Kinetics of Sn electrodeposition from Sn(II)-citrate solutions

Article abstract of DOI:10.1016/j.electacta.2008.06.037

The influence of solution chemistry on the electrodeposition of Sn from Sn(II)-citrate solutions is studied. The distribution of various Sn(II)-citrate complexes and citrate ligands is calculated and the results presented as speciation diagrams. At a SnCl₂·H₂O concentration of 0.22 mol/L and citrate concentration from 0.30 mol/L to 0.66 mol/L, SnH₃L⁺ (where L represents the tetravalent citrate ligand) is the main species at pH below about 1.2 and SnHL^- is the main species at pH above about 4. Polarization studies and reduction potential calculations show that the Sn(II)-citrate complexes have similar reduction potentials at a given solution composition and pH. However, the Sn(II)-citrate complexes become more difficult to reduce with higher total citrate concentration and higher solution pH. Nevertheless, SnHL^- which forms at higher pH is a favored Sn(II)-citrate complex for Sn electrodeposition due to better plated film morphology, likely as a result of its slower electroplating kinetics. Precipitates are formed from the Sn(II)-citrate solutions after adding hydrochloric acid (to lower the pH). Compositional and structural analyses indicate that the precipitates may have the formula Sn₂L.

Full text of DOI:10.1016/j.electacta.2008.06.037

Electrochimica Acta 53 (2008) 8332–8340
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Kinetics of Sn electrodeposition from Sn(II)–citrate solutions
Chunfen Han∗, Qi Liu, Douglas G. Ivey
Department of Chemical and Materials Engineering, University of Alberta, Edmonton, Alberta, Canada T6G 2G6
a r t i c l e i n f o
a b s t r a c t
Article history:
Received 24 May 2008
Received in revised form 19 June 2008
Accepted 20 June 2008
Available online 27 June 2008
The influence of solution chemistry on the electrodeposition of Sn from Sn(II)–citrate solutions is studied.
The distribution of various Sn(II)–citrate complexes and citrate ligands is calculated and the results pre-
sented as speciation diagrams. At a SnCl2·H2O concentration of 0.22 mol/L and citrate concentration from
0.30 mol/L to 0.66 mol/L, SnH3L+ (where L represents the tetravalent citrate ligand) is the main species at
pH below about 1.2 and SnHL− is the main species at pH above about 4. Polarization studies and reduction
potential calculations show that the Sn(II)–citrate complexes have similar reduction potentials at a given
solution composition and pH. However, the Sn(II)–citrate complexes become more difficult to reduce with
higher total citrate concentration and higher solution pH. Nevertheless, SnHL− which forms at higher pH
is a favored Sn(II)–citrate complex for Sn electrodeposition due to better plated film morphology, likely
as a result of its slower electroplating kinetics. Precipitates are formed from the Sn(II)–citrate solutions
after adding hydrochloric acid (to lower the pH). Compositional and structural analyses indicate that the
precipitates may have the formula Sn2L.
Keywords:
Sn(II)–citrate complexes
Distribution diagram
Polarization curves
Electrodeposition
Solution chemistry
© 2008 Elsevier Ltd. All rights reserved.
1. Introduction
of sodium carbonate to the solution did not result in any pre-
cipitation. A kinetic model for Sn electrodeposition from citrate
Increasing health and environmental concerns worldwide
regarding the toxicity of Pb, and legislation to prohibit or restrict
the use of Sn–Pb solders, have stimulated the development of dif-
ferent Sn-rich alloys as alternative solders in electronic assemblies
[1–3]. As Sn is the main ingredient in these alloys, studies on Sn
solution chemistry and, in particular, the role of complexing agents
in Sn electrodeposition are beneficial.
Different complexing agents have been investigated and citrate
is one of the most extensively studied ones, particularly in the elec-
trodeposition of pure Cu [4,5] and Cu and Co alloys, such as Cu–Ni
[6–8], Cu–Co [9], Co–Mo [10], and Co–W [11]. Citrate has also been
used in Sn, Sn–Au, Sn–Cu, and Sn–Co electrodeposition [12–15].
Citrate has been shown to be a suitable complexing agent, because
good quality Sn and Sn-rich alloys have been obtained and more
importantly because of its inherent low toxicity. Moreover, citrate
can function as a brightening, leveling, and buffering agent, thus
eliminating the need for other plating bath additives [4,6].
The formation of Sn(II)–citrate complexes in solutions contain-
ing both stannous chloride and ammonium citrate was confirmed
by Hess and Taft [16] in 1938 from the observation that the addition
solutions was proposed by Survila et al. [17] for the codeposi-
tion of Sn and Co. SnL2− and SnHL− (where L represents the
tetravalent citrate ligand and citric acid can be represented as
H4L) were considered as the most likely complexes formed in
acidic solutions. A stability constant of 3.2 × 1019 for SnHL− was
determined experimentally by Survila et al. [17]. Their distribu-
tion diagrams of Sn(II)–citrate complexes showed that at a Sn2+
concentration of 0.05 mol/L and a citrate concentration of either
0.1 mol/L or 0.3 mol/L, SnHL− was the main species between pH
2 and 4, SnL2− was the main species at pH > 4 and Sn2+ was the
dominant species below pH 2. Although several other researchers
have also studied the Sn(II)–citrate system [18–21], there are dif-
ferences among the proposed stability constants and distributions
for Sn(II)–citrate complexes. Jean and Sinicki [19] proposed three
mono-nuclear complexes, SnH3L+, SnH2L, and SnHL−, and calcu-
lated their stability constants, which were 3.2 × 105, 2.5 × 107, and
2.0 × 1010, respectively. Tikhonov and Kurolap [20] used potentio-
metric methods to study the complexes formed between citric acids
and stannous ions and showed that SnHL− formed when the pH was
below 7 and SnH(OH)L2− formed when the pH was greater than or
equal to 7. Elbourne and Buchanan [21] proposed the formation
of two different Sn(II)–citrate complexes in solutions containing
0.4 mol/L citrate and different concentrations of stannous ions.
They proposed that in a solution containing 1 × 10−3 mol/L to
2 × 10−3 mol/L stannous ions a bi-nuclear structure (Sn2(H2L)2) was
formed, where two citrate groups chelated with two stannous ions.
∗
Corresponding author at: Department of Chemical and Materials Engineering,
University of Alberta, 536 CME Building, Edmonton, Alberta, Canada T6G 2G6.
Tel.: +1 780 492 8902; fax: +1 780 492 2881.
E-mail address: chunfen@ualberta.ca (C. Han).
0013-4686/$ – see front matter © 2008 Elsevier Ltd. All rights reserved.
doi:10.1016/j.electacta.2008.06.037

C. Han et al. / Electrochimica Acta 53 (2008) 8332–8340
8333
Table 1
However, in solutions containing higher concentrations of stannous
ions (2.5 × 10−3 mol/L), a mono-nuclear structure (Sn(C6H6O6)2)
was formed, where two citrate groups chelated with one stannous
ion. The stability constants for the complexes were not determined.
One concern about a citrate electroplating solution is its insta-
bility. Although precipitation has occurred in the Sn(II)–citrate
solution, the precipitates have never been characterized. The prin-
cipal objective of this investigation was to study Sn(II)–citrate
solution chemistry, to study the reduction behavior of Sn(II) from
the Sn(II)–citrate solution and to identify precipitates formed in
the Sn(II)–citrate solutions to better understand the kinetics of Sn
electrodeposition from Sn(II)–citrate solutions.
Stability constants for Sn(II)–citrate complex ions and Sn(II)–hydroxyl species (“L”
represents tetravalent citrate ligands)
Equilibrium
log K
5.69
4.35
2.87
Ref.
H+ + HL3− → H2L2−
H+ + H2L2− → H3L−
H+ + H3L− → H4L
Sn2+ + HL3− → SnHL−
[22]
[22]
[22]
19.5
10.3
[17]
[18,19]
Sn2+ + H2L2− → SnH2L
7.4
5.5
[18,19]
[18,19]
[23]
[23]
[23]
[23]
[23]
[23]
Sn2+ + H3L− → SnH3L+
Sn2+ + H2O → SnOH+ + H+
Sn2+ + H2O → Sn(OH)2 + H+
Sn2+ + H2O → Sn(OH)3− + H+
Sn2+ + H2O → Sn2(OH)22+ + (H+)2
Sn2+ + H2O → Sn2(OH)3+ + H+
Sn2+ + H2O → Sn3(OH)42+ + H+
−3.8
−7.8
−17.5
−2.4
−6.66
−5.6
2. Experimental
2.1. Solution preparation
The Sn(II)–citrate solutions were prepared by dissolving tri-
ammonium citrate in deionized water, followed by the dissolution
of SnCl2·2H2O in the citrate solution. The deionized water was
obtained from an Aqua MAXTM-ultra purification system. Solution
pH was adjusted by the addition of HCl.
Table 2
Selected free energy data for the Sn(II)–citrate system at 298 K
◦
Species
ꢀGf (kJ/mol)
Ref.
Citric acid (H4L)
H+ (aq.)
−1236.4
[24]
[24]
[24]
[24]
0
−27.2
0
Sn2+ (aq.)
Sn (white)
2.2. Identification of precipitates
Precipitates were formed in the Sn(II)–citrate solutions when
the pH was decreased. The precipitates were analyzed using wide
angle X-ray diffraction (XRD) with a Rigaku rotating Co anode sys-
tem, operating in continuous scanning mode at a voltage of 40 kV
and a current of 160 mA. X-ray photoelectron spectroscopy (XPS)
was also performed on samples using an AXIS-165 XPS system with
an Al K␣ (1400 eV) X-ray source with a resolution of 0.5 eV. The
binding energies of the C 1s, O 1s, and Sn 3d5/2 levels were detected
and the C 1s (284.65 eV) level was taken as the standard for calibra-
tion. Atomic ratios of carbon to oxygen, carbon to tin and oxygen to
tin (C/O, C/S, O/Sn) were determined from fits to the survey spec-
tra. Citric acid and tri-ammonium citrate were used as reference
specimens.
Two different sets of stability constants for SnHL− were found in
the literature, as shown in Table 1. To determine which stability
constants were applicable to the solutions in this work, calculations
with data from both the literature and experiments were done and
are described below.
From the stability constants for citric acid (Table 1), the following
data can be obtained:
3H+ + HL3− → H4L K = 1012.9
ꢀG◦ = −RT ln K = ꢀGf◦,H L − ꢀG◦
(1)
− 3ꢀGf◦,H
+
3−
f,HL
4
From the above equations and the free energy data in Table 2, the
standard Gibbs free energy of formation for HL3− can be calculated
as −1162.81 kJ/mol.
2.3. Polarization measurements
Speciation calculations were initially performed on a solution
containing 0.22 mol/L SnCl2 and 0.30 mol/L tri-ammonium citrate,
which is labeled as S1 (Table 3). This particular solution was chosen
because 0.30 mol/L of tri-ammonium citrate has been shown to be
the minimum amount required to completely dissolve 0.22 mol/L
SnCl2 [25]. The natural pH of solution S1 is 4.1. At this pH, all
three ammonium ions should fully dissociate from the citrate ions
and each citrate ion can carry up to four protons, thus form-
ing H4L, H3L−, H2L2− and HL3−. Only H3L−, H2L2− and HL3− are
considered as ligand species with Sn2+ ions in solution. Survila
et al. [17] considered both SnHL− and SnL2− in their model for
Sn(II)–citrate complexes. The difference is that they used citric acid
instead of tri-ammonium citrate. One important consideration for
tri-ammonium citrate is that the proton in the ternary hydroxyl
group is difficult to dissociate at a pH of 4.1, so that SnHL− was more
likely to form in solution S1. To simplify the calculation, all Sn2+ ions
were assumed to form SnHL− complexes with the citrate, so that the
Polarization measurements were performed using a Gamry
electrochemical system operating at a scanning rate of 1 mV/s. Plat-
inum was used as both the working and the counter electrodes,
and a saturated calomel electrode (SCE) was used as the reference
electrode.
2.4. Electrodeposition experiments
All electrodeposition was done on Si wafer pieces, metallized
with a 25-nm Ti adhesion layer and a 200-nm Au seed layer. A Dyna-
tronix DuPR 10-0.1-0.3 pulse plating power supply was used for
electrodeposition, with a 2-ms forward on-time and 8-ms forward
off-time. All electrodeposition was done with a current density of
5 mA/cm2 at room temperature for 30 min. The microstructure and
composition of the Sn films plated under different conditions were
characterized using a Hitachi H2700 scanning electron microscope
(SEM), equipped with an ultra thin window energy dispersive X-ray
(EDX) spectrometer.
Table 3
Composition of Sn–citrate solutions used in this study
3. Results and discussion
S1
S2
S3
S4
SnCl2·2H2O (mol/L)
0.22
0.30
0.22
0.33
0.22
0.44
0.22
0.66
In order to determine the distribution of various Sn(II)–citrate
species in solution, it is necessary to know their stability constants.
Tri-ammonium citrate (mol/L)

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C. Han et al. / Electrochimica Acta 53 (2008) 8332–8340
Fig. 1. Polarization curves for Sn–citrate solutions. In all solutions, the concentra-
tion of SnCl2·2H2O was kept at 0.22 mol/L, while the concentration of tri-ammonium
citrate was varied as follows: S1, 0.30 mol/L; S2, 0.33 mol/L; S3, 0.44 mol/L; S4,
0.66 mol/L.
Fig. 2. Sn(II)–citrate speciation diagram as
a function of pH for solution S1
(0.22 mol/L Sn2+ and 0.30 mol/L tri-ammonium citrate). ꢁ is the molar fraction of
Sn(II)–citrate complexes.
concentration of SnHL− would be 0.22 mol/L and the concentration
of HL3− would be 0.08 mol/L. From the polarization measurements
shown in Fig. 1, the reduction of SnHL− started at about −0.64 V ver-
sus SCE (i.e., −0.398 V versus NHE). A further decrease in applied
potential caused an almost linear increase in the current density,
until at about −0.75 V, where a plateau appeared which is the lim-
iting current density for SnHL− reduction in solution S1.
The reduction reaction for SnHL− during the electrodeposition
of Sn can be given as:
calculated with the aid of spreadsheet software (e.g., Microsoft
Excel). To perform a calculation, the concentrations of all possi-
ble Sn(II)–citrate complexes and Sn2+ hydrolysis products were
expressed as a function of the concentration of SnHL−, Sn2+ and/or
the pH of the solution, respectively. Then, mass balance equations
for both Sn and citrate were set up, so that the concentration of
SnHL− could be calculated at a given pH. Subsequently, the con-
centrations of all the other species were calculated.
The two mass balance equations used during the calculations
are:
SnHL− + 2e → Sn + HL3−
(2)
A potential of −0.398 V versus NHE, from the polarization measure-
ments, was used as the reduction potential for reaction (2) in the
following calculation. It follows from the Nernst equation that:
Ccitrate = [H4L] + [H3L−] + [H2L2−] + [HL3−] + [SnH3L+]
+ [SnH2L] + [SnHL−]
(4)
RT
nF
[SnHL−]
E = E◦ +
ln
= −0.398 V versus NHE
[HL3−
]
CSn(II) = [SnH3L+] + [SnH2L] + [SnHL−] + [Sn2+] + [SnOH+]
The standard electrode potential E◦ for reaction (2) can be calcu-
lated as −0.41 V versus NHE. The standard Gibbs free energy change
for this reaction is equal to:
2+
+ [Sn(OH)2] + [Sn(OH)3−] + 2[Sn2(OH)2
]
+ 2[Sn2(OH)3+] + 3[Sn3(OH)4
]
(5)
2+
ꢀG◦ = −nFE◦ = ꢀGf◦,HL + ꢀGf◦,Sn − ꢀG◦
= 79.3 kJ/mol
f,SnHL
The distribution of Sn(II)–citrate complexes (SnH3L+, SnH2L, and
SnHL−), in molar fraction as a function of pH for solution S1, is
shown in Fig. 2. The concentration of Sn2+ ions is in the range of
10−3 mol/L to 10−9 mol/L and concentrations of all the Sn2+ hydrol-
ysis products shown in Eq. (5) are in the range of 10−5 mol/L to
10−40 mol/L and are not shown in the distribution diagram. It can
be noted that at the solution natural pH of 4.1, the main species in
solution S1 is SnHL−, with a very small amount of SnH2L. This result
may explain the observation that when S1 is stored in a sealed bot-
tle for 2 months, a small amount of precipitates appears and settles
to the bottom of the container. Assuming the precipitates are in fact
SnH2L, S1 should not be stable in the pH region where the SnH2L
concentration is high enough to precipitate out.
Hydrochloric acid was added to solution S1 to decrease the pH.
During the pH adjustment, the solution remained clear until the
pH reached about 3.3. When the pH was decreased to about 3.1,
a small amount of precipitates started forming. Once the solution
was stirred for about 30 s, more precipitates formed, and eventually
solution S1 became milky in appearance, with white precipitates
settling on the bottom of the container. During precipitation, the
solution pH increased to about 3.3 and remained at that value.
−
3−
From the above equations and data shown in Table 2, the standard
Gibbs free energy of formation for SnHL− is −1242.11 kJ/mol.
The SnHL− complex forms according to:
Sn2+ + HL3− → SnHL−
(3)
The standard Gibbs free energy change for the formation of SnHL−
is:
ꢀG◦ = −RT ln K = ꢀGf◦,SnHL − ꢀGf◦,HL − ꢀGf◦,Sn
−
3−
2+
Thus:
◦
K = eꢀG /−RT = 1.36 × 109 and log K = 9.1
Comparing this value with the data from the literature, the stability
constant for SnHL− from [18,19] (i.e., log K = 10.3) is much closer to
our results and thus was used in the subsequent calculations.
Other than the formation of Sn(II)–citrate complexes, Sn2+ ion
hydrolysis was also taken into account and the available stability
constants are listed in Table 1. In this study, the equilibrium con-
centrations of the citrate complexes and all other compounds were

C. Han et al. / Electrochimica Acta 53 (2008) 8332–8340
8335
When the pH was lowered to 2.5, more precipitates formed; this
was likely due to the formation of more SnH2L based on Fig. 2. As
more HCl was added, the precipitates started to dissolve and the
solution became almost clear at a pH of 1.1. This disappearance of
precipitates can be explained by the dissolution of SnH2L and the
formation of SnH3L+ as the main Sn(II)–citrate complex (Fig. 2).
The same calculations were carried out for solutions S2, S3,
and S4 (Table 3). For solutions S1, S2, S3, and S4, the concentra-
tion of SnCl2·2H2O was fixed at 0.22 mol/L, while the concentration
of tri-ammonium citrate was varied from 0.30 mol/L to 0.66 mol/L.
The concentrations of Sn(II)–citrate complexes (SnH3L+, SnH2L, and
SnHL−) and citrate species (H4L, H3L−, H2L2−, and HL3−) as a func-
tion of pH are shown in Fig. 3a–d. There is almost no change in the
distribution of the different Sn(II)–citrate complexes as the con-
centration of tri-ammonium citrate is increased from 0.30 mol/L to
0.66 mol/L, but the distribution of the citrate species changes. The
peak positions of the curves for these four citrate species stayed the
same for S1, S2, S3, and S4, but the height of each peak increased
with increasing tri-ammonium citrate concentration. The standard
electrode potentials (E◦) for the reduction of Sn(II) to Sn0 from
SnHL−, SnH2L, and SnH3L+, calculated from Gibbs free energy data,
were −0.69 V, −0.60 V, and −0.54 V versus SCE, respectively. These
results show that SnH3L+ is the easiest to reduce, followed by SnH2L
and then SnHL−. To find out the effect of citrate species distribution
on the reduction behavior of the Sn(II)-complexes, the reduction
potentials (E) for SnHL−, SnH2L, and SnH3L+ were calculated using
the Nernst equation for solutions S1, S2, S3, and S4 and the results
are shown in Fig. 4a. In each solution, the reduction potentials for
the three Sn(II)–citrate complexes were similar, although there is a
slight increase from SnHL− to SnH2L and from SnH2L to SnH3L+.
In addition, as the concentration of tri-ammonium citrate was
increased from 0.30 mol/L to 0.66 mol/L, the reduction potentials
for the Sn(II)–citrate complexes decreased from −0.61 V to about
−0.68 V versus SCE. The calculated and experimental reduction
potentials are shown in Fig. 4b; both follow the same trend. The
values for each Sn(II)–citrate complex are so close that they overlap
one another and thus only one data point is shown for each solution.
The reduction potentials from the polarization measurements were
more negative than the calculated values, because during the polar-
ization measurements, in addition to the overpotential needed to
support the electrochemical reactions, the applied potential also
must cover the ohmic drop required to drive the ionic current in
solution. Since the reduction potential for all three Sn(II)–citrate
complexes were very close to one another, it is also quite possible
that they would participate in the reduction reaction at the same
time, and be reduced to Sn metal on the cathode.
The same calculations for reduction potentials were also car-
ried out for solution S1, with different pH values, and the results
are shown in Fig. 5a. The pH was adjusted in the same way as
before. At each pH value, the reduction potentials for SnHL−, SnH2L,
and SnH3L+ were again similar to one another, with a very slight
increase from SnHL− to SnH2L and then to SnH3L+. As the pH was
Fig. 3. Speciation diagrams as a function of pH for solutions S1 (a), S2 (b), S3 (c), and S4 (d). Note that all solutions contain 0.22 mol/L Sn2+, but different concentrations of
tri-ammonium citrate: 0.30 mol/L (a), 0.33 mol/L (b), 0.44 mol/L (c), and 0.66 mol/L (d).

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C. Han et al. / Electrochimica Acta 53 (2008) 8332–8340
Fig. 4. (a) Calculated reduction potentials for SnHL−, SnH2L, and SnH3L+ for solutions S1, S2, S3, and S4 at natural pH. (b) A comparison of calculated and experimental
reduction potentials as a function of tri-ammonium citrate concentration.
decreased from 4.1 to 2.5 and then to 1.1, the reduction potentials
became less negative, moving from −0.615 V to −0.53 V and then to
−0.48 V versus SCE, which means that the Sn(II)–citrate complexes
became easier to reduce at lower pH.
To help explain the above calculations, polarization measure-
ments were performed for solution S1 with pH values of 1.1, 2.5,
and 4.1, and the results are shown in Fig. 5b. The reduction potential
change as a function of pH showed a similar trend to the calculated
Fig. 5. (a) Calculated reduction potentials for SnHL−, SnH2L, and SnH3L+ for solution S1 (0.22 mol/L Sn2+ and 0.30 mol/L tri-ammonium citrate) at pH 1.1, 2.5, and 4.1. (b)
Polarization curves for solution S1 at pH 1.1, 2.5, and 4.1. (c) A comparison of calculated and experimental reduction potentials for solution S1 as a function of pH.

C. Han et al. / Electrochimica Acta 53 (2008) 8332–8340
8337
results. The current density for solution S1, with a pH of 2.5, was
very low, lower than the values for the other two solutions. This
is due to the formation of a significant quantity of precipitates at
a pH value of 2.5. In addition, both the Sn and citrate concentra-
tions in the solution were much lower. The polarization curve for
solution S1, with a pH of 1.1, showed an almost linear dependence,
due to significant hydrogen evolution at this pH. By comparison,
the calculated and experimental reduction potentials for different
pH values are combined in Fig. 5c. The data follow the same trend,
with the measured reduction potential being more negative for the
reasons stated previously.
Electrodeposition of Sn from solution S1, with pH values of 1.1,
2.5, and 4.1, was carried out on Au-coated Si substrates. SEM sec-
ondary electron (SE) images for both plan views and cross-sections
of the as-deposited Sn films are shown in Fig. 6a–c. The magnifica-
tion for Fig. 6a is 5 times lower than that for Fig. 6b and c. EDX point
analysis at the positions identified in the plan views of Fig. 6a and
b is shown in Fig. 7a and b. EDX spectra from point 1 (the smooth
area) and point 2 (one of the large particles) in Fig. 6a are shown in
Fig. 7a-1 and a-2. The EDX spectrum in Fig. 7a-1 is mostly Au with
a small Sn peak. The Au is from the underlying seed layer and the
Sn is from a thin continuous Sn layer (ꢀ1 ␮m thick), which forms
during the initial nucleation stage from the reduction of SnH3L+,
SnH2L, and SnHL−. The larger particles are pure Sn (Fig. 7a-2). At a
pH of 2.5, a thicker continuous Sn layer (about 0.8 ␮m thick) forms
before isolated Sn particles grow preferentially (Fig. 7b-1 and b-
2). The reason why the continuous Sn layer is thicker at pH 2.5
than at pH 1.1 is that the concentration of SnHL− is much higher
at pH 2.5 than at pH 1.1. This is important because SnHL− has a
slower plating rate and is more likely to form a continuous layer.
Consequently, fewer and smaller Sn particles are formed at this
pH. At the natural pH (solution S1), a dense Sn film forms with
a thickness of about 5 ␮m. All three cross-section samples were
prepared by cleaving with a diamond scribe, and the cross-section
image in Fig. 6c is not perfectly edge-on, so some of the surface is
visible.
Fig. 6. SEM secondary electron (SE) images (plan view and cross-section) of Sn films electrodeposited from solution S1 (0.22 mol/L Sn2+ and 0.30 mol/L tri-ammonium citrate)
at pH 1.1 (a), 2.5 (b), and 4.1 (c).

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C. Han et al. / Electrochimica Acta 53 (2008) 8332–8340
Fig. 8. XRD pattern for precipitates obtained from solution S1 at pH 2.5.
Fig. 9 shows an XPS survey spectrum for the precipitates from
solution S1 (pH 2.5). Several elements, including Sn, C, O, Cl, and
N were detected. There was a small amount of N (about 0.7 at.%)
and Cl (about 0.5 at.%) present in the precipitates. XPS analysis was
done on standard citric acid and tri-ammonium citrate samples
for comparison. Similar N levels (about 0.2 at.%) were detected in
the citric acid specimen and similar Cl levels (about 0.3 at.%) were
detected in the tri-ammonium citrate specimen, and may be due to
contamination.
Fig. 10 shows the C 1s XPS spectra for citric acid, tri-ammonium
citrate and the Sn(II)–citrate solution precipitates. Three peaks
were identified by peak fitting and these correspond to dif-
ferent surroundings for carbon, which match the number of C
bond types in both citric acid and tri-ammonium citrate. The C
1s peaks at 284.65 eV (Table 4) are assigned to methyl carbons
[18,26,27], while the peaks at 286.4 eV (citric acid), 286.53 eV
(tri-ammonium citrate), and 286.27 eV (Sn(II)–citrate precipitates)
correspond to carbon bonded to carbonyl groups [18,27]. The
peaks at 288.78 eV (citric acid), 288.65 eV (tri-ammonium citrate)
and 288.36 eV (Sn(II)–citrate precipitates) are attributed to carbon
atoms of COO groups, which is also in agreement with the binding
energy values reported for acetyl carbons [18,26–28].
Fig. 11 shows the O 1s XPS spectra for citric acid, tri-ammonium
citrate and Sn(II)–citrate precipitates. Two peaks were deconvo-
luted for both citric acid and tri-ammonium citrate, and these also
matched the number of O bonds in both reference specimens. The
high binding energy for O 1s at 533.07 eV (citric acid) and 533.03 eV
Fig. 7. EDX point analysis for Sn plated from solution S1 at pH 1.1 (a-1 and a-2) and
2.5 (b-1 and b-2).
(tri-ammonium citrate) can be assigned to oxygen atoms in
C OH
[27,29] groups and the oxygen peaks at 531.84 eV (citric acid) and
As shown in Fig. 2, at a pH of 1.1, the main complex in the solution
is SnH3L+. With a positive charge, SnH3L+ will be easily attracted
to cathode. In addition, as shown in Fig. 5a, SnH3L+ becomes easier
to reduce at pH 1.1 than at higher pH values. All these contribute
to the high growth rate of Sn on the cathode. In contrast to SnH3L+,
the SnHL− species (which forms at higher pH values) has a nega-
tive charge. The charge distribution within the polar complex ions
must be aligned to facilitate attraction to the cathode, so that mass
transfer will be slower and the deposited film should have sufficient
time to densify.
531.77 eV (tri-ammonium citrate) can be assigned to COO− groups
XRD and XPS were used in an attempt to identify the precipitates
from solution S1, which formed at different pH values, and precipi-
tates from solution S2, S3, and S4 when the pH was decreased. XRD
and XPS analyses indicated that the precipitates had the same struc-
ture and composition. An XRD pattern, from precipitates which had
formed from solution S1 at a pH of 2.5, is shown in Fig. 8. The pattern
is quite complex, with the strongest peak at 2ꢂ = 11.75◦. The pattern
does not match XRD patterns obtained for tri-ammonium citrate,
citric acid, tin chloride dihydrate or tin oxide. We were unable to
index the pattern.
Fig. 9. XPS survey spectrum for precipitates obtained from solution S1 at pH 2.5.

C. Han et al. / Electrochimica Acta 53 (2008) 8332–8340
8339
Table 4
Binding energy for main components of XPS C 1s, O 1s, and Sn 3d5/2 peaks for citric acid, tri-ammonium citrate, and precipitates obtained from solution S1 at pH 2.5 (peak
indexing was made with reference to Refs. [18,26–32])
Binding energy (eV)
XPS peak parameters
Components
C 1s
C
O 1s
COOM
Sn 3d5/2
Sn O
C
C
OH
COO
COOH
C
O
H
O
Citric acid
Tri-ammonium citrate
Sn–citrate precipitate
284.65
284.65
284.65
286.49
286.53
286.27
288.78
288.65
288.36
531.84
533.07
533.03
531.77
531.47
486.8
be attributed to a decrease in the electron density of the carboxyl
oxygen by an increase in the positive environment produced by
Sn(II). Two other Sn(II)–citrate precipitates that formed in solution
S1, with a pH of 2.5, produced similar XPS spectra.
The XPS Sn 3d core level spectrum, with a binding energy
of 486.8 eV, from the precipitates is shown in Fig. 12. It is
difficult to distinguish between O–Sn(II)–O and O–Sn(IV)–O
since the positions of the Sn 3d5/2 peaks for Sn(II) and
Sn(IV) oxy/hydroxy compounds partially overlap: 486.3 0.6 eV
(SnO) and 486.9 0.6 eV (SnO2) [18,30–32]. Considering that tri-
ammonium citrate was not likely an oxidizing agent for stannous
ions in solution, the Sn 3d5/2 line at 486.8 eV could indicate the
presence of a Sn(II) compound. Quantitative analysis, based on the
area estimated from the XPS C 1s, O 1s, and Sn 3d5/2 spectra, was
carried out and the atomic concentrations of C and O in citric acid
and tri-ammonium citrate and C, O and Sn in the Sn(II)–citrate pre-
cipitates are shown in Table 5. The relative concentrations of C/O for
citric acid and tri-ammonium citrate are 1.23 and 1.11, respectively,
which are higher than the ratio of 0.86 calculated from the chem-
ical formulas of citric acid (C6H8O7) and tri-ammonium citrate
((NH4)3C6H5O7). This indicates that some C contamination from
air occurred and the calculated C concentrations are too high. The
same effect would be expected for Sn(II)–citrate precipitates under
the same working conditions, giving higher C/Sn relative concentra-
tions than expected. Therefore, the actual C/Sn atomic ratio in the
Sn(II)–citrate precipitates is lower than 3.86, e.g., 3, 2 or 1. Taking
Fig. 10. XPS C 1s spectrum from: (a) citric acid, (b) tri-ammonium citrate, and (c)
precipitates obtained from solution S1 at pH 2.5.
[26–29]. It should be noted that the binding energies for O 1s in
COOH (531.84 eV) and COO M (531.77 eV) are very close. For
the Sn(II)–citrate precipitates, the O 1s spectrum has only one peak
with a binding energy of 531.47 eV, suggesting equivalent oxygen
atoms and implying equivalency of the two Sn atoms to which
they are bound. Comparison with the reference O 1s peaks in cit-
ric acid and tri-ammonium citrate indicates that the O 1s peak for
into account that there are no C
O H bonds for the O 1s spectrum
for the precipitate (Fig. 11), the Sn2L structure shown in Fig. 13
is proposed, with a C/Sn atomic ratio of 3. Each citrate molecule is
combined with four stannous ions through O bonds, and connected
with another four citrates. Each stannous ion is connected with two
citrates through O bonds.
the precipitates is associated with COOH bonds and/or COO
M
bonds, with no C H bonds. The O 1s binding energy for the
O
Sn(II)–citrate precipitates was lower than that for citric acid and
tri-ammonium citrate by 0.37 eV and 0.3 eV, respectively, which can
Fig. 12. XPS Sn 3d core level spectrum for precipitates obtained from solution S1 at
pH 2.5. The thick solid line is the experimental spectrum; the thin line is the fitted
spectrum.
Fig. 11. XPS O 1s spectrum from: (a) citric acid, (b) tri-ammonium citrate, and (c)
precipitates obtained from solution S1 at pH 2.5.

8340
C. Han et al. / Electrochimica Acta 53 (2008) 8332–8340
Table 5
Composition results for citric acid, tri-ammonium citrate and precipitates obtained from solution S1 at pH 2.5
Element
Atom concentration (%)
Atomic ratio
Atomic ratio
C/Sn
C
O
Sn
C/O (theoretical)
C/O (actual)
O/Sn
3.3
Citric acid
Tri-ammonium citrate
Sn–citrate precipitate
55.0
51.1
46.5
44.8
46.1
40.2
0.86
0.86
0.86
1.2
1.1
1.2
12.0
3.9
Acknowledgments
The authors wish to thank Micralyne, Inc. for providing met-
allized Si wafers. This project was supported financially by the
Natural Sciences and Engineering Research Council (NSERC) of
Canada. The authors are also grateful to the Alberta Centre for Sur-
face Engineering and Science (ACSES) and to A. He and D. Karpuzov,
in particular, for assistance with the XPS analysis, as well as to
graduate student Weifeng Wei for his helpful discussion.
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