Absorption Bands of Cuprous and Cupric Salts in Concentrated Alkali Halide Solutions, and Their Analytical Implications SIR: The characteristic absorbance of many rations in concentrated halide solutions in the ultraviolet is of considerable analytical interest. A number of papcrs havc appeared in recent years proposing the quantitative determination of these ions in the parts per million range (1, 4-6). One of the sprious interferences is the presence of coppcr salts (6),the exact position of the maximum absorption (AmBx) of which is controversial. With respect to the cuprous halides, it has been observed by Fromherz and
Ill
06/--
r
LA_--
200
fairly narrow, having a half width of 28 mfi, the same as reported by Fromherz and Menschik (9). Spectra with exactly the same form and peak position were obtained when reducing a cupric solution. Reduction was accomplished by three different reagents: SOa-*, HzOa, and metallic copper (Figure 1). Hydrogen peroxide reduces the cupric ion in concentrated halide solutions and absorbs in this part of the spectrum. Therefore, the solutions were boiled t o decompose it, before recording the spectrum. Distinct from other cation complexes, Amax of the cuprous halide complex is only very slightly influenced by the nature of thc concentrated halide 80111-
Menschik (3) that Amsx in concentrated potassium chloride solutions is a t 272 mp, while in concentrated potassium bromide solutions it is only slightly shifted to longer wave lcngths, A,,,%* = 276 mp. I n Figure 1 are shown the absorption spectra of cuprous chloride in sodium chloride and in potassium chloride solution, both being 4 M in halogen ion (alkali halide solutions of the same concentration were used with all expcriments reported here). Both spectra peak a t 272 mp, and the hands are
220
L_L---
i
260
280
240
WAVE
-L \ E 300
d
320
L E N G T H IN
340
-
1
--
d 360
3R0
400
mp WAVE
Figure 1 ,
Absorption spectra of Cu(l) solutions
{For sake of clarity, scale on ordinate w a i moved for each spectrum as lndlcated In bracketi) I. 8 p.p.m. Cu(l) In 4N NaCl II. 8 p.p.m. Cu(l) In 4N KCI (+0.1) 111, 8 p.p.m. Cu(l1) 100 p,p.m. SOI-' In 4N KCI (+0.2) 50 p.p.m. HzOs In 4N KCI (4-0.2) IV. 1 0 p.p.m. Cu(l1) V. Cu metal dissolved b y bolllng In 4N KCI (rlght hand icale) 10 p.p.m. VI. Cu metal dissolved b y bolllng In 4N KCI Cu(ll) (rlght hand scale, -0.2)
++
+
1122
ANALYTICAL CHEMISTRY
Figure 2. I. 11. Ill. IV.
V. VI.
LENGTH
IN mp
Absorption spectra of Cu(ll) solutions
10 p.p.m. Cu(1I) in 4N KCI 1 6 p.p.m. Cu(l1) In 4N NaCl Trace of CI, In 4N KCI 19 p.p.m. Cu(i1) f CIS In 4N KCI (rlght hand scale) 19 p.p.m. Cu(ll) CI, In 4N KCI after bolllng (add 0.16 to scale on left) 10 p.p.m. Cu(ll) stream of Nz In 4N KCI 20 p.p.m, Cu(l) Clt In 4N KCI (double scale unlts on left)
+ ++
tions, being always situated a t the same wave length. The spectra of the cupric salts in concentratcd halide solutions are far more complicatcd. In Figure 2 are given two broad absorption Rpectra in potassium and sodium chloride solutions. They penk a t 258 and 255 mp, respectively. [It appears that the spectrum shown in Figure 11 of (6) and indicated to be that of Cu(1) in 6 M hydrochloric acid, is really that of a Cu(I1) solution.] The Amsx of CuBr2 in KBr solutions was found by Fromherz et al. (3) to be at 293 mp, and in a later publication (9)a t 281 mp. The latter also observed one or two additional absorption bands in the longer wave length range, in highly concentrated lithium bromide solutions, and noted that the cupric salt decomposes to a large extent (about 19%) into the cuprous complex and bromine. These observations offer an explanation for the inconsistency in the stated peak positions, as well as for the exceptional broadness of the absorption spectra. It is suggested by the present authors that the broad bands obtained with the cupric halides result from the overlap of three distinct spectra-namely (with chloride solutions), the cupric complex at h,.”250 mp; the cuprous complex at A b . = 272 m l ; and the long wave length tail of chlorine, peaking in KCl solutions at 234 mp (corresponding to the
Vrband in crystalline KC1) (curve 111, Figure 2). The equilibrium: C U + ~(complex)+ Cu+ (complex) leans to the right in the alkali halide solutions studied according to the order: HCl < NaCl < KC1 < Kbr. Further proof for the above proposition has been obtained by trying to shift the equilibrium one way or the other, as follows : When chlorine gas is bubbled through halide solutions of Cu(II), the copper band is entirely hidden under the strong absorbance of the chlorine (curve IV, Figure 2). The excess of chlorine may be removed by careful boiling, whereupon the copper band, now peaking a t shorter wave lengths, reappears. The spectrum thus obtained (curve V) peaks a t 251 mp, i t is appreciably narrower than that of the untreated Cu(I1) solutions, and, in common with curve 111,does not show the rising tail of the second absorption band of Cu(1) at shorter wave lengths. (Solutions containing excess chlorine are sensitive to ultraviolet light, and their absorbance diminishes during a measurement.) When nitrogen gas is passed through a Cu(I1) solution, the absorption peak, as well as the minimum aksoktion at shorter wave lengths, are gradually shifted to longer wave lengths, clearly, owing to the equilibrium shift toward
the increased formation of Cu(1) (curve
VI). Last, when chlorine was bubbled through a Cu(1) solution and then boiled (to remove excess), absorption spectra identical in form and peak position to the Cu(I1) solutions were obtained. This is shown by dots close to curve I. Figure 2 also shows the minimum absorbance a t 220 mp of Cu(1I) solutions in comparison to a minimum appearing a t about 245 mp with Cu(1) solutions. LITERATURE CITED
(1) De Sesa. M. A.. Rogers. L. B.. Anal.
.-,(B) 3 , l (1929).
.
.
’
(4) Kanzelmeyer J., Freund, H., ANAL. CHEW 25, 18& (1953).
(6) Merritt, C., Jr:, Hershenson, H. M., Rogera, L. B., Zbzd., 25,572 (1953). (6) Yamamoto, Y., Bull. Inst. Chem. Research, Kyoto Unav. 36, No. 6, 139 (1958).
ABRARAM GLASNER PINCHAS AVINUR Department of Inorganic and Anal tical Chemistry The Sebrew Univeralty Jerusalem, Israel RECEIVEDfor review March 27, 1961. Acce ted A ril 11, 1961. Taken from the $h.D. tfesis to be submitted b P. Avinur to the Senate of The Hetrew University.
Adsorption as a Source of Inconstancy of the Chronopotentiometric Constant for Short Transition Times SIR: Recently, Bard (4) presented experimental data demonstrating that the chronopotentiometric constant, ir1I2/C, increases markedly at short transition times for a number of reactions at platinum electrodes. To account for this increase in the values of i9l2/C, Bard suggested that the short (10-l to lo-* second) transition t i e s were lengthened relative to longer (1 to 10 seconds) transition times as a result of oxidation of the electrode, or the time necessary to charge the double layer, or because the effective area of the electrode increased at short transition times because of the surface roughness of the electrode. This communication offers evidence that a major catme of the increases in if1‘2/Cobserved by Bard was adsorption of the species undergoing the electrode reaction on the surface of the electrode. Experimental values of i~l‘z/C for the oxidation of iodide ion in 1F Ha04 are presented by Bard for 10, 6, 1, and O.lmM solutions of iodide ion. I n the
extends into the solution from the case of iodide ion, oxidation of the surface of a chronopotentiometric workelectrode can be ruled out aa a conis approximately equal to tributing factor to the increases in i ~ ~ / ing ~ electrode / (Dt)1/2,where D is the diffusion coC because the transition times were efficient of the species that is reacting a t measured at a potential (0.7 volt us. the electrode and t is the time since the S.C.E.) where no appreciable oxidation electrolysis was started. The roughness of the platinum electrode could have of the electrode surface will begin to occurred. Qualitatively, t.he deviations from cause increases in the effective area of constancy in the chronopotentiometric the electrode when distance (Dt)’/* becomes commensurate with the average constant resulting from double layer distances from “hilltops” to “valleys” charging are very similar to deviations on the electrode surface. The possicaused by reactant adsorption (vide infra) and it is not always n simple bility of contributions from roughness matter to distinguish between the two causing increases in W/Z/C is more likely as T decreases; however, for a effects. However, Bard calculated that under the conditions he employed the constant value of T, i ~ 1 ” / C should be observed increases in i ~ l / ~ /were C much independent of concentration because (Dt)l I 2 is independent of concentration. too large to be even approximately accounted for by contributions to T Thus, chronopotentiograms obtaiied from the charging of the double layer. for solutions of varying concentration Electrode roughness can also be ruled with the current density adjusted in out as a major source of the increases in each case to keep T constant should give ~ T ~ / ~by/ Cmeans of the following values of i~l’q/C that are independent argument, which is due to Lorenz (6): of the concentration even for rough The distance that the diffusion layer electrodes. Figure 3 in Bard’s paper VOL. 33, NO. 8, JULY 1961 e
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