J Solution Chem (2011) 40: 247–260
Polarographic Behavior of Manganese(II) in the Presence
of Oxalate Ions in Perchlorate and Sulfate Solutions
Jadwiga Urba´ nska
Received: 3 November 2009 / Accepted: 7 July 2010 / Published online: 20 January 2011
© The Author(s) 2011. This article is published with open access at Springerlink.com
Abstract The dc polarographic method has been applied to study coordination equilibria
between Mn(II) and oxalate ions in perchlorate and sulfate solutions. The stoichiometries of
complexes formedinsolutionandthosereducedatadroppingmercuryelectrode wereestab-
lished. The stability constants of the Mn(II) oxalate and sulfate complexes, as well as their
diffusion coefficients, were determined at a constant ionic strength 0.5 mol·L−1and 25°C.
The stabilities of these Mn(II) complexes were compared with the corresponding complexes
of other divalent metal ions. The polarographic method was able to identify complexes that
have not been established by other methods and to determine their stability constants with
Keywords Polarography · Manganese(II) complexes · Oxalate complexes · Stability
constants · Electrode reduction mechanism
Manganese is among the metals most frequently used in the industry, which results in el-
evated levels of this metal in the environment: in water, soil and in biological materials.
Manganese is involved in the production of steel (as a hardening agent), and in Mn alloys,
dry-cell batteries, and in the chemical industry for coloring glasses, ceramics and pigments,
as well as in agriculture as a fertilizer and fungicide (e.g., MANEB, ZINEB). In gasoline, it
is applied as anti-knocking agent and to improve octane ratings (MMT). Recently, its com-
pounds have been used as the contrasting agent in nuclear magnetic resonance tomography
(Mn-DPDP for liver and pancreas scans) [1, 2].
On the other hand, manganese is important in the activation of many enzymes involved
in metabolic processes of all organisms. It is required for protein and fat metabolism, for
healthy nerves and a healthy immune system, for blood sugar regulation, the production of
cellular energy and bone growth [1, 3, 4]. Manganese elevates the level of anti-oxidative
J. Urba´ nska (?)
Faculty of Chemistry, University of Wrocław, 14 F. Joliot-Curie, 50-383 Wrocław, Poland
248 J Solution Chem (2011) 40: 247–260
protection by decreasing the concentration of free radicals , assists in the utilization of
vitamins B1and E, and prevents clotting effects .
However, increased manganese levels are known to damage the central nervous system,
resulting in the motor coordination abnormalities and psychic disorder and, finally, can even
result in symptoms similar to Parkinson’s disease [1, 4, and references cited therein]. Man-
ganese toxicity is also a serious constraint to crop cultivation since manganese is taken up
by plants and can easily be passed into the food chain, again causing symptoms similar to
Parkinson’s disease .
The exact mechanism of toxicity of manganese is not fully understood. To understand
better the role of this element in biological reactions, it is necessary to investigate its inter-
action with various inorganic and organic ligands occurring in the environment, in order to
identify which species influence its bioavailability, mobility and toxicity.
A survey of the literature reveals that manganese speciation, in spite of its biological im-
portance, has been rarely investigated [6–9] in comparison to studies of other 3d-electron
elements. The polarographic method has seldom been used in manganese coordination equi-
libria studies [10–13] although, in some favorable cases, it can detect complexes that have
not been established by other methods and can be used to determine their stability constants
with high accuracy or, at least, comparable to the accuracy achieved by the potentiometric
method [14, 15].
Mn(II) is reduced on a mercury electrode at strongly negative potentials, and in acidic
media its reduction wave (or that of its complexes) overlaps with the hydrogen ion wave, and
thus it is extremely difficult to separate them. On the other hand, in alkaline media Mn(II)
precipitates. Secondly, the reduction wave of Mn(II) to Mn(0) in some media has a slope
higher than that for a fully reversible 2-electron process [10, 16], and it is uncertain which
equation should be used to determine the stability constants of Mn(II) complexes, that for
a reversible reduction or that for an irreversible one. Application of both equations yields
The purpose of the present work is study Mn(II) speciation that may help in recognizing
its beneficial or toxic role in living organisms. The main goal was to demonstrate a simple
method of analysis of Mn(II) waves that provides valuable information concerning the elec-
trode reduction mechanism and to determine stability constants of the complexes. As model
systems, the reduction of Mn(II) in the presence of oxalate ions in perchlorate and sulfate
solutions have been chosen. The discharge of Mn(II) in the presence of oxalate ions in per-
chlorate solutions was partly studied by Verdier and Piro [10, 11]. However, the conclusions
reported in that study need to be revised since their detailed analysis of the current-potential
curves for the reduction mechanism differed from that proposed in the present paper. The
results obtained are presented and discussed below.
Measurements were made on a Radelkis OH-105 polarograph using a dropping mercury
electrode. The mercury container was placed at a height of 50 cm, the mercury flow rate was
2.36 mg·s−1, and the resulting drop time (t1) was 3.8 s. A saturated calomel electrode (SCE)
with a large surface area was used as both anode and reference electrode. It was connected
to the examined solution with an electrolytic bridge, filled with saturated sodium perchlorate
(or sulphate) solution.
All stock solutions were prepared from analytical-reagent grade chemicals. The concen-
trations of manganese(II) in MnSO4solutions was fixed at 5 × 10−5mol·L−1while those
J Solution Chem (2011) 40: 247–260 249
of oxalate (Na2C2O4) were varied over the range 0.001–0.1500 mol·L−1. A constant ionic
strength equal to 0.5 mol·L−1was maintained by adding the necessary amounts of sodium
perchlorate or sodium sulfate. The pH values of all investigated solutions were in the range
6.0–6.5. The pH values were checked by means of a Radelkis OP-211 digital pH meter.
All experiments were carried out under an argon atmosphere at 25°C. The current of
interest was recorded as the difference between the total current and the residual current
measured in the pure supporting electrolyte and ligand solution at the same potential.
3 Results and Discussion
The wave shifted towards more negative potentials as the oxalate ions concentration in-
creased and its limiting current decreased in both supporting electrolytes (sodium perchlo-
rate and sodium sulfate), see Fig. 2, indicating the formation of complexes. Logarithmic
analyses were performed for all of the waves to determine the half-wave potentials (E1/2)
and their slope coefficients (2.303RT/αnαF). Some of the curves, characteristic of this in-
vestigation, are shown in Fig. 3.
Plots of log10[i/(il− i)] versus E for the reduction of the free Mn(II) ion (aqua ion)
and of its oxalate complexes formed at concentrations of oxalate lower than 0.03 mol·L−1,
in both supporting electrolytes NaClO4and Na2SO4, were straight lines with a slope of
33 mV, which is a little higher than that for a fully reversible two-electron process. When
the concentration of oxalate ions exceeded 0.03 mol·L−1, the slope in the lower part of
the wave was still 33 mV, whereas it increased in the upper part. The half-wave potentials
E1/2for these waves were obtained by extrapolation of the lower sections of each wave
(decreased by 33 mV) to zero value of log10[i/(il−i)]. The values of E1/2obtained plotted
against log10coxare presented in Fig. 4.
Fig. 1 Selected polarographic waves for Mn(II)–oxalate complexes in 0.5 mol·L−1NaClO4
250 J Solution Chem (2011) 40: 247–260
Fig. 2 Dependence of the limiting current on log10coxin NaClO4(1) and in Na2SO4(2) solutions
Fig. 3 Selected logarithmic wave analysis for Mn(II)–oxalate complexes in 0.5 mol·L−1NaClO4cox: 0 (1),
0.001 (2), 0.003 (3), 0.008 (4), 0.015 (5), 0.03 (6), 0.06 (7), 0.10 (8), and 0.15 (9) mol·L−1
This method of determination of the half-wave potential of the wave was previously
applied in the Ni(II)–oxalate system , and the stability constants of the complexes deter-
mined from the shift of this potential with respect to the oxalate ions concentration produced
values very close to those obtained in a different way by Crow .
J Solution Chem (2011) 40: 247–260 251
Fig. 4 Changes in the half-wave potentials of Mn(II) with log10coxin NaClO4(1 and 3) and in Na2SO4
(2 and 3) solutions; for curves 1 and 2, the E1/2values were obtained from the wave with slope of 33 mV,
whereas the values for curve 3 were taken directly from the logarithmic curve
Agreement of the slope for reduction of the free Mn(II) ion in perchlorate solution with
that for reduction of Mn(II) oxalate complexes indicated a common character of the elec-
trode process mechanisms in both electrolytes, i.e. reduction of the Mn(II) aqua ion .
However, following the Brönsted theory, the perchlorate ion cannot form complexes with
Mn(II) but the sulfate ion has same tendency for complexation. The results for these sys-
tems are discussed separately below.
3.1 Reduction of Mn(II) in NaClO4Solutions
To determine the stoichiometry of Mn(II)–oxalate complexes formed in solution, two tests
wereperformed:plotting ?E1/2versus log10coxandplottinglog10[i/(il−i)] versus log10cox
at a fixed potential E. The slope of both plots gave the difference in the average ligand
numbers between the complexes predominating in the solution and those that were directly
reduced at the mercury electrode .
The limiting slope of the plot ?E1/2versus log10coxwas 2.32 for (E1/2)cextrapolated
from the lower part of the wave, and 2.51 for (E1/2)ctaken directly from the logarithmic
curve. In the sulfate solutions these values were 2.39 and 2.50 respectively.
The diagram of log10[i/(il− i)] versus log10coxplotted at the three potentials −1.53,
−1.56 and −1.59 V (Fig. 5), at oxalate concentration higher than 0.03 mol·L−1, gave a
difference in the average ligand numbers of 2.40. These results evidently confirm that, in
this system, only the free Mn(II) aqua ion undergoes direct reduction at the mercury elec-
trode, whereas in the solutions there exist three oxalate complexes: Mn(ox), Mn(ox)2−
then the difference of the average ligand numbers should be less than two. These results dif-
fer from those obtained by Verdier and Piro . They have stated that Mn(II) formed only
two oxalate complexes, i.e. Mn(ox) and Mn(ox)2−
ion, were reduced irreversibly at the mercury electrode.
3. When any other species besides of the free Mn(II) ion is reduced at the electrode,
2, and that both, together with Mn(II) aqua
252 J Solution Chem (2011) 40: 247–260
Fig. 5 Variation of log10[i/(il− i)] with log10coxat fixed potentials of: −1.53 V (1), −1.56 V (2), and
−1.59 V (3) in NaClO4solutions
The overall stability constants of these complexes were calculated from the relation given
by Biernat , based on a modification of the DeFord and Hume equation , which
consists in replacing n by αnα. Thus, the modified De Ford and Hume equation can be
where F0(cox) is a polynomial function representing the sum of βjcj
is the overall stability constant of the j-th complex, coxis the oxalate ion concentration, αnα
is the transfer coefficient of the electrode reaction (α < 1), E1/2is the half-wave potential,
and il is the limiting current of the wave (the subscripts s and c refer to the simple and
complex ions, respectively). R, T and F have their usual meanings.
The log10[(il)s/(il)c] term in the F0(cox) function was neglected in calculations at high
concentrations of oxalate ions since (E1/2)cwas extrapolated from the lower part of the
Stability constants of complexes were obtained as follows:
where β0= 1. The function F1(cox) is also introduced:
F1(cox) = (F0(cox)−1)/cox= β1+β2cox+β3c2
A plot of F1(cox) values against the oxalate ion concentration yields a curve from which
values of βj were determined numerically using a commercial computer program in con-
junction with the non-linear least-square regression method. The stability constants obtained
are: β1= 102.65, β2= 104.54and β3= 105.38. All constants were determined with errors not
exceeding 0.02 logarithmic units. The use of the unmodified form of the De Ford-Hume
oxfor all complexes, βj
J Solution Chem (2011) 40: 247–260 253
Table 1 Published stability constants for Mn(II)–oxalate complexes
Glass electrode0.1 KNO3
Polarographic 0.1 NaClO4
Present work 5.38
aCalculated assuming a fully reversible 2e−process
Fig. 6 Distribution of Mn(II) among different oxalate forms as a function of log10coxin NaClO4solutions
equation (i.e., assuming a fully reversible process) gave somewhat higher values of βjrel-
ative to those obtained from the modified form: β1= 102.72, β2= 104.67and β3= 105.91.
Stability constants reported by other authors are summarized in Table 1. The distribution of
Mn(II) among various species is presented in Fig. 6 as a function of log10cox.
The majority of literature studies report the stability constant only for the first oxalate
complex [22–27]. Generally, the value of log10β1obtained herewith is 1.1–1.3 units lower
than the corresponding values obtained by other methods, both at an ionic strength of
254 J Solution Chem (2011) 40: 247–260
0.1 mol·L−1or extrapolated to zero concentration, but is approximately equal to that ob-
tained polarographically by Verdier and Piro [10, 11]. In turn, the log10β2value displays
good concordance with Verdier and Piro’s values obtained at ionic strengths of 0.1 and
0.4 mol·L−1. No data were found in the literature for the stability constant of the third com-
Increasing the oxalate ion concentration is accompanied by a decrease in the limiting
current of the waves. At a concentration of 0.15 mol·L−1, its value was 67.3% of the limiting
current of the Mn(II) wave in the absence of oxalate ions, but it is still diffusion controlled.
This is illustrated by the linear dependence of ilwith h1/2.
From examining changes in the distribution diagram for complexes with changes in the
oxalate ion concentration (Fig. 6), it follows that at an oxalate concentration of 0.15 mol·L−1
there is 48.1% Mn(ox)4−
ically, and the limiting current was dependent only upon the concentrations of the remaining
forms in solution.
According to the basic equation of Ilkoviˇ c, the intensity of the diffusion limiting current
at any given condition is proportional to the concentration of the reduced component :
3present. This suggests that this complex was inactive polarograph-
il= k ·c
k = 6.3nFD1/2
For a constant difference of mercury levels between the container and capillary, the k value
depends only upon the diffusion coefficient. Thus, measurement of the limiting current of
the wave at an appropriate concentration of oxalate ions allows determination of the value
of average diffusion coefficient of complexes, Dav, whose values are presented in Table 2.
The average diffusion coefficient of complexes, on the other hand, is expressed in the
general form as:
Equation 7 allows the possibility of calculating the diffusion coefficients for all species
present in the solution at equilibrium. To do an evaluation it was necessary to calculate
values of x0in terms of the previously determined formation constants. The diffusion co-
efficient for the Mn(II) aqua ion was calculated from the limiting current of its wave in the
absence of oxalate.
The diffusion coefficients of individual complex species were calculated using the equa-
J Solution Chem (2011) 40: 247–260 255
Table 2 Values of the limiting currents (il), average diffusion coefficient (Dav), and average number of
oxalate ions coordinated to the Mn(II) ion (jav) as a function of the oxalate ion concentration in NaClO4and
Na2SO4solutions at 25°C
Table 3 Values of the diffusion
coefficients of individual
complex species in the
Mn(II)–oxalate system at 25°C
First, the graph of Z versus coxwas plotted and the values of βjDjwere determined numer-
ically using a commercial computer program in conjunction with a nonlinear least-square
regression method. Because the values of βjare already known, it was thus possible to cal-
culate the values of Djfor all species. Their values are collected in Table 3. The contribution
of the diffusion coefficient of each particular complex species (the values of xjDj) to the
average diffusion coefficient Davis shown in Fig. 7.
256 J Solution Chem (2011) 40: 247–260
Fig. 7 The average diffusion coefficient (Dav) as a function of log10coxin NaClO4solutions (1), and the
contributions of individual Mn(II)–oxalate complexes to the average diffusion coefficients: Mn(aq)2+(2),
Mn(ox) (3), Mn(ox)2−
(5)(4), and Mn(ox)4−
As expected, the diffusion coefficient of the species Mn(ox)4−
parison with the diffusion coefficients of other species, and thus its flow to the electrode
surface is almost negligible. The very slow diffusion of the large and negatively charged
concentration of the electroactive form at the electrode surface, resulted in a decrease of
the limiting current with increase of oxalate ions concentration in the solution. On the other
hand, the species Mn(ox) is uncharged and its diffusion coefficient slightly exceeds that of
the Mn(II) aqua ion. Similar phenomena were observed earlier in the Ni(II)–oxalate sys-
tem , but for that system the difference between the diffusion coefficients of those two
forms was more evident.
Summarizing the results obtained above, the electrode reduction of Mn(II)–oxalate com-
plexes in aqueous NaClO4solutions can be expressed by the following scheme:
is extremely low in com-
3complex molecule to the strongly negative charged mercury electrode lowers the
Mn(aq)2+? Mn(ox) ? Mn(ox)2−
E1/2= −1.4975 V
2.303RT/αnαF = 0.033 V
3.2 Reduction of Mn(II) in Na2SO4Solutions
The replacement of perchlorate ions by sulfate ions in the solution caused the half-wave
potential of the Mn(II) reduction wave to shift by 21 mV towards more negative values
and simultaneously diminished its limiting current. This indicated that Mn(II) ions formed
complexes with sulfate ions.
J Solution Chem (2011) 40: 247–260257
Fig. 8 Distribution of Mn(II) among various complex species as a function of log10coxin Na2SO4solutions
Several potentiometric and conductometric literature studies reported the formation of
only one Mn(II)–sulfate complex with stability constant β1values falling within the range
102.28to 102.40[6–8], whereas Jain et al.  reported from polarographic measurements
the formation of three complexes, Mn(SO4), Mn(SO4)2−
constants β = 8.5,9.0 and 93, respectively. The last value seems to be rather unlikely.
In our study the concentrations of sulfate ions in the solution were decreased with in-
creasing oxalate ion concentration, from 0.1667 mol·L−1to zero, because the sulfate ions
were substituted by oxalate ions while keeping the ionic strength constant.
The separate study of complexation of Mn(II) ions with sulfate ions showed that E1/2
of Mn(II) reduction wave changes linearly with log10[SO2−
ion concentrations. This fact implies the formation of a single Mn(SO4) species. Its stability
constant β1was estimated to be 24. Although the obtained value of β1is ten times lower
than reported in the literature it seems to be quite reliable, and correlates well with respective
values obtained polarographically for other metal ions, e.g., for Ni(II) [30, 31], Cd(II) ,
and Zn(II) .
Thus, in sulfate solutions Mn(II) is present as mixtures of the Mn(II) aqua ion and Mn–
SO4complexes. Addition of oxalate ions to sulfate solutions simultaneously results in the
formation of Mn(II)–oxalate complexes and a decrease in the concentrations of MnSO4
Taking the values of the stability constants for MnSO4and of the three Mn(II)–oxalate
complexes obtained as described above, the expected values of E1/2were estimated for the
Mn(II) reduction wave at various concentrations of oxalate and sulfate ions. It appears that
in the oxalate concentration range of 1×10−3to 4×10−2mol·L−1, the experimental values
of E1/2somewhat exceed the calculated ones. This result suggests formation of the ternary
species Mn(SO4)(ox)2−in solution, which dominates in this oxalate concentration region.
Its stability constant was estimated to be 103.08. The distribution of Mn(II) among different
species in Na2SO4solutions is shown in Fig. 8.
3, with stability
4] in the studied range of sulfate
258J Solution Chem (2011) 40: 247–260
The diffusion coefficient of the Mn(SO4)(ox)2−complex species was determined in a
similar manner to that reported above, from the limiting current of the Mn(II) wave recorded
at various concentration of oxalate ions in Na2SO4solutions. Its value is given in Table 3.
The reduction mechanism in this system is similar to that in NaClO4solutions; the free
Mn(II) aqua ion is the only species being reduced on the mercury electrode, whereas in
the solutions five complexes exist in labile equilibrium, namely: Mn(SO4), Mn(SO4)(ox)2−,
lowing scheme, were the numerical values are base 10 logarithms of the corresponding
3. The complex equilibria may be described by the fol-
Since manganese is both a toxic and a biologically essential trace metal, its speciation is of
fundamental importance for better understanding of its role in biological reactions.
A careful analysis of the dependence of the current–potential curves on the concentration
of free ligand enabled the determination of the composition of the species reduced at the
mercury electrode and of those existing in the solutions.
Reported potentiometric and conductometric studies indicated the formation of only one
Mn(II)–oxalate complex, i.e. Mn(ox) (Table 1). In contrast, the polarographic measurements
reported here indicate the formation of two other complexes: Mn(ox)2−
perchlorate solution and the ternary species Mn(SO4)(ox)2−in sulfate solution, and allowed
the determination of their stability constants with relatively high precision.
Mn(II) complexes are characterized by low stability. The values of β obtained in this
work for Mn(II)–oxalate complexes are concordant with those determined for other divalent
metal ions. The radius of Mn(II) lies between those of Mg(II) and Ca(II). The half-filled 3d5
shell of the Mn(II) ion causes it to behave as a spherically polarizable ion, with no crystal-
field stabilization energy; hence regarding the structure of its complexes it will behave in a
manner intermediate between those of Mg(II) and Ca(II). Comparison of the values of β for
oxalate complexes of these three divalent cations (Table 4) led to the conclusion that Mn(II)
binds to oxalate ions with a strength roughly equal to that for Ca(II) and only slightly more
strongly than that for Mg(II). On the other hand, the obtained values of β for Mn(II)–oxalate
complexes also display good concordance with those of 3d transition metal ions complexes
and follow the Irving-Williams sequence: Mg < Mn∼= Ca < Fe < Co < Ni < Cu > Zn.
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J Solution Chem (2011) 40: 247–260259
Table 4 Stability constants of
selected oxalate complexes
4.04 5.16 
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