The electroreduction of azides bound to nickel(II) ions in weak acidic media

Journal of

Electroanalytical Chemistry Journal of Electroanalytical Chemistry 574 (2005) 311–320 www.elsevier.com/locate/jelechem

The electroreduction of azides bound to nickel(II) ions in weak acidic media Rafał Jurczakowski, Marek Orlik

*

Laboratory of Electroanalytical Chemistry, Department of Chemistry, University of Warsaw, ul. Pasteura 1, PL-02-093 Warsaw, Poland Received 28 June 2004; accepted 18 August 2004 Available online 28 September 2004

Abstract Azide ions undergo reduction at the mercury electrode in slightly acidified media if they are coordinated to Ni(II) ions. Under such conditions, the electroreduction of Ni(II) to nickel amalgam is followed, at more negative potentials, by the parallel reduction of the ligand. The macro-scale electrolysis of slightly acidified solutions of NiðIIÞ–N
3 complexes showed the formation of ammonia and the evolution of bubbles of colorless gas. The electrode mechanism involves the direct electroreduction of N
3 ions in the coorþ þ
dinated state according to the reaction scheme: N
3 þ 4H þ 2e ! N2 þ NH4 as the suggested main reaction pathway, accompanied by a side reduction of N
3 by nickel amalgam. In terms of this mechanism, the average coordination number of the reacting NiðIIÞ–N
3 complexes could be deduced. A molecular mechanism of the electroreduction of the azide ion bound to Ni(II) central ion was suggested.
2004 Elsevier B.V. All rights reserved. Keywords: Azide complexes of nickel; Ligand electroreduction; Chronocoulometry; Streaming mercury electrode

1. Introduction Pseudohalogenide anions are known to exhibit distinct similarities when their chemical properties are compared [1,2]. However, electrochemical studies of such ions are limited. Exceptions are studies of thiocyanate and selenocyanate ions which were found to undergo catalytic reduction, accompanying the nickel(II) electroreduction at mercury electrodes [3–5]. Contrary to results obtained for SCN
and SeCN
, analogous studies carried out on the azide complexes of nickel(II) [6] showed only the electroreduction of the central ion at the mercury electrode. Based on these results one could conclude that it is more difficult for the azide ions to undergo heterogeneous electron uptake. However,

*

Corresponding author. Tel.: +48 22 822 02 11x245; fax: +48 22 822 59 96. E-mail address: [email protected] (M. Orlik). 0022-0728/$ - see front matter
2004 Elsevier B.V. All rights reserved. doi:10.1016/j.jelechem.2004.08.012

under certain conditions, azide ions appear to be reducible. Masˇek [7] showed that hydrazoic acid HN3 undergoes electroreduction on mercury, but only in strongly acidic medium (8–18 mol dm
3 H2SO4), giving rise to a polarographic wave of ([H2SO4] dependent) E1/2 varying from
1.2 to ca.
0.7 V (vs. the Hg2SO4jHg electrode in 17 mol dm
3 H2SO4). The predominant reduction path was found to be the one-electron reduction of HN3 to molecular nitrogen and hydrazine. In a more recent study Dalmia et al. [8] found that azide ions can be reduced from non-acidified solutions of NaN3 at the platinum electrode at potentials more negative than
0.9 V (vs. the saturated HgjHg2SO4 electrode), with the formation of N2, N2H4 and possibly NH3. Also the oxidation of N
3 at more positive potentials appeared possible under such conditions. During our recent research into new dynamic electrochemical systems we found the conditions, for which the complicated electroreduction scheme of azide complexes of nickel(II) at the streaming mercury electrode gives

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R. Jurczakowski, M. Orlik / Journal of Electroanalytical Chemistry 574 (2005) 311–320

rise to bistable and – particularly rarely reported – tristable behavior [9]. One of the key pathways of this mechanism is the electroreduction of the coordinated N
3 ions in weak acidic media, the possibility of which was only briefly mentioned by us in our previous work [9]. To our knowledge the N
3 electroreduction under such conditions has never been reported before in the literature and therefore the aim of this paper is to describe in more detail these new aspects of the electrochemical behavior of the NiðIIÞ–N
3 complexes and to elaborate the appropriate electrode mechanism of the N
3 electroreduction.

2. Experimental Commercially available p.a. NaN3 (Fluka) was used without further purification. For varying NaN3 concentrations, the constant ionic strength of the solutions was maintained constant by addition of appropriate amounts of p.a. NaClO4 (APOLDA), purified by prior recrystallization from water. In order to remove the Al component, the pure nickel–aluminum catalyst alloy (Koch-Light), was treated three times with hot 30% NaOH and washed with distilled water. The remaining wet powder, being a reactive form of pure nickel, was used. Crystalline nickel(II) perchlorate, Ni(ClO4)2 Æ 6H2O was obtained by neutralization of the p.a. basic nickel(II) carbonate (POCh) with the appropriate amount of p.a. perchloric acid (APOLDA), followed by recrystallization from water. P.a. nickel(II) ammonium sulfate (NH4)2Ni(SO4)2 Æ 6H2O (POCh), p.a. copper sulfate CuSO4 Æ 5H2O (POCh), spectral pure NH4Cl (Ciech), p.a. 35% hydrochloric acid (POCh), p.a. KMnO4 (POCh), p.a. potassium bromate(V) KBrO3 (POCh), p.a. KBr (POCh), p.a. NaOH (POCh), p.a. phenol (Loba-Chemie), p.a. acetone (POCh), p.a. ethanol (POCh) and p.a. dimethylglyoxyme (POCh) were used without further purification. For all solutions thrice-distilled water, additionally purified in a final step using Millipore filters, was used. Two types of working electrode were used in typical voltammetric experiments: a classical dropping mercury electrode DME (with capillary flow m = 0.354 or 0.503 mg s
1) and a home-made streaming Hg electrode, the construction of which was described by us in detail previously [10]. Both electrodes were filled with p.a. mercury (POCh, Poland). At the DME the voltocoulometric (charge–voltage) dependences were recorded with the use of a homemade chronocoulometer. In these experiments the total duration time of the single mercury drop was divided into time intervals: t1 (1–4 s long) within which the electrode was maintained at either an indifferent or a negative pre-

electrolysis potential, and the interval t2 (10–200 ms long) within which the potential was changed into the desired value at which the current was integrated over time t2. In some experiments this potential program included also a third potential jump of duration t3 (10– 200 ms), and then the integration of the current occurred over this time. Using the streaming electrode [10], the voltammetric current–potential (I–E) dependences were recorded with a EG&G/PARC potentiostat mod. 273, controlled by an IBM PC compatible computer with commercial M270 software. For electrolysis on the macro-scale, a mercury pool cathode with a surface area of 2.83 cm2 was used. In all measurements a three-electrode potentiostatic circuit was applied. The saturated (KCl) calomel reference electrode (SCE) was separated from the solution studied by a salt bridge and the uncompensated ohmic potential drops were minimized by using a Luggin capillary. A platinum wire (A  2 cm2) served as the counter electrode, separated from the working electrode space by the glass frit. All measurements were done for a temperature of 298.15 ± 0.05 K, controlled by an ultrathermostat U4 (VEB MLW Pru¨fgera¨te-Werk Medingen Dresden). Prior to the electrochemical measurements the sample solutions were thoroughly deaerated using pure argon (5.0, Praxair, Poland). This had been additionally purified from traces of oxygen by passage through a glass column filled with heated copper wires, obtained from p.a. cupric oxide (CuO + Cu2O) wires (POCh, Poland), which had been reduced to the metal by gaseous hydrogen at the appropriately enhanced temperature. Spectrophotometric measurements were made by recording the UV/Vis spectra for samples of optical length = 1.0 cm, with a Hewlett–Packard diode-array spectrophotometer, mod. 84852A.

3. Results 3.1. The voltammetric characteristics of NiðIIÞ–N
3 electroreduction The sample solutions prepared by direct mixing of pure Ni(ClO4)2 and NaN3 solutions always yielded somewhat drawn-out and slightly irreproducible voltocoulometric (Q–E) dependences, presumably due to the partial precipitation of nickel(II) hydroxides, forming in the slightly alkaline medium of the hydrolyzing sodium azide. In order to avoid these perturbations, prior to the measurements the samples were always acidified with HClO4 solution, in the minimum proportion of 6 mmol HClO4 per 1 mol of NaN3. Fig. 1 shows a set of typical voltocoulometric curves, obtained for the solution prepared in this way, for various sampling

R. Jurczakowski, M. Orlik / Journal of Electroanalytical Chemistry 574 (2005) 311–320

5

2.0

313

0.20

4

4

0.15

3

-I / mA

-Q / µC

1.5

2

1.0

3

0.10 2 1

0.05

0.5

1

0.00 0.0 0.4 0.6

0.8

1.0

1.2

1.4

0.6

0.8

1.6

1.0

1.8

1.4

1.6

1.8

-E / V

-E / V Fig. 1. Voltocoulometric normal pulse faradaic curves (Q–E) recorded for 8 · 10
4 mol dm
3 Ni(ClO4)2 + 2 mol dm
3 NaN3 + 1.2 · 10
2 mol dm
3 HClO4, for different sampling (pulse duration) times t2: (1) 10 ms, (2) 50 ms, (3) 90 ms, (4) 130 ms and (5) 170 ms. For each drop a potential pulse of duration t2 was applied after t1 = 4 s, maintaining the electrode at the indifferent potential E1 =
0.4 V. The recorded total charge was corrected for the capacitive contribution by subtracting of the charge recorded under the same conditions for the pure 2 mol dm
3 NaN3 solution. Electrode surface area at the sampling time: A = 0.0107 cm2.

times. It is noteworthy that the Q–E dependences presented were obtained from the subtraction of the capacitive charge (recorded for the pure 2 mol dm
3 NaN3 solution) from the total charge recorded for this electrolyte after addition of appropriate amounts of Ni(ClO4)2 and HClO4. Assuming that the double layer charge did not change significantly upon addition of Ni(ClO4)2 and HClO4 one can conclude that Fig. 1 accurately shows the isolated faradaic charge associated with the NiðIIÞ–N
3 electroreduction. Although the waves do not form an ideally horizontal plateau, their limiting charges Ql are close to the diffusion-controlled values, calculated for the two-electron reduction of the Ni(II) species from the integrated Cottrell equation (the minus sign indicates the cathodic charge): 1=2

Qd ¼


1=2 0 2nFADox cox t2 1=2 p

ð1Þ

with Dox = 6.9 · 10
6 cm2 s
1 [6,9]. For example, for t2 = 10 ms, the theoretical charge Qd =
0.49 lC and for t2 = 170 ms, Qd =
2.02 lC (cf. similar values for the corresponding curves (1) and (5) in Fig. 1). It is also noteworthy that in these curves, a sequence of two regions of negative differential resistance (NDR) is observed (the first at E =
0.85 to
1.05 V and the next at E <
1.2 V). The existence of the NDR regions for this process was shown to be a source for its bistable and tristable behavior in our earlier studies [9]. When the sampling time is markedly shorter than in the voltocoulometric experiments, the kinetic effects

Fig. 2. Dependence of the faradaic current I on the potential E of the streaming mercury electrode for the solutions containing a constant concentration of 2 · 10
3 mol dm
3 Ni(ClO4)2 and varying concentrations: (1) 0.331, (2) 1.0, (3) 1.25 and (4) 1.58 mol dm
3 NaN3. A constant ionic strength (I = 2) of the solutions was maintained by addition of NaClO4 and the solutions were acidified in the proportion c(HClO4):c(NaN3) = 6 mmol:1 mol. Parameters of the streaming electrode: capillary flow m = 212 mg s
1, electrode length lmax = 2.85 mm, electrode radius r = 0.055 mm.

within the potential region of the formation of the quasi-plateau of the waves are much more strongly pronounced. We observed these effects under conditions of a slow linear potential scan with the streaming mercury electrode replacing the dropping one. For the parameters of the streaming electrode used in our experiments the maximum sampling time (at the end of a mercury stream of length lmax) was equal to 1.74 ms. Typical I–E curves recorded for the different NaN3 and HClO4 concentrations (with the ionic strength I = 2 kept constant) and corrected for the contribution from the capacitive current in the same way as described above for the voltocoulometric experiments, are shown in Fig. 2. First, these curves reveal more distinctly, than for the voltocoulometric conditions, the formation of two regions of negative differential resistance (NDR), originating from the overlapping of a single NDR region of the first, main wave with a small post-wave. Second, the predominant control of the current by the rate of the charge-transfer step now manifests itself over the entire potential range, since even the maximum current of the curves from Fig. 2 remains below the diffusion-controlled value Id =
0.270 mA, calculated for the two-electron reduction of Ni(II) from the following expression [11]: Id ¼


1=2 1=2 0 4nFD1=2 lmax cox ox m 1=2

qHg

ð2Þ

with the minus sign indicating the cathodic current, qHg is the density of mercury, lmax is the length of the mercury stream, and other quantities having their usual meaning.

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25

2.5 IV

5 4

-Q / µC

III

2.0 3

15 II 2

10 5

Q (II,III) / Q (I) l l

20

1.5

1.0

1

I

0.5

0 0.0 0.6

0.8

1.0

1.2

1.4

0

1.6

-E / V Fig. 3. Influence of the addition of HClO4 on the voltocoulometric faradaic Q–E dependences for the electroreduction of 4.0 · 10
3 mol dm
3 Ni(ClO4)2 + 0.6 mol dm
3 NaN3. Concentrations of added HClO4: (1) 6; (2) 16; (3) 36; (4) 76; (5) 156 mmol dm
3. The dropping mercury electrode was held for t1 = 2 s at the indifferent potential E =
0.5 V prior to the potential pulse (t2 = 100 ms); electrode surface area at the sampling time A = 6.6 · 10
3 cm2.

40

60

80

100

120

140

160

c(HClO4) /mmol dm-3 log10[HN3] ≈ log10c(HClO4) 2.5

-0.6 -0.8 -1.0 -1.2 -1.4 -1.6 -1.8 -2.0 -2.2 -2.4

2.0

Q (II,III) / Q (I) l l

3.2. The influence of HClO4 on the I-E characteristics of NiðIIÞ–N
3 electroreduction

20

(a)

1.5

1.0

For the diagnosis of the electrode mechanism underlying the Q–E or I–E characteristics from Fig. 1 or 2, the role of the acidity of the samples was studied further. We found that the magnitude of the ‘‘hump’’ (small post-wave) in Fig. 2, separating the two NDR regions, was strongly dependent on the amount of added HClO4. Fig. 3 shows how this effect manifests itself on the voltocoulometric curves of nickel(II) in 0.6 mol dm
3 NaN3. Upon addition of increasing amounts of HClO4, a pattern of two (II, III) distinct cathodic waves, following the first one (I) that corresponds to the Ni(II) electroreduction, evolves. In the following report and discussion the waves (II) and (III) will be always considered as a double wave of a common electrochemical origin. Obviously, considering the limiting charge of the waves (II) + (III), compared to that of the wave (I), the former pair have to be ascribed to the electroreduction of species that is present in the solution in an amount greater than that of nickel(II), but – on the other hand – significantly lower than the total NaN3 concentration: 0.6 mol dm
3, which is 150 times higher than that of the nickel(II) ions. Undoubtedly hydrogen ions participate in this process but it is particularly noteworthy that the dependence of the maximum charge of the waves (II, III) (and thus also of the total limiting charge of the waves (I, II, III)) on the total concentration of added HClO4 is non-linear and apparently tends to a certain limiting value. This effect, shown in Fig. 4(a), forms a dependence formally proportional to the logarithm of c(HClO4) which suggests linear correlation with the log-

0.5

0.0 (b)

4.6

4.8

5.0

5.2

5.4

5.6

5.8

6.0 6.2

6.4

pH

Fig. 4. (a) Dependence of the ratio of the limiting charges of the (II, III) and (I) cathodic waves (Ql(II,III)/Ql(I)) from Fig. 3 on the total concentration of added HClO4. The solid line y = 0.593log10(0.232x) was fitted (determination coefficient r2 = 0.9997) to the experimental points (); (b) the corresponding linear dependences of Ql(II,III)/Ql(I) on pH: y = 7.965
1.253x(r2 = 0.9979) and on log [HN3] (logc(HClO4)); y = 3.229 + 1.365x (r2 = 0.9997), the values of the concentrations in mol dm
3 were used for the calculation of their logarithms.

arithm of concentration of the various acid–base species. Therefore for all the samples, the thermodynamic activity of H+, as well as the equilibrium concentration of HN3 (being practically equal to c(HClO4) due to high excess of N
3 ions), was calculated from the set of nonlinear equations for the acid–base equilibria in the NaN3 + HClO4 system studied, with the dissociation constant for HN3 pKa 4.4, the ionic product of water in 25 C pKw = 14.0, and the ionic activity coefficients corresponding to the ionic strength equal to 0.6 (determined practically by the NaN3 concentration). Very good linear dependences of the charge ratio on pH and log[HN3] were indeed found (see Fig. 4(b)), showing that this type of correlation of the electrochemical prop-

R. Jurczakowski, M. Orlik / Journal of Electroanalytical Chemistry 574 (2005) 311–320

erties of the system with its bulk acid-base equilibria, is observed at least within the concentration range of HClO4 studied. Similar linearities were observed for the dependence of either the total limiting charge of the waves (I, II, III) or of the limiting charge of the waves (II, III) alone, on pH and log c(HClO4). All these correlations were better than the correlations with the square root of the respective concentrations. If higher excess of HClO4 is applied, the strong protonation of ligands causes the decomposition of the azide complexes of nickel(II), which is known a phenomenon [12]. For the diagnosis of the origin of the waves (II) and (III) the following two experimental observations are also of crucial importance: (i) these waves were observed only for samples containing simultaneously all three species, Ni(ClO4)2, HClO4 and NaN3; thus acidified solution of NaN3, in the absence of Ni2+ ions did not yield any faradaic response within the potential region studied, and (ii) the double potential-step voltocoulometric experiments with the cathodic pre-electrolysis of NiðIIÞ–N
3 at different potentials, followed by measurement of the charge recorded at the positive potentials corresponding to the oxidation of nickel amalgam, proved that a significant increase of the cathodic waves (II, III) upon addition of HClO4 did not cause a proportionally enhanced production of nickel amalgam, which was thus unambiguously related to wave (I) only. þ 3.3. Preparative electrolysis of the NiðIIÞ–N
3 –H system at the mercury electrode

If the reduction of azide ions is a source of the cathodic waves (II, III), one can expect the formation of at least one of the following products of this process: NH3, N2 or N2H4. In order to verify this assumption we performed coulometric electrolysis of acidified Ni(ClO4)2 + NaN3 solutions at a mercury pool as the macro-cathode, under potentiostatic conditions. The resulting solutions were analyzed for the NH3 content and the remaining Ni(II) in the following way. For the detection and determination of ammonia, first the residual azide ions were removed from the solution after the electrolysis, in the following way. 1.0 mol dm
3 H2SO4 was added to 5 cm3 of the sample (in the proportion – initial c(NaN3):c(H2SO4) = 1.0:1.2). This solution was further diluted to ca. 100 cm3 and evaporated to 1/3 of its original volume; during this procedure the volatile HN3 was expected to evaporate. The resulting solution was transferred to the distillation flask and made alkaline with 30% NaOH (boiled for 60 min prior to this experiment, to remove the

315

eventual trace amounts of NH3). From this solution the volatile products were distilled off in the closed equipment and absorbed in distilled water, slightly acidified with HCl. The presence of ammonia in this distillate was confirmed qualitatively by the classical Nessler test. The blank probe, i.e., the analysis of the distillation products of the sample of the same initial composition, but not electrolyzed, indicated only trace amounts of NH3. Ammonia in the samples after the electrolysis was further quantitatively determined spectrophotometrically using the indophenol method [13]. The 5 cm3 sample of the solution after the macro-electrolysis was again acidified with 1 mol dm
3 H2SO4 (in the same molar proportion as above) and the resulting solution was diluted to ca. 100 cm3, evaporated to 1/3 of its original volume and transferred to a 50 cm3 measuring flask. There appropriate amounts of sodium phenolate and of sodium chlorate(I) were added. Sodium phenolate was prepared ex tempore by mixing appropriate amounts of an acetone–ethanol solution of phenol and 30% aqueous solution of NaOH. Sodium chlorate(I) was also prepared ex tempore by passage of gaseous chlorine, obtained from KMnO4 and HCl, through NaOH solution. The calibration curve was obtained using standard solutions of ammonia, prepared from a solution of NH4Cl, dried at 100 C prior to weighing. The amount of ammonia determined in this way for the samples after the electrolysis was recalculated into the equivalent faradaic charge (under the assumption of the stoichiometry of Eq. (3), discussed later) and compared with the total electric charge measured through the integration of the current over the electrolysis time. The results of these calculations and determinations are collected in Table 1. It is important to note that ammonia was detected after the electrolysis was performed, not only for the electrode potential E =
1.3 V, at which waves (II) and (III) are formed in voltocoulometric measurements (cf. Fig. 3), but also at the potentials
0.7 and
0.8 V, corresponding to the foot and the plateau, respectively, of wave (I). For the determination of nickel(II) in the same solutions after the electrolysis also, the spectrophotometric method was applied. It was based on the formation of a soluble, red complex compound of dimethylglyoxime (DMG) with Ni(IV). Ni(II) in the sample was oxidized to Ni(IV) with bromine, prepared ex tempore from an acidified (HCl) solution of the stoichiometric (in the molar proportion 5:1) amounts of potassium bromide and potassium bromate(V) [13]. The calibration curve was prepared using standard solutions of p.a. nickel(II) ammonium sulfate. From the results of these analyses, compared to the initial content of nickel(II), the amount of nickel(II) which underwent reduction and the equivalent faradaic charge were determined. These results are collected in the last two columns of Table 1. Since, for

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R. Jurczakowski, M. Orlik / Journal of Electroanalytical Chemistry 574 (2005) 311–320

Table 1 Results of the electrolysis of Ni(ClO4)2 + NaN3 samples, acidified with HClO4, at the mercury macro-cathode c(NaN3) (mol dm
3)

c(HClO4) (mol dm
3)

c(Ni(ClO4)2) (mmol dm
3)

E (V)

t (s)a

Qel/Cb

n(NH3)c (mol)

Q(NH3)d/C

Qel
Q(NH3)/C

Dn(Ni2+)e (mol)

Qred(Ni2+)f/C

2.0 2.0 2.0 2.0

0.10 0.05 0.10 0.05

3.74 2.50 3.74 2.50


1.3
1.3
0.8
0.7

1000 3000 3000 10000

20.28 4.566 13.62 1.17

1.19 · 10
4 5.84 · 10
6 6.05 · 10
5 2.85 · 10
6

22.96 1.126 11.68 0.55

<0 3.44 1.94 0.62

0 1.19 · 10
5 7.05 · 10
6 8.13 · 10
7

0 2.3 1.36 0.16

a b c d e f

Electrolysis time. Charge from the integration of current over the electrolysis time. Determined by the indophenol method. Charge equivalent to the amount of ammonia in (c), based on the stoichiometry of Eq. (3). The decrease of Ni(II) content in solution determined by the dimethylglyoxime method. Charge equivalent to the decrease of Ni(II) content in solution, assuming two-electron reduction to the nickel amalgam.

the interpretation of data presented in Table 1, it is useful to consider the possibility of the redox interaction between Ni(0) and the azide ions, the following additional experiments were performed. Solutions of 2 mol dm
3 NaN3 acidified with 0.05 mol dm
3 HClO4 were shaken with either nickel amalgam, prepared ex tempore electrolytically from the Ni(ClO4)2 solution or with reactive Raney nickel. Although the quantitative analysis of the resulting samples was not easy, since it was difficult to centrifuge the colloidal nickel, the oxidation of nickel was observed in both cases, but the formation of ammonia was detected unequivocally only in the samples treated with the Raney nickel. Finally, no strong reduction properties of the solution after the electrolysis were detected (e.g., copper(II) ions were not reduced to metallic copper in alkaline solution [14]) and in this way the formation, at least of detectable amounts of hydrazine, was excluded.

4. Discussion As we emphasized in the Introduction, before our present studies the electroreduction of NiðIIÞ–N
3 complexes at mercury electrodes was considered to be a process involving only the central ion. The NiNþ 3 complex was proposed to be a directly reacting species [6]. Thus if complexes of a higher coordination number existed in the solution, their partial dissociation to NiNþ 3 species had to precede the charge-transfer step [6]. No manifestation of the region of negative differential resistance was reported. We do not question any of the main points of these earlier mechanisms leading to the formation of nickel(0) amalgam, as they were based on solid kinetic studies. We show here, however, that these mechanisms do not explain all the aspects of the electrochemical behavior of the azide complexes of nickel(II) at mercury electrodes and should be extended to include the additional reaction pathways, discovered by us.

4.1. Inner- and outer-sphere electroreduction of the NiðIIÞ–N
3 complexes First, the formation of the region of negative differential resistance needs at least an outline explanation. Although in this paper, we do not analyze this particular effect in detail, its very probable origin can be derived from the comparison of the electrochemical characteristics of both the azide complexes of nickel(II) (discussed here) and of thiocyanate complexes of nickel(II) [15]. As far as the NDR region without the ‘‘hump’’ in the I–E or Q–E characteristics of the NiðIIÞ–N
3 electroreduction is considered, it exhibits quite a similar shape to the I–E (Q–E) dependences for Ni(II)–SCN
[15], and also for In(III)–SCN
electroreduction [16,17], which are a result of the desorption of the catalytic SCN
ions from the electrode surface, when its electric charge becomes increasingly negative. We suggest thus an analogous mechanism for the NDR formation in NiðIIÞ–N
3 electroreduction. In conjunction with the earlier findings on the stoichiometry of the reacting particle [6], the electroreduction of NiðIIÞ–N
3 , forming wave (I) in Fig. 3, should thus presumably proceed via the intermediate NiNþ 3 species, adsorbed on mercury due to the surface affinity of the azide anions. At potentials even preceding the formation of the NDR region, the ligand-induced adsorption of this complex compound is, however, so weak that it is difficult to detect it directly, but it still remains kinetically significant. With an increasingly negative potential (with respect to the pzc that was estimated by us as close to
0.583 V vs. SCE in 1 mol dm
3 NaN3 and close to
0.666 V in 2 mol dm
3 NaN3), the progress of desorption of the azide ions from the electrode surface causes a decrease of the surface concentration of N
3 and, consequently, a decrease of the rate of the electrode process, which effect leads to the NDR region. This mechanism seems to be at least semi-quantitatively justified, since it is known [18] that the adsorption of azide ions on mercury, being stronger than that of chlorides, becomes negligible at potentials more negative than ca.
1 V. In other words, we postulate that the electro-

R. Jurczakowski, M. Orlik / Journal of Electroanalytical Chemistry 574 (2005) 311–320

reduction of the central ion in the NiðIIÞ–N
3 complex starts from the inner-sphere route leading to nickel(0) amalgam, for which the pathway involves adsorbed N
3 ligand as a bridge for the facilitated electron transfer. At much more negative potentials (ca. <
1.6 V) there is also a probable contribution from the highly irreversible outer-sphere reaction pathway (cf. wave (IV) in Fig. 3), but the current of that process significantly overlaps with the high current of hydrogen evolution at the end of the available potential range at mercury electrodes. 4.2. Electroreduction of the azide ligand bound to nickel(II) ion The fact that ammonia is produced within the wide potential range of NiðIIÞ–N
3 electroreduction proves that the azide ions must undergo electroreduction under such conditions. Since simultaneously, bubbles of colorless gas evolve (suggesting thus its low solubility in water) and the solution does not exhibit distinct reductive properties, we suggest – at least as the predominating pathway - the following reaction scheme of N
3 electroreduction: þ
N
3 þ 3H þ 2e ! NH3 þ N2

ð3Þ

or, at lower pH, at which ammonia is largely protonated: HN3 þ 3Hþ þ 2e
! NHþ 4 þ N2

ð4Þ

These reaction schemes explain why the addition of HClO4 to the samples was necessary for the above processes to proceed, as H+ ions are required to form a thermodynamically stable product: NH3 or NHþ 4, respectively. In the following discussion it will be assumed that the waves (II) and (III), the limiting charge of which rises with H+ concentration, are both due to the electroreduction of the azide ligand, facilitated by the simultaneous electroreduction of Ni(II) ions. Let us first consider theoretically the simple catalytic scheme: nickel(II) ions are reduced to Ni(I) or Ni(0) and this (these) species reduce the azide ions in the solution. However, this scheme – at least as a source for the (II) + (III) waves – has to be rejected: waves (II) and (III) would then largely overlap with the first electroreduction wave (I) of Ni(II) ions, since obviously Ni(I) or Ni(0) species appear already at the foot of this wave. Instead, since the electroreduction of N
3 ions occurs at significantly more negative potentials, giving rise to a separate set of waves (II) + (III), it becomes obvious that the azide ligands undergo a facilitated heterogeneous electron uptake, still remaining in the state coordinated to the nickel(II) ions. Evidently in this state, the activation energy for the azide electroreduction process is significantly lower than in the unbound state. The enhanced reactivity of the lig-

317

and in the coordinated state is a well-known phenomenon (cf., e.g., [19]). This mechanistic conclusion is supported further by the fact that the limiting charge of the cathodic waves (II) + (III), ascribed by us to N
3 electroreduction, tends to an asymptotic value, in spite of the increasing concentration of added HClO4. The charge of the Ni(II) electroreduction, forming the first (I) cathodic wave, is due to the amount of nc moles of the complex compound that reaches the electrode surface at unit time, so 2nc moles of electrons flow at this time, when nickel amalgam is produced. At more negative potentials, when the limiting height of the waves (II) + (III) is reached, the coordinated azide ligands also undergo electroreduction. If the coordination number of the NiðIIÞ–N
3 complex is equal to p and the amount of H+ ions is sufficient, then all p ligands undergo simultaneous electroreduction with Ni(II). According to the reaction scheme (3) or (4), this process consumes an additional 2pnc moles of electrons. Since, in terms of the mechanism discussed, the diffusion-controlled height of the waves (I) and (II) + (III) is determined by the diffusion coefficient of the entire complex particle, and not by the individual diffusion coefficients of the central ion and the ligands, the ratio of the limiting (asymptotic) height of the waves (II, III) to that of the first wave (I) is equal simply to the coordination number of the reacting complex, or – more strictly and more generally - to the average coordination number of the mixture of complexes of various coordination numbers: Ql ðwaves II; IIIÞ 2pnc ¼ ¼ p: Ql ðwave IÞ 2nc

ð5Þ

From our measurements for 0.6 mol dm
3 NaN3 it follows that the value of p, determined from Eq. (5), tends to a value of ca. 2.5 (cf. Fig. 4). It should be noted that this value is quite reasonable if compared with the earlier studies of the complexation equilibria in the NiðIIÞ–N
system [6]: starting from c(NaN3) = 0.5 3 mol dm
3, not only are [NiN3]+ and [Ni(N3)2] species present, but also higher complexes (anionic species) should begin to form, whereas only for c(NaN3) = 2–4 mol dm
3 should the four-coordinated species [Ni(N3)4]2
predominate. The above mechanism can be summarized in one equation of the electrode process, which occurs at such negative potentials that both the central ion and the ligand undergo electroreduction: ½NiðN3 Þp 

ð2
pÞþ

þ 4pHþ þ 2ð1 þ pÞe


! NiðHgÞ þ pNHþ 4 þ pN2

ð6Þ

In terms of this reaction route one can suppose that the partial pre-dissociation of the [Ni(N3)p](2
p)+ species to the directly reducible [NiN3]+ form, postulated in [6], is – at the potentials corresponding to the electroreduction

R. Jurczakowski, M. Orlik / Journal of Electroanalytical Chemistry 574 (2005) 311–320

4.3. Molecular mechanism of the parallel electroreduction of Ni(II) and coordinated N
3 ligands Based on knowledge of the structure of azide ions and their complexes with transition metals [1,2,20] one can suggest the following detailed mechanism of the ligand electroreduction, which may help in understanding why just ammonia (or ammonium ion) and molecular nitrogen are the final products of this process. The bonds in the free azide ions can be formally expressed in terms of the appropriate resonance structures. According to the schematic Fig. 5, it is suggested that the formation of the [NiN3]+ complex enhances the relative contribution of this resonance structure of the ligand which contains the triple N„N bond, being a precursor for the N2 molecule. When the electrons are transferred from the electrode, the Ni-N bond is repelled + back to the N
3 ion, which particle, in the presence of H ions is not released as the entire unit, but decomposes

[

−

+

−

N

N

N

3

2

1

− −

+

N

N

] + Ni

N

2+

N

+

−

N

N

Ni2+

n tio uc ed r tro ec el

+ 4H+

N2 +

N

N

N

+ 2e

+ 2e

Ni2+ cleavage

of the entire complex – coupled or replaced with the electrochemical destruction of the azide ligands. On the basis of our results it is however, difficult to decide whether H+ ions have to protonate the ligands prior to the electroreduction or if they are transferred simultaneously or later from the bulk HN3 species. In the presence of a large excess of N
3 ions, practically all added H+ ions are transformed into HN3 particles. Considering the electronic structure of the azide ions, one should expect them to be protonated more strongly in the unbound than in the coordinated state. However, the fact that the coordinated azide ions undergoing electroreduction give rise to a double wave (II) + (III), the relative proportion of which varies with the amount of added HClO4, suggests that this process involves at least two different complex species of NiðIIÞ–N
3 , differing presumably in their degree of protonation or with the composition of the inner sphere (e.g., by the formation of mixed complexes with H2O or OH
). Evidently the complex form, which predominates at the higher H+ concentrations, gives rise to the N
3 electroreduction wave at less negative potentials (i.e., to the wave (II)). The lack of protonation constants for the NiðIIÞ–N
3 complexes makes a more quantitative explanation of the phenomena discussed difficult. Another source of complication is the fact that the pH of the solution in the pre-electrode layer should change as a function of the potential, since hydrogen ions are consumed in the azide electroreduction, and the reactant HN3 and the product NHþ 4 have different pKa values. Finally, the limiting charge of the (II) and (III) waves is slightly perturbed by the contribution from the region of the negative differential resistance, affecting the limiting charge of wave (I) as their background. Thus, there are several factors that make the I–E dependence for the process studied rather complex, at least in detail.

cleavage

318

+

NH4

Hg

Ni(Hg)

Fig. 5. Possible molecular mechanism of [NiN3]+ electroreduction explaining the formation of ammonium cations (ammonia) and molecular nitrogen N2 in the presence of a sufficient amount of H+ ions.

into NHþ 4 ion and N2 molecule as the more thermodynamically stable products. It is useful to compare this mechanism with the available data on the structure of azide ions in the unbound state and bound to nickel(II) ions. In the free, symmetrical N
3 ion, the N–N distance is equal to 115 ± 2 pm [1,2] and the formal charge at either of the peripheral N-atoms (i.e., N(1) and N(3)) is equal to
0.8, whereas at the middle, N(2) atom it is +0.6. Recent theoretical calculations for the NiNþ 3 and Ni(N3)2 species [20] give certain insight into the redistribution of electrons when the azide ligands become coordinated to the nickel(II) ion. As an example the NiNþ 3 ion in its singlet ground state will be invoked as a particularly stable species. For this species, the following Ni–N(1), N(1)–N(2) and N(2)–N(3) distances were obtained, respectively: 172.4 pm, 119.6 pm and 111.2 pm. The N(2)–N(3) bond is thus (i) the shortest, (ii) shorter than the N–N distance in the free N
3 ion, and (iii) close to the N–N distance in the free dinitrogen molecule: 109.8 pm [2]. This comparison confirms the postulated enhanced contribution of this resonance structure of the N
3 ligand, which includes the precursor of the N2 molecule. According to [20], the relatively long Ni–N(1) distance is formed by the predominantly covalent r-bonding, without a contribution from the p-bonding. It is noteworthy that the latter conclusion differs from more common, earlier [1] suggestions that the azide ions serve rather as the pdonor ligands in their complexes with the transition metal ions. The calculated [20] charge densities for the NiNþ 3 species are: 1.6927 (Ni),
0.7427 (N(1)), 0.0059 (N(2)), 0.0440 (N(3)). The negative charge is thus localized on the N(1) atom, whereas the N(3) atom carries only a small positive charge, both conclusions being roughly in accord with the scheme in Fig. 5. A very small positive charge at the central N(2) atom is however not in accord with the +1 value predicted by the respective resonance structure in Fig. 5. Considering this

R. Jurczakowski, M. Orlik / Journal of Electroanalytical Chemistry 574 (2005) 311–320

319

difference and an N(1)–N(2) bond length only slightly greater than that of the N(2)–N(3) bond, one can conclude that, in the coordinated azide ion, the order of the N(1)–N(2) bond is higher than one, presumably due to the non-bonding electrons of N(1) moving towards the N(2) atom. This conclusion implies a certain contribution from other resonance structures to the real structure of the N
3 ion as a ligand. Of course, the applicability of all the conclusions drawn here to the interpretation of the results of our studies of the N
3 electroreduction requires an assumption that the solvation (hydration) processes which occur in our real system, as well as the electric field of the double layer at the electrodejsolution interface, do not significantly affect the theoretical electron distribution in the free NiNþ 3 ion. It is noteworthy that similar relations between the bond lengths N(1)–N(2)–N(3) were found also for the Ni(N3)2 species [20].

However, from Table 1 it follows that the amount of Ni2+ that undergoes reduction is always lower than the amount equivalent to the value of [Qel
Q(NH3)] (note that the data in the first row, producing slightly negative charge difference, are most probably due to a certain experimental error). This means that there is also an additional side path of nickel(0) oxidation, not involving azide ions, and this is most probably the interaction of Ni(Hg) with water containing H+ ions and presumably some amount of oxygen absorbed during the preparation of the sample for the analysis. The amount of nickel(II) remaining in the samples appears thus to be a rather uncertain quantity. The possible participation of the intermediate Ni(I) in these reactions is difficult to verify in terms of our experimental data and methods used.

4.4. Possible interaction of Ni(0) and N
3 species

The similarities between the electrochemical properties of the pseudohalogenide ions in our studies found one more manifestation in the, so far, unknown ability of the azide ligand to undergo electroreduction in the state bound to nickel(II) ion in the complex compound. There is also a possibility of a side redox interaction of N
3 with Ni(0) (and presumably also Ni(I) species). Compared to the Ni(II)–SCN
system, studied earlier, the main difference is the formation of the distinct separate I–E (or Q–E) wave for the electroreduction of N
3, whereas for the Ni(II)–SCN
the electroreduction of the central ion and the ligand largely overlap, with a relatively low efficiency of the SCN
electroreduction at ambient temperatures. Instead, all the azide ions bound to Ni(II) can be reduced under such conditions, provided that the amount of hydrogen ions is sufficient (taking into account the acid–base equilibria) to produce ammonia as one of the products. Based on the electronic structure of the azide ligand and its modification in the state coordinated to Ni(II) ion, it is possible to understand why, in the presence of H+ ions, the reaction pathway leads to ammonia (ammonium ions) and molecular nitrogen as the most stable products. The above mechanism explains the key electrochemical role of the azide ligand in the generation of the multistability in NiðIIÞ–N
3 reduction at a streaming electrode [9].

The mechanism discussed so far explains the formation of the waves (II) + (III) but for a more complete understanding of the electrode process under study it is useful to discuss a slight increase of the limiting charge of the wave (I) in Fig. 3, observed upon addition of HClO4. We think that this phenomenon is a consequence of the direct reduction of N
3 by the Ni(0) (or/ and also Ni(I) species), with the reoxidation to Ni(II), a reaction which, of course, also requires H+ ions. This process appears to occur with a relatively low efficiency, but a small amount of Ni(II), regenerated in this way in the diffusion layer, is sufficient to make the concentration gradient of Ni(II) slightly steeper and thus to ensure the appropriately faster transport of this species to the electrode surface. The most convincing proof of this pathway is the production of ammonia at electrolysis potentials less negative than those corresponding to waves (II, III) (cf. Table 1). One should remember that reactive Raney nickel was also able to reduce azide ions in the acidified media. A quantitative explanation of the amount of NH3 compared to the amount of Ni2+ after the macro-electrolysis is however more complex. In terms of the electrode process (6) the difference between the total charge of the electrolysis and the charge equivalent to the amount of ammonia produced, i.e., the value of [Qel
Q(NH3)] should always be equal to the charge equivalent to the net amount of reduced Ni2+. This conclusion would remain valid even if the nickel(0) formation were followed by its reoxidation to Ni(II) by azides, since then only the relative proportions between Q(NH3) and [Qel
Q(NH3)] would change, with the latter difference tending asymptotically to zero if the nickel reoxidation by azides occurred with 100% efficiency.

5. Conclusions

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