i initi temp (2) (2012) 196‐2 hemistry (Online)  2012 196‐201.583 of Chem eurjchem.com opentam hane‐2‐di a 321, Egypt Ibrahimia, Alexan M. Ismail). mopentammine ations of succina /w) within the L) and the obse ons. The therm ms of solvent ef n discussed on t relation between d (L) has been es pared using do mplex was prep verted to the n D.H.) was dried a Equipments an Spot Galvanom sensitivity of nected in seri O3/Buffer, H2Q rated bromid entiometerically to its high rep constant val rically [2]. esults and disc A series of kin minecobalt(III) % neutralizatio taining differen w) within the te are method wa a according to th ere Vt is the pot is the infinity t al concentrati perature and s 201 2 EURJCHEM mistry m mine cob ol‐water dria, 21321, Egypt. cobalt(III) ion ate ion in aqueou temperature ran erved rate cons odynamic param ffects. Also the e the basis of solu n the rate coeff stablished and a ouble distilled ared by the m nitrate salt and at 300 oC for 3 h nd kinetic mea meter with an in f 2x10‐8 A. p es with the c Q, Q, Pt. For the de ion from y [8]. This met producibility an lues with tho cussion netic runs for ion were per on of succinic nt percentages o emperature ran as used to fit th he kinetic Equa tentiometric tit titre which wa on of the co solvent compo alt(III) ca mixtures . has been inve us medium and nge (35‐65 oC). stant have been meters of activa xtrathermodyna ute‐solvent and ficient and the c a suggested reac water. The [Co method of Booth d recrystallised hr. asurements nternal resistan er mm. scale chemical cell, e purpose of d m the aquat thod of analys nd the good ag ose obtained the aquation o rformed in su c acid (m1) by of ethane‐1,2‐d nge (35‐65 oC) he obtained firs ation 1, ration reading s calculated fro omplex salt (m osition, a set o ation estigated in the that mixed with The ion‐pairing n determined at ation have been amic analyses of solvent‐solvent concentration of tion mechanism o(NH3)5Br] Br2 h [9] and then d form. Na2CO3 nce of 60 Ohms division was Ag/AgBr(s)/Br‐ etermining the tion reaction is is preferred greement of its spectrophoto‐ of bromopenta uccinate buffer y Na2CO3 (m2) iol (up to 50%, . A linear least st order kinetic (1) at time (t) and om the known m3). For each of kinetic runs e h g t n f t f m 2 n 3 s s ‐ e n d s ‐ a r ) , t c d n h s Ismail et al. / European Journal of Chemistry 3 (2) (2012) 196‐201 197 were conducted for different stoichiometric concentrations of the dicarboxylate ligand. The obtained least square rate constant kobs, at different experimental conditions are listed in Table 1. Garrick [10] found that the rate of aquation of chloropenta amminecobalt(III) ion was slightly increased by chloride, nitrate, chlorate, formate and acetate anions while sulphate cause marked acceleration. Garrick expressed his results in terms of, Equation 2. (2) where, kobs is the observed pseudo‐unimolecular rate constant, k1 its value in the absence of L, k2 is the catalytic coefficient and CL is the stiochiometric concentration of the univalent ligand. Jones, Harris and Wallace [11] on testing the last equation found that the plots of kobs against CL of univalent ligands were linear and extrapolated to a common k1 at CL = 0 while for divalent ligands the plots are nonlinear. In the present investigation, it is now necessary to examine the possible empirical correlations in the succinate ion solutions containing different concentrations of ethylene glycol (up to 50%) where the added ethylene glycol lowered the dielectric constant which assist the ion‐association process. The stoichiometric concentration of the succinate ion can be expressed by the stoichiometric cabroxylic acid concentration (m1). Firstly most plots of kobs versus m1 for different succinate ion at different solvent compositions gave nonlinear plots radiating from k1 (rate constant in absence of succinate ion). This nonlinear correlation is agreed with the work of Jones, Harris and Wallace [11] for divalent ion‐pairing ligands (like dicarboxylates). These plots are refined by using the free succinate ion concentration [L] instead of m1which is reprehensive in Figure 1. Figure 1. Variation of kobs versus [L] for succinate ion at 45 °C for 50% solvent composition. Different attempts were tested and one of them was only valid in which the plots of 1/(kobs–k1) versus 1/[L] (L is L2‐) give a straight line with a positive slope and positive intercept, as shown in Figure 2. This linearity leads to the following empirical correlation (Equation 3). ]L[ b a )kk( 1 1obs   (3) where a and b are empirical constants and their values for different cases are listed in Table 2. Equation 3 can take another form, Equation 4.   ]L[)ba(1 ]L[b1)ba(kk k 11 obs    (4) Figure 2. Plot of 1 / (kobs – k1) versus [L] for succinate ion at 45 oC for 50% solvent composition. Data in Table 2, shows that values of the empirical constants a and b depend on the solvent composition which arises from the solute‐solvent and the solvent‐solvent interactions. In presence of succinte ions, an ion‐pair is formed and the aquation reaction can be represented as: CpX2+ + H2O [CpXH2O]2+ k1 (5) (6) where, is the ion‐pair concentration, is the free complex ion concentration and kobs is related to kip according to Wyatt and Davies treatment [12]. The principle of calculations of ion‐pair rate constant kip, were discussed before [2]. K1, K2 and KD values taken from references [13,14].The calculation of the free ion‐pairing succinate concentration [L] is based on the following set of equations, and Table 3 collects the average kip values at different conditions. CpXL  CpX2+ + L2‐ KD (7) NaL‐  Na+ + L2‐ KNaL‐ (8) H2L  HL‐ + H+ K1 (9) HL‐  L2‐ + H+ K2 (10) KD = [CpX2+][L2‐] f22/ [CpXL] (11) K1 = [H+][HL‐] f12 / [H2L] (12) K2 = [H+][L2‐] f2/ [HL‐] (13) KNaL‐‐ = [Na+][L2‐] f2/ [NaL‐] (14) where f1 and f2 are the activity coefficients of the univalent and divalent ions respectively. Their values were obtained from Debye‐Huckel equation in the form log f = ‐A (I½/1+I.3I½)‐0.3I where I is the ionic strength, 198 Ismail et al. / European Journal of Chemistry 3 (2) (2012) 196‐201 Table 1. Values of rate constant kobs, for the aquation of [Co(NH3)5 Br]2+ in succinate media at different composition of ethylene glycol. % 35 oC 45 oC 55 oC 65 oC 103 [L] kobs* x107 103 [L] kobs * x107 103 [L] kobs * x107 103 [L] kobs * x107 0 5.39 236.5 5.26 679.8 4.78 2463.9 5.28 9113.7 8.09 245.8 8.19 774.1 7.31 2489.2 7.97 7081.7 10.84 250.0 10.72 784.4 9.73 2527.9 10.76 9063.9 13.57 255.0 13.48 806.6 12.20 2698.4 13.47 11845.8 16.20 257.3 16.14 848.4 14.64 2629.8 16.11 11602.4 18.81 263.1 18.66 885.1 17.09 2925.4 18.76 12195.5 10 5.34 246.2 5.33 702.7 5.26 2329.2 5.31 10713.5 8.10 259.4 8.07 762.8 8.02 2567.6 8.056 10293.1 10.79 268.1 10.77 733.8 10.76 2647.4 10.83 9583.9 13.53 275.5 13.46 813.3 13.48 2656.7 13.56 11534.6 16.21 280.8 16.06 793.5 16.07 2728.6 16.17 11327.6 18.77 285.1 18.77 839.9 18.66 2816.0 18.95 11412.2 20 5.34 178.0 5.19 475.0 5.25 2551.2 5.23 8267.4 8.13 173.5 8.05 500.7 8.05 2909.7 8.00 10725.0 10.91 190.0 10.83 526.8 10.80 2848.7 10.79 9498.6 13.65 194.5 13.48 577.1 13.50 2581.6 13.51 10368.4 16.27 198.3 16.20 612.2 16.19 2919.6 16.22 13870.6 18.99 187.4 18.81 650.9 18.82 2402.3 18.87 11777.3 30 5.24 180.4 5.18 625.4 5.17 2382.0 5.16 7660.3 8.03 188.8 7.94 704.5 7.99 2566.3 7.96 8782.6 10.84 193.8 10.76 688.8 10.79 2681.2 10.76 8420.1 13.58 196.8 13.56 742.6 13.55 2780.0 13.52 9623.7 16.34 202.1 16.31 731.8 16.27 2682.8 16.28 9874.5 18.98 203.3 18.85 762.6 19.02 2792.1 19.02 9299.4 40 5.17 186.7 5.27 612.5 5.09 2188.7 5.04 8008.4 7.99 220.0 8.05 622.6 7.90 2426.3 7.78 9165.7 10.81 196.1 10.84 667.8 10.72 2407.6 10.65 9698.0 13.62 199.8 13.67 696.7 13.36 2382.2 13.43 7498.0 16.34 215.0 16.39 720.4 16.13 2351.3 16.17 10538.5 19.07 204.6 19.08 698.2 18.91 2468.0 18.98 8680.1 50 5.14 211.0 5.15 544.3 4.94 2095.3 5.00 6831.5 7.97 539.6 7.90 602.5 7.74 1651.3 7.79 7963.1 10.80 692.8 10.81 650.5 10.67 2416.9 10.71 7810.8 13.65 285.0 13.63 661.1 13.38 2464.0 13.44 7600.1 16.39 438.0 16.43 654.2 16.24 2559.6 16.31 8064.5 19.15 356.0 19.15 705.7 18.97 2661.8 19.09 8476.9 Table 2. Values of (a, b and a/b) (a = Intercept, b = Slope) of equation 1 for succinate ion at different composition of ethylene glycol. W/W, %a Succinate ion a b a/b 0 ‐8.7453 0.7465 ‐11.715 10 23.1520 0.1045 221.550 20 8.7441 0.6695 13.061 30 19.8020 0.0699 283.290 40 21.8470 0.0661 330.514 50 19.6120 0.0999 196.316 a w/w, % of ethan‐1,2 diol. Table 3. Average values of ion‐pair rate constant 105 kip (s‐1) at different experimental conditions for succinate ion. W/W, %a Temp. °C 35 45 55 65 0 3.4 10.9 34.8 179.7 10 4.2 12.7 38.7 180.9 20 2.6 8.4 44.6 146.6 30 2.7 11.3 39.4 131.4 40 2.9 10.2 33.4 116.3 50 4.6 9.3 35.3 110.3 a w/w, % of ethan‐1,2 diol. I = 0.5 ([H+] + [HL‐] + 4 [L2‐] + 4 [CpX2+] + 2m3 + [Na+] + [NaL‐]) (15) m1 = [H2L] + [HL‐] + [L2‐] + [CpXL] + [NaL‐] (16) m3 = [CpX2+] + [CpXL] (17) where m3 is the stoichiometric concentration of the complex ion, (m3 = [CpX2+] + [CpXL]), k1 is the observed rate constant in the absence of the succinate ion‐pairing ligand (X = Br). The principle of calculations performed by computer programs can be summarized as. For the first cycle [H+] = 0, [CpXL] = 0, [NaL‐] = 0 (18) [CpX2+] = m3 ‐ [CpXL], [HL‐] = 0.5 m2 [H2L] = 0.3 m1 (19) [L2‐] = m1 ‐ [HL‐] ‐ [CpXL] ‐ [NaL‐] ‐ [H2L] (20) [Na+] = 2 m2 ‐ [NaL‐] (21) Then the ionic strength takes its first approximated value and then f1 and f2 (f2 = 4 f1) after which the following terms take their new values as, [H+] = K2 [HL‐] / [L2‐] f2 (22) [H2L] = [HL‐] [H+] f 21 / K1 (23) [HL‐] = 2 m1 ‐ 2m2 ‐ 2[H2L] ‐ [H+] (24) Ismail et al. / European Journal of Chemistry 3 (2) (2012) 196‐201 199 [L2‐] = m1 ‐ [HL‐] ‐ [H2L] ‐[CpXL] ‐[NaL‐] (25) [NaL‐] = [Na+][ L2‐] f2 / K NaL‐ (26) [CpXL] = m3 / [(KD /f22 [L2‐]) +1] (27) [CpX2+] = m3 ‐ [CpXL] (28) and then I, f1 and f2 recalculated again. These steps of calculations were repeated many times until the difference between two successive values of [L] becomes equal to (or less than) 10‐7. 3.1. Variation of activation parameters with solvent composition The thermodynamic parameters of the ion‐pairing aquation reactions, ∆G*ip, ∆H*ip and ∆S*ip were calculated at 25 oC using least square procedure program, and these values with their standard deviations are given in Table 4. ∆G*ip values show a small increase with increase the mole fraction of the co‐solvent, giving a good indication of the compensation between ∆H*ip and ∆S*ip .Variation of ∆H*ip and ∆S*ip versus the mole fraction of the co‐solvent (χ2) displayed minimum at χ2 = 0.03 and maximum at χ2 = 0.07 as shown in Figure 3. These values are found to be close to those obtained of other cobalt complexes in the same solvent system. In dilute aqueous solution (χ2 < 0.1) ethan‐1,2 diol can exist as the gauche cyclic conformer with strong intermolecular hydrogen bonding, which enhances the hydrophobicity of the co‐solvent [15]. The enhancement of the solvent structure and structure breaking, as reflected by the maxima and minima in these plots, strongly suggest that the effect of structure perturbations in the bulk phase are effectively transmitted to the reaction zone through the solvent shell of the reactant. Table 4. Values of the thermodynamic parameters ∗ , ∗ , ∗ of different solvent compositions at 25 oC. W/W, %a ∗ (Kj/mol) ∗ (J/K.mol) ∗ (kJ/mol) 0 110.38±9.21 26.38±18.57 102.51±17.73 10 104.78±8.45 10.03±16.21 101.79±16.26 20 116.38±6.55 44.02±20.31 103.26±12.61 30 108.56±1.22 20.04±3.78 102.59±2.35 40 103.69±2.06 6.40±4.22 102.43±3.97 50 91.35±9.31 ‐33.22±18.87 101.23±17.92 a W/W % of ethan‐1,2 diol. Figure 3. Plot of (∆G*ip, ∆H*ip, ∆S*ip) versus X2. 3.2. Extrathermodynamic analysis In the present work the plot of ∆H*ip versus ∆S*ip for the ion‐ pair aquation reactions at different compositions is linear as shown in Figure 4. The parallel changes in ∆H*ip and ∆S*ip lead to only small changes in ∆G*ip and for such a closely related series, a common reaction mechanism is supported. The obtained isokinetic temperature (β) is 322.4 K. Figure 4. Isokinetic plot of the aquation of [Co(NH3)5Br]2+ at different solvent composition. The genuine nature of the isokinetic relationship was verified by the Exner criterion [16] by plotting log k(328 K) versus log k(318 K) . The value of β was calculated from Equation 29, where b is the slope of Exner plot, and the ratio T1/ T2 must be smaller than unity. (29) The calculated value of β is 330 K, which lie within the studied temperature range. This means that the compensation effect must be born in mind. The true explanation of compensation effect lie in terms of solvent‐solute interactions. Any effect that, for example, leads to stronger binding between a solute molecule and the solvent molecules will lower the enthalpy; it will also, by restricting the freedom of vibration and of rotation of the solvent molecules, lower the entropy. Application of more exact theories to these effects leads to the result that they generally give rise to a fairly exact compensation between ∆H*ip and T∆S*ip and therefore to a very small effect on ∆G*ip. Although the effect of solvent on the rate and the position of chemical equilibrium has been known, there is still no reliable or exact methods for a quantitative description and prediction of such solvent effects. In the present work the correlations ∆G*ip‐∆Goass, ∆H*ip‐∆Hoass and ∆S*ip‐ ∆Soass were found to be linear at different solvent compositions (0‐ 50%w/w) and can be represented by the following Equation 30. y = mx + c (30) in which y represents ∆G*ip, ∆H*ip or ∆S*ip and x represents ∆Goass, ∆Hoass or ∆Soass where these correlations can describe a Linear Free Energy , Linear Enthalpy, and Linear Entropy Relationships . The m and c values were found to be 70.44 and 109, ‐1.23 and 116, ‐1.04 and 86.1 for the above relations respectively. These linear correlations can refer to a common reaction mechanism existing within the studied mixed solvent compositions. 3.3. Variation of ion‐pair rate constant with dielectric constant and water concentration Elsemongy and Amira [17] proposed a general equation for the variation of specific rate constant (k) with the dielectric 200 Ismail et al. / European Journal of Chemistry 3 (2) (2012) 196‐201 constant (D) for any reaction, in which the transition state may or may not be polarized. This equation was proved to be applicable to usual reactions, as well as those exhibiting minima or maxima with solvent composition variation. Their equation takes the form, Equation 31), . . (31) (Log C = Log a + 293.15 b) (32) where A is the frequency factor, a and b are Akerlöf`s empirical constants [18]. The log A/k versus Eb/log (C/D) gives a good straight line passing through origin with a slope equal to 0.0520, which is in consistent with the theoretical one (0.0522). This finding add, a further support to this equation. Also log kip was plotted against reciprocal of the dielectric constant D at different temperatures, where the dielectric constant values of different compositions are obtained from Akerlöf data [18]. As shown in Figure 5, the plot was found to be non linear in accordance with the general observations found in the aquation of a large number of other cobalt(III) complexes [19‐25] in water‐cosolvent mixtures, which led to a conclusion, that the contribution of the non electrostatic part of solvent effect, overcomes the electrostatic component part. Also, this parameter measures macroscopic properties, while specific solute‐solvent interactions occur on a microscopic scale are completely neglected. In such cases, the differential solvation of the initial and transition states is the controlling factor for the changes in the rate constant [26]. Figure 5. Variation of log kip versus 1/D at 35 °C. The plot of the logarithm of ion pair rate constant versus the logarithm of water concentration at constant temperature and varying solvent composition is found to be nonlinear. This nonlinearity can be attributed the complex structure of the mixed solvent medium. 3.4. Proposed reaction mechanism A proposed reaction mechanism which takes into considerations, the Wyatt and Davies treatment [12], the extrathermodynamic analysis of the obtained kinetic data, the kinetic solvent effects and the empirical correlation of kobs with the free concentration of the ion‐pairing succinate ligand is described by the following scheme. (33) (34) (35) and are five coordinate intermediates (kinetic solvent effect) (36) (37) where . (Watt and Davies treatment) and (38) i.e. . (39) Therefore . k k (40) Applying steady state treatment on and [ (41) k (42) or [ ] ) (43) i.e. (44) Therefore, (45) From Equations 41 and 45 (46) Substituting in Equation 40, and rearrangement Therefore, (47) Equation 47 discuss the empirical correlation between kobs and [L2‐] in which k1 in Equation 4 is replaced by k4 in Equation 47, where and 1 (48) and consequently the empirical constant (a) equal to . Ismail et al. / European Journal of Chemistry 3 (2) (2012) 196‐201 201 4. Conclusion Our aim in the present investigation was to look for possible correlations between the thermodynamic properties of the activated complex and the corresponding thermodynamic functions of the ion‐pair formation reactions within the studied binary composition of the mixtures under investigation. Further we investigated the kinetic solvent effects and the empirical correlation of kobs with the free concentration of the ion‐pairing succinate. 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