Nikel (II) ve demir (III) nitrit komplekslerinin sulu ortamdaki ardışık oluşumunun potansiyometrik ve spektrofotometrik yöntemle incelenmesi

Abstract

Ill ÖZET Bu çalışmada, geçiş metal iyonlarından nikel (II) ve de mir (I II) katyonlarımın nitrit iyonu ile verdiği komplekslerin sulu ortamdaki ardışık dengeleri araştırılmş ve (25 ± 0.1) C da ve birim iyonik şiddetteki perkloratlı bir ortama oluşum sa bitleri saptanmıştır. Çalışmada hem potansiyometrik ve harı de spe>rrrofotometrik yöntemler uygulanarak kompleks oluşum başarmakları araştırılmış tır. Sulu ortamlarda kompleks oluşumuna aix gerekli kuramsal bilgiler ayrıca sunulmuştur. Kullanılan derişiklik aralığında, nikel(II) katyonu nit rit ligandı ile tek çekirdekli iki zayıf kompleks oluşturmakta dır. Demir(III) iyonu ise nitrit ligandı ile üç tane daha kuvvet li tek çekirdekli kompleksler vermektedir. Nikel (II) -nitrit sis teminde oluşan iki kompleksin oluşum sabirleri için sırasıyla, -1 -2 3X = (6 ± 1) M ; B2 = (12 ± 2) M değerleri bulunmuştur, öte taraftan demir ( III) -nitrit sisteminin tik:, oluşum sabitleri içinj gj - (3.9 ± 0.7)102 M`1 ; 02 ~ (5 ± 2)103 M*2 ; g3 - (2.8 ± 0.6)105 M~3 değerleri saptannıştır. Yanılgılar % 99.9 güvenirlik sınırlarına göre verilmiştir.

IV SYNOPSIS POTEKTIOMETRIC AND SPECTROPHOTOMCTRÎC STUDIES ON THE STEPWISE FORMATION OF COMPLEXES OF NÎCKEL(II) AND ERON(III) NÎTRÎTES IN AQUEOUS SOLUTION The most important way to characterize the complex for mation in solution is to determine the equilibrium constants of complexes formed. The majority of quantitative equilibrium measurements have been carried out in aqueous solution because of enhanced solubilities and, moreover,- more is known about how ions behave in this solvent. In recent years, many differ ent techniques and the new methods of computation to evaluate the stability constants of complexes have been developed. After a large amount of stability constant data had been determined, many chemists began to examine these data in order to make a generalization on metal-ligand association. For this purpose 5 some reliable classification schemes have been proposed. Ac cording to the most important classification, acceptors and donors are classified as hard (or class a) and soft (or class b), but there are many cases where acceptors and donors fall on the borderline between two classes. This classification enables us to understand why a particular metal ion shows a preference for one ligand rather than another. In aqueous solution, the formation of complexes between hard, or class a, acceptors and hard donors is usually entropy- controlled. In other words, the decrease of the free energy is mostly due to a large gain of entropy. As a rule» the enthalpy change is often positive for this interaction. Hard acceptors, as well as hard donors, are characterized by a high charge//radius ratio. The bonds formed between them are mainly elec trostatic in nature. In aqueous medium they also interact strongly with water molecules, forming ordered hydrate struc tures. In a complex formation, these structures are broken. This causes a large gain in entropy change. For soft, or class b, acceptors and soft donors the situation is almost reversed. Tne complex formation between soft acceptors and soft donors is invariably enthalpy - controlled i.e. the decrease of the free energy is mainly due to a large decrease of enthalpy while the entropy term is either minor important, or fairly strongly negative, thus counteracting the reaction. The large negative values of enthalpy changes are due to the formation of essentially covalent bonds. For almost any d configuration the stereochemistry of ligand ion around the metal can be predicted, and the relative stability of various d structures is known. Thus the majority 3 5 fi of d, d, and d complexes, such on Cr(III), Fe(III), Fe(II) Q and Co(III) are octahedral. The d systems viz. Pd(II), Pt(II) Q and Ir(I) are typically square planar. As a d acceptor Ni(II) exhibits six-coordination too. The variable valency of accep tors plays an important role in many complex reactions. The nitrite ion, N0«, is an ambidendate ligand which can be bonded to a metal ion either via N atom or over 0 atom. Thus forming the nitro and nitrito complexes, respectively. The metal ion itself has a dominant influence on this bonding. The class a or hard acceptors prefer the 0 atom while the soft acceptors prefer the N atom. Nitrogen could therefore be clas sified as considerably softer donor atom than oxygen. Conse quently, bonds formed via the N atom are more covalent inVI nature compared the bonds formed via the 0 atom. In this study, the stepwise complex formation of nitrite ion with iron (III) and nickel (II) has been investigated, both potentiometrically and spectrophotometrically, in an aqueous medium of unit ionic strength with sodium perchlorate as a supplementary salt at 25.0 ± 0.1 C. The potentiometric measurements have been carried out by means of a glass electrode in different buffer solutions. On account of the low stability of nickel (II) nitrite complexes, their formation could be most accurately studied by potenti ometric determination of the free central ion concentration. Since for this system, however, no suitable working electrode is available, the free ligand concentration, [L], has been determined by glass electrode measurements in nitrite-nitrous acid buffers containing nickel (II) perchlorate. This can, in principle, be performed for the numerous ligands which are.anions of weak acids, and constitutes a very general method for the determination of the free ligand concentration. These measurements are, on the other hand, well suited for the nitrite system of iron(III) as stronger complexes are formed in this system. In order to evaluate the values of formation constants from these measurements, the acidity constant of the nitrous acid has to be known. It has therefore been determined in separate measurements. Values of [H J have been found by measuring the e.m.f.s of cells of the following type. tG.E/C^ M Me(C10u)v,CH H HCIO^C^ M NaN02, NaClO^ to-1 M // RE-VII where RE// - Ag,AgCl/0.025 M NaCl, 0.975 M NaClO^/1.0 M NaClO^/. Here, Me denotes nickel or iron and v is the valency. The measurements have been arranged as titrations at constant metal ion concentrations, C,, and at constant hydrogen ion concentration, C`. Equal volumes of the solutions T, and T` were added from piston burettes to v ml of the solution S. These solutions had the following compositions : XL. M NİCCİO.). rl. H l S :4CU M HC1CX n h 1(1.0 - 30J M NaCIO 1.00 M NaNO, 2^ M NKCIO^ T2 H2C^ M HC10U v(1.0 - 6^) M NaC10H for the nickel (II) nitrite system and, 1.00 M NaNO, ^ M Fe(C10l)3 S :ı(^M HCIO^,(1.0 - eC^) M NaClO^ for the iron (III) nitrite system. x2 '2C^ M Fe(C10l)3 2^ M HCIO^ 1(1.0 - 12^) M NaClO^ Thus, up to C, - e,, the ionic strength is higher than 1.0 M. But at CT ?= e.» which is reached after a few points, all the free acid will practically be converted into HN0«. The ionic strength of the resulting solution will then be 1.0 M.VIII Both of the systems have also been investigated spectro- photometrically. Since, however, the nitrite complexes of nickel( II) have not a separate absorption band differs from the absorption bands of the metal and the ligand, these measurements are not possible for the nickel (II) nitrite system. Spectrophoto- metric measurements have been performed for the nitrite system of iron(III). Due to the limited ligand concentrations could be used in such measurements, no formation constants beyond the first one could, however, be determined. Two mononuclear complexes are formed in the concentration range studied for the nickel (II) nitrite system with 0! « (6 ± 1) M`1 and 32 = (12 ± 2) M~2. In the iron(III) nitrite system, on the other hand, the stability constants of the first three complexes have been determined as 3j - (3.9 ± 0.7)lû2 M***`; 32 - (5 ± 2)103 M~2; g3 - (2.8 ± 0.6)105 M~3. The errors given correspond to confidence limits on the 99.9 % level of signifi cance. In the nitrite system of iron(III), the formation of the fourth complex seems to take place in same extent. It is more likely, however, that. this indication is due to changes in the activity coefficients brought about by complex formation and by the progressive exchange of CIO` with NCL.