Erratum: Formation of Cu Nanoparticles by Electroless Deposition Using Aqueous
CuO Suspension [J. Electrochem. Soc., 155, D474 (2008)]
Shunsuke Yagi∗,z , Hidetaka Nakanishi∗ , Eiichiro Matsubara∗ , Seijiro
Matsubaraa , Tetsu Ichitsubo∗ , Kazuo Hosoyab , Yorishige Matsubab
∗
Department of Materials Science and Engineering,
Kyoto University，Kyoto 606-8501, Japan
a
Department of Material Chemistry,
Kyoto University, Kyoto 615-8510, Japan
b
Tsukuba Research Laboratory,
Harima Chemicals, Inc., Ibaraki 300-2635, Japan
z
Corresponding author. Email: [email protected]
However, pH normally changes with temperature in
an aqueous solution because the ionic product for water, K w = [H+ ][OH− ], changes with temperature. In the
present method, the pH of the reaction suspension was
adjusted to 12.0 using NaOH aqueous solution at 298 K,
and pH values at higher temperatures must differ from
the initial value 12.0 at 298 K. Figure 2 shows the calculation and experimental results of the change in pH
with temperature. The calculation result was obtained
on the assumption that the concentration of OH− ions is
almost constant with temperature because most of OH−
ions arise from the dissociation of strongly basic NaOH.
The calculation result agrees well with the experimental
results below about 330 K, but the difference becomes
large toward high temperature above about 330 K. The
extra change in pH seen above about 330 K does not
seem to be accounted for by considering only the main
reactions in this system (i.e. Cu deposition, hydrogen
generation, hydrazine oxidation, and CuO dissolution)
as written in the original paper at p.D477, l.19 from the
bottom in the right column. In our preliminary exper-
16
14
−0.15
12
10
(b)
−0.20
8
6
−0.25
4
(a)
2
0
290
Oxidation-reduction potential / V vs SHE
In the original paper, the activity of Cu2+ aquo ions
and the oxidation-reduction (redox) potential of the
Cu2+ /Cu redox pair were calculated as shown in fig. 5
at p.D477 in the original paper. However, some thermodynamic data (the standard entropy of hydrogen and
specific heats of hydrogen and proton) were not included
for the calculation of the redox potential of the Cu2+ /Cu
redox pair. Compensating for the deficient data,1,3 the
precise redox potential at a constant pH 12 can be determined as in fig. 1. Table I shows all the data considered
and used for the calculation in the original paper and the
present erratum.
iment, such an extra pH change with temperature was
also observed in a hydrazine aqueous solution. Thus, it
can be concluded that the extra pH change is caused by
the decrease in the equilibrium concentration of protonated hydrazine N2 H5 + due to the decrease in the total
amount of hydrazine species by decomposition or oxidization during temperature increase. Actually, the amount
of N2 H5 + is calculated to be ca. 5.1×10−6 at pH 12,
298 K, and [N2 H4 ]total =0.5 M, considering N2 H4 , N2 H5 + ,
and N2 H6 2+ as hydrazine species, which is sufficient to
change pH from 12 to around 10.
Considering the pH change with temperature according to the experimental results and calclulation, the activity of Cu2+ aquo ions and the redox potential of the
Cu2+ /Cu redox pair are calculated as shown in fig. 3.
In the calculation result, the activity of Cu2+ aquo ions
linearly increases with temperature and the redox potential of the Cu2+ /Cu redox pair monotonically decreases
with temperature. This tendency is the same below 330
K in the experimental result. However, the activity of
Cu2+ aquo ions exponentially increases with temperature
above 330 K, and the redox potential of the Cu2+ /Cu redox pair starts to increase with temperature above about
340 K. In the present experiment, the immersion poten-
Activity of Cu2+ ions aCu2+ / 10−17
The paper “Formation of Cu Nanoparticles by Electroless Deposition Using Aqueous CuO Suspension” was
published in Journal of the Electrochemical Society, 155,
D474 (2008). The authors found some errors and misleading expressions in the paper. These, however, do not
affect the experimental results, the outline of the discussion, or the conclusion. We thank Mr. Shohei Shiomi
from Kyoto University for his help in conducting the experiments for this erratum. Corrections and supplemental remarks follow.
−0.30
300
310 320 330 340
Temperature / K
350
360
FIG. 1: (a) Activity of Cu2+ ions and (b) redox potential of
Cu2+ /Cu redox pair calculated from thermodynamic data.
2
0
Potenti
al/ V vsSHE
12.0
11.5
pH
Experimental result
11.0
Calculation result
10.0
300
310
320
330
340
Temperature /K
350
360
FIG. 2: Calculation and experimental results of the change
in pH with temperature.
tial measured in the reaction suspension at 353 K is still
lower than both the redox potentials obtained only by
calculation (−0.19 V vs SHE) and by calculation in combination with the experimental result (−0.15 V vs SHE),
which does not affect the outline of the disscusion. Consequently, fig. 9 in the original paper is modified using
the value obtained only by calclulation, which is a severer value (−0.19 V) for Cu deposition, as fig. 4 in the
erratum. In addition, at p.D479, l.11 from the bottom in
the left column, the redox potential should be −0.16 V
and −0.19 V vs SHE without the effect of the decrease
in pH by the hydrazine consumption at 323 K and 353
K, respectively. It should be noted that the initial pH of
the reaction suspension is 12.0 at 298 K, but is no longer
12.0 at 323 K and 353 K.
In the original paper, Cu2+ aquo ion was only considered as Cu(II) ionic species dissolved in the solution
to determine the redox potential of the Cu2+ /Cu redox
pair. This was for simplicity because the redox poten−0.12
Experi
mentalresul
t
Acti
vi
tyofCu2+ i
onsaCu2+ /10−14
7
6
5
4
3
2
1
Cal
cul
ati
on resul
t
310
320
330
340
Temperature /K
Oxi
dati
on-reducti
on potenti
al/V vsSHE
(a)
8
300
−0.4
−0.6
0
experimental result
0
ECu2+/Cu
Immersi
on potenti
al
20
40
60
Ti
me / mi
n
80
100
120
Fitting curve of
10.5
9
−0.19
−0.2
(b)
Experi
mentalresul
t
−0.14
−0.16
−0.18
Cal
cul
ati
on resul
t
−0.20
350
360
300
310
320
330
340
Temperature /K
350
360
FIG. 3: (a) Activity of Cu2+ ions and (b) oxidation-reduction
potential of Cu2+ /Cu redox pair calculated from thermodynamic data considering calculation and experimental results
of the change in pH with temperature.
FIG. 4: Immersion potential of a QCM substrate during
liquid-phase reduction and the oxidation-reduction potential
of Cu(II)/Cu redox pairs at 353 K for 2h.
tials of any Cu(II) ionic species in equilibrium and metal
Cu are the same at constant pH and temperature in the
presence of abundant solid CuO powder. However, other
Cu(II) ionic species, e.g. HCuO2 − or Cu(OH)2 − and
CuO2 2− or Cu(OH)3 2− , in fact exist in the solution.5,6
Thus, the solubility of CuO is not always equal to the
equilibrium concentration or activity of Cu2+ aquo ions,
and the expression “the solubility of CuO (particles)” in
the original paper should be further specified as “the solubility of CuO (particles) as Cu2+ aquo ions” or revised
as “the activity of Cu2+ aquo ions”. For example, the solubility of CuO at 298 K can be determined as fig 5 in the
erratum considering Cu2+ aquo ions, HCuO2 − ions, and
CuO2 2− ions as Cu(II) ionic species in the solution, and
the lowest solubility (2.3 × 10−10 mol dm−3 ) is achieved
at pH 9.
In their introduction in the original paper, the authors
mentioned that “Muramatsu et al. fabricated Cu2 O
nanoparticles using hydrazine from an aqueous CuO suspension of pH 9.3, where E N2 /N2 H4 and E N2 -NH3 /N2 H4 are
−0.88 and −2.13 V vs SHE at 298 K, respectively. Because these potentials are also low enough to reduce Cu2+
ions to Cu metal, the Cu2 O particles they obtained were
kinetically stabilized.”. This may be ambiguous or misleading. Accurately, the most stable chemical species in
solution can be “thermodynamically” determined by potential at a pH and temperature. However, in electroless
deposition systems, the mixed potential in the solution
is “kinetically” determined at the value where the total
of anodic currents I a,total , balances the total of cathodic
currents I c,total . In the present work, the mixed potential
in the reaction solution or suspension was determined by
measuring the potential of an immersed gold-sputtered
QCM substrate. The measured value was called the “immersion potential”, which was assumed to be almost the
same as the value of mixed potential in the paper. More
rigorous discussion was reported by the authors.7
Trivial typing errors are also found:“share modulus”
should be “shear modulus” at p.D475, l.1 in the right
column. “358 K” should be “353 K” at p.D477, l.18
from the bottom in the right column.
3
8
log[aCu2+ + aHCuO2− + aCuO22−]
6
4
2
0
−2
−4
−6
Cu2+
CuO22−
−8
−10
HCuO2−
0
2
4
6
8
pH
10
12
14
16
FIG. 5: Influence of pH on the solubility of CuO at 298 K considering Cu2+ , HCuO2 − , and CuO2 2− as Cu(II) ionic species.
1
2
3
4
C. M. Criss and J. W. Cobble, J. Am. Chem. Soc., 86, 5385-5393
(1964).
W. M. Latimer, The Oxidation States of the Elements and their
Potentials in Aqueous Solutions, 2nd ed., Prentice-hall, INC. Englewood Cliffs, N. J. (1959).
O. Kubaschewski and C. B. Alcock, Metallurgical Thermochemistry, 5th Edition Revised and Enlarged, Elsevier, NY (1979).
D. R. Stull and H. Prophet, JANAF Thermochemical Tables,
5
6
7
2nd ed., NSRDS-NBS, Washington, DC (1971).
B. Beverskog and I. Puigdomenech, J. Electrochem. Soc.,
144(10), 3476-3483 (1997).
M. Pourbaix, Atlas of Electrochemical Equilibria in Aqueous Solutions, Cebelcor, Brüssel, 384 (1966).
S. Yagi, H. Nakanishi, T. Ichitsubo, and E. Matsubara, J. Electrochem. Soc., 156(8), D321-D325 (2009).
4
Table I. List of standard free energy of formation, heat of formation, entropy at 298 K and 1 atm, and specific heat at constant pressure (1 atm) considered and used
for thermodynamic calculation.
Chemical species
H+ (aq)
−
OH (aq)
Cu2+ (aq)
HCuO2− (aq)
Standard free energy
of formation
Standard heat of formation
Standard entropy
0.0
0.0
0.0
a
129.7
−157.4
−230.0
−10.5
−242.8
(1), (2)
65.0
64.4
267.8
(1), (2)
b
c
References
(1), (2)
CuO22− (aq)
−257.0
−182.0
(2)
(2)
(2)
(2)
NH3 (aq)
−26.6
N2H4 (aq)
127.9
N2H5+ (aq)
N2H62+ (aq)
87.9
94.2
H2O (l)
0.0
70.0
75.4
130.6
27.3
(2), (3)
3.3
0.0
Cu (s)
CuO (s)
(2)
−237.3
H2 (g)
N2 (g)
(2)
(3)
(2)
0.0
−127.2
0.50
33.1
22.6
6.3
42.7
38.8
20.1
(3)
(2), (3)