Yellow transient forms in the decomposition in acidic solution of the blue–violet nickel(II) complex of a trifurcated hexamine

Abstract

In the decomposition in acidic solution of the octahedral high-spin nickel(II) complex with the hexamine sen, a thermodynamically unstable, kinetically controlled, square-planar low-spin species forms, with a lifetime of 5 s.

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Metal complexes with ammonia and polyamines remain fundamental to coordination chemistry and their investigation promoted the development of crucial concepts of transition metal chemistry. The dual propensity of the amine group to behave both as a Brønsted base and as a ligand offers the opportunity to investigate the stability of metal amine complexes in water, or in other protic solvents, through competitive titration experiments. The interaction of divalent 3d metal ions (from Fe^II to Zn^II) with a variety of non-cyclic polyamines has been extensively investigated during the last fifty years.¹ In particular, both the formation of metal polyamine complexes and their destruction (e.g. on addition of acid) are fast processes and their temporal development must be studied by stopped-flow techniques. We describe here an apparent paradox observed in the demetallation in acidic solution of the nickel(II) complex of the trifurcated hexamine sen (1: N,N′-bis-(2-amino-ethyl)-2-[(2-amino-ethylamino)-methyl]-2-methyl-propane-1,3-diamine), a process displaying conflicting features in its thermodynamic and kinetic aspects.

A 1 M solution in NaClO₄, containing the polyamine 1 plus excess HClO₄ acid, was titrated with a standard solution of NaOH and the pertinent pK_a values were determined through non-linear fitting of the titration curve.† Then, a similar titration experiment was carried out on a solution containing equimolar amounts of 1 (L) and Ni^II(ClO₄)₂. Best fitting of the titration curve was obtained on assuming the formation of the following metal complex species: [Ni^II(LH₂)]⁴⁺, [Ni^II(LH)]³⁺, [Ni^II(L)]²⁺: in the case of the protonated species, one or two of the terminal amine groups are protonated. Fig. 1 shows how the concentration of the species present at the equilibrium varies in the investigated pH interval. Over the course of the titration, the solution took a pale blue–violet colour, while spectrophotometric measurements showed the development, since pH 4, of the d–d bands, at 350 and 520 nm, typically observed for high-spin Ni^II polyamine complexes of octahedral geometry. Fig. 1 shows that the intensity of the band at 520 nm superimposes well with the left branch of the concentration profile of the [Ni^II(LH₂)]⁴⁺. The spectrum does not change shape on further pH increase and reaches limiting absorbance at pH ≥ 6. This indicates that the three complexes: [Ni^II(LH₂)]⁴⁺, [Ni^II(LH)]³⁺, [Ni^II(L)]²⁺ display a similar d–d spectrum, thus possessing the same spin state and geometry.

Fig. 1 Concentration of the species present at the equilibrium in a 6.7 × 10⁻³ M solution in 1 (L) and Ni^II(ClO₄)₂, at 25 °C (ionic strength 1 M NaClO₄). Symbols: molar absorbance of the band centred at 520 nm.

Possible geometrical arrangements of the three complexes are sketched in Scheme 1, moving from high to low pH values. The structure of [Ni^II(L)]²⁺ (a in Scheme 1) reflects that determined in the solid state through X-ray diffraction studies.⁴ In particular, complex a shows full amine coordination, according to a slightly distorted octahedral geometry. On addition of the first H⁺ ion, one of the three terminal –NH₂ groups is protonated and removed from the coordination sphere, being replaced by a water molecule, to give b. Notice that the b form of the mono-protonated complex species [Ni^II(LH)]³⁺, displaying a cis conformation, can rearrange to the trans isomer c. It is difficult to assess which one of the two forms is actually more stable and predominates in solution. Undoubtedly, the two isomers are related by a fast equilibrium. The second H⁺ goes to protonate one of the two remaining coordinated –NH₂ groups, to give the cis-diaquo complex d, [Ni^II(LH₂)]⁴⁺, which maintains octahedral geometry, colour and paramagnetism. Addition of a further H⁺ leads to the destruction of the complex. Thus, progressive addition, according to a titration mode, of standard acid to a solution of [Ni^II(L)]²⁺, a, adjusted to pH 9, induces the decomposition of the complex, passing through the intermediate species b, c and d, while the violet solution gradually discolours. The described reaction pathway refers to a sequence of steady states, in which, after addition of acid, the investigated system is allowed to reach equilibrium. Then, a different experiment was carried out, in which excess acid was added to a solution of [Ni^II(L)]²⁺ in a single shot: quite surprisingly, on acid addition to the blue–violet solution, a yellow flash appeared, to vanish in a few seconds. Such an evidence prompted us to investigate the decomposition process through a stopped-flow spectrophotometric technique. In particular, syringe A of the stopped-flow apparatus (BioLogic MPS-51, connected to a J & M Tidas-10 spectrophotometer) contained a 5 × 10⁻³ M solution of [Ni^II(L)]²⁺, adjusted to pH 9 with NaOH and to ionic strength 1 M with NaClO₄; syringe B contained a solution of HClO₄, of concentration varying from 0.05 to 1 M, and adjusted to ionic strength 1 M with NaClO₄.

Scheme 1 Species involved in the decomposition in acidic solution of the [Ni^II(1)]²⁺ complex.

Fig. 2 shows the family of spectra obtained in a typical run (syringe B: HClO₄ 0.5 M). It is observed that the band at 520 nm, pertinent to the octahedral [Ni^IIL]²⁺ complex (dark grey spectra in Fig. 2), disappears, while a new band centred at 445 nm (grey spectra) rapidly develops, which reaches a limiting value of absorbance (ε = 12 M⁻¹ cm⁻¹) after 750 ms, then disappears more slowly. A d–d band at ca. 450 nm is typically observed for low-spin Ni^II complexes with tetramines, displaying a square-planar geometry.⁵ Thus, it is suggested that the transient yellow colour and band at 445 nm pertains to the [Ni^II(LH₂)]⁴⁺ complex, isomer e. Such a species originates from the protonation of the secondary amine nitrogen atom of an already protonated arm of the c isomer of the [Ni^II(LH)]³⁺ complex. The yellow diamagnetic isomer e is intrinsically less stable than the blue paramagnetic isomer d, due to the presence of two proximate charged ammonium groups (primary and secondary) on the same arm, and its formation must be therefore kinetically controlled. Analysis of stopped-flow spectrophotometric data indicated that the band at 445 nm develops with a k_obs linearly dependent upon [H⁺], according to eqn. (1):

k_obs = k₁ + k₂[H⁺] (1)

in which k₁ = 3.7 ± 0.2 s⁻¹ and k₂ = 2.9 ± 0.2 M⁻¹ s⁻¹. Such an equation corresponds to the mechanism illustrated in Scheme 2 and has been previously observed in the decomposition of [Ni^II(en)]²⁺ (en = 1,2-diaminoethane).⁶ In particular, it was observed that decomposition of [Ni^II(en)₃]²⁺ involves as a first step full protonation and release of one en molecule.⁶

Fig. 2 Selected spectra taken over the course of a stopped-flow experiment on the decomposition in acid of the [Ni^II(1)]²⁺ complex, at 25 °C and ionic strength 1 M. The dark grey colour refers to the spectra of the high-spin [Ni^II(L)]²⁺ complex, grey colour to the low-spin [Ni^II(LH₂)]⁴⁺ complex, white colour to the Ni^II aquaion formed on decomposition. Inset: time dependance of the band centred at 445 nm (yellow species).

Scheme 2 Proposed mechanism for the acid dissociation of the [Ni^II(1)]²⁺ complex and formation of the yellow species [Ni^II(LH₂)]⁴⁺ (Structure e in Scheme 1).

The structurally analogous [Ni^II(1)]²⁺ complex seems to undergo a similar decomposition pathway, with full protonation of an –NH(CH₂)₂NH₂ fragment (i.e. the equivalent of an en molecule). In the resulting [Ni^II(LH₂)]⁴⁺ complex, the chelating linear tetramine exerts in-plane interactions strong enough to stabilise the singlet state of Ni^II, responsible for yellow colour and for the band at 445 nm. Such a band decays with a lifetime τ = 5.0 ± 0.7 s, and decay is independent upon acidity. The relatively long persistence of this species has to be ascribed to the inertness of a low-spin d⁸ cation in a square environment. However, the measured value of ε of the band at 445 nm (12 M⁻¹ cm⁻¹) is distinctly lower than typically observed for low-spin Ni^II tetramine complexes (60–100 M⁻¹ cm⁻¹). This may be due to the coexistence at the equilibrium of the trans-diaquo octahedral high-spin form.⁵ The blue-to-yellow equilibrium can be displaced toward the yellow form on increasing the concentration of the inert electrolyte. In general, in an NaClO₄ saturated solution (∼7 M), the 100% of the square low-spin form is achieved. Thus, a further stopped-flow experiment was carried out, on adjusting ionic strength to 7 M with NaClO₄. Due to the viscosity of the solutions, the BioLogic apparatus could not be used. Thus, the experiment was carried out with a manually driven Hi-Tech Scientific SFA-20 Rapid Kinetics Accessory unit connected to an HP-8452 Diode Spectrophotometer. Due to the extended dead time, the preliminary development of the band at 445 nm was lost. However, the limiting absorbance of the band could be safely measured (98 M⁻¹ cm⁻¹), which corresponds to 100% of the square form. It derives that in 1 M NaClO₄ solution there exists an equilibrium mixture of 12% of low-spin square-planar (calculated from the ratio of the molar absorbance measured in 1 M and in 7 M NaClO₄) and 88% of the high-spin trans-diaquo octahedral complexes. Quite significantly, the lifetime of the yellow form in 7 M NaClO₄ (τ = 5.3 ± 0.4 s) is the same as observed in 1 M NaClO₄.

References

1. NIST Standard Reference Database 46: Critically Selected Stability Constants of Metal Complexes: Version 8.0; R. M. Smith, A. E. Martell, Critical Stability Constants, vol. 2: amines, Plenum Press, New York, 1975.
2. P. Gans, A. Sabatini and A. Vacca, Talanta, 1996, 43, 1739–1753.
3. G. K. Hollingshed, G. A. Lawrance, M. Maeder and M. Rossignoli, Polyhedron, 1991, 10, 409–413.
4. A. McAuley, S. Subramian and T. W. Whitcombe, J. Chem. Soc., Dalton Trans., 1993, 2209–2214.
5. L. Fabbrizzi and L. Sabatini, Inorg. Chem., 1979, 18, 438–444.
6. R. A. Read and D. W. Margerum, Inorg. Chem., 1981, 20, 3143–3149.

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Footnote

† Constants of protonation and complex formation equilibria were determined by processing titration data with HYPERQUAD:² pK_a1 = 2.61(1); pK_a2 = 5.08(1); pK_a3 = 6.99(1); pK_a4 = 8.87(1); pK_a5 = 9.64(1); pK_a = 9.56(1); logK([Ni^II(L)]²⁺ = 18.15(1); logK([Ni^II(LH)]³⁺ = 24.21(1); logK([Ni^II(LH₂)]⁴⁺ = 29.43(1). In a previous study on metal complexation equilibria of 1, Ni^II had not been considered.³

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