<ref name=lee2007>Sung Oh Lee, Tam Tran, Byoung Hi Jung, Seong Jun Kim, and Myong Jun Kim (2007): "Dissolution of iron oxide using oxalic acid". ''Hydrometallurgy'', volume 87, issues 3–4. pages 91-99. {{doi|10.1016/j.hydromet.2007.02.005}}</ref>
The dissolution of iron oxide is believed to take place via a
photo-electrochemical reduction process, involving a complicated
mechanism of charge transfer between the predominant oxalate species,
namely ferric oxalate, Fe(C2O4)33−, ferrous oxalate, Fe(C2O4)22−
acting also as an auto-catalyst, and the oxalate ligand on the iron
oxide surface (Taxiarchou et al., 1997b, Blesa et al., 1987). In the
absence of light the reaction proceeds slowly which complicates the
reaction further.
The solution pH governs the distribution of various oxalate ions in
the leach system. Below pH 1.2, oxalic acid exists mainly as H2C2O4,
whereas HC2O4− is the most predominant species (mole fraction > 0.92)
at pH 2.5–3.0. Above pH 4, C2O42− is the predominant species. The
speciation of Fe(III) oxalate and Fe(II) oxalate is also governed by
pH and total oxalate concentration (Panias et al., 1996). For a
solution having pH > 2.5 and an oxalate concentration higher than 0.1
M, the most predominant Fe(III)-oxalate species is Fe(C2O4)33−. At
these conditions (pH > 2.5 and oxalate concentration higher than 0.1
M) the predominant Fe(II) complex species is Fe(C2O4)22−.
The dissolution process also has to be optimized with respect to
oxalate concentration and pH to minimize the precipitation of ferrous
oxalate. On Eh–pH diagrams (Sukhotin and Khentov, 1980) reproduced in
Fig. 1, the predominance of FeC2O4(s) is clearly shown for the system
containing 0.21 M oxalate (right-sided graph). Without oxalate, Fe2O3
and Fe3O4 will be dissolved in acid forming Fe2+, whereas in the
presence of oxalic acid, solid FeC2O4(s) is the predominant species
existing over a wide range of pH from acidic zone to pH > 7 in the
potential range where reductive dissolution of iron oxides takes place
for 0.21 M oxalate. This implies that solid FeC2O4(s) will be finally
formed when the oxalate concentration is 0.21 M (as shown in this
graph). Unfortunately there is no reference to the concentration of
total Fe used for these diagrams, making it difficult for the
interpretation of the process involved. As a result, these diagrams
however could not be used to explain the fact that iron oxide could
finally be dissolved by oxalate. There must be another reaction step
involved which causes the solid ferrous oxalate to re-dissolve if
formed, or there must be conditions which allow the dissolution to
take place, indicating the shortfall of Sukhotin and Khentov's Eh–pH
diagrams.
The iron dissolution process therefore takes place via an
electrochemical process, summarised below:
Oxidation of oxalate to form carbonic acid or carbon dioxide,(1)HC2O4−
= H+ + 2CO2 + 2e−Reduction of hematite forming Fe(II) oxalate,(2)2H+ +
Fe2O3 + 4HC2O4− + 2e− = 2Fe (C2O4)22− + 3H2OThe dissolution reaction
is therefore:(3)H+ + Fe2O3 + 5HC2O4− = 2Fe (C2O4)22− + 3H2O + 2CO2.
The overall reaction indicates that species involved in the leaching
would be hydrogen ions, oxalate and iron oxide particles. At the
optimum pH 2.5–3.0 temperature, concentration of oxalate, iron oxide
mineralogy and its particle size will determine the reaction kinetics.
The charge transfer mechanism could also be assisted by the presence
of Fe(II) as experienced in previous studies.
The solubility of Na HC2O4(s) is 17 g/L and 210 g/L at 25 and 100 °C,
respectively (CRC Handbook of Chemistry and Physics, 1982).
Na2C2O4(s) of which the solubilities are 37 g/L and 63.3 g/L at 25 and
100 °C, respectively.
The solubilities for (NH4)2C2O4(s) of 25.4 g/L and 118 g/L at 25 and
50 °C, respectively (CRC Handbook of Physics and Chemistry, 1982)
It is predicted that the ammonia-bioxalate (NH4)HC2O4(s) solid is very
soluble (no figure on solubility given in the CRC Handbook of Physics
and Chemistry, 1982) and its re-dissolution must form the stable
complex (NH4)C2O4−.
The dissolution of iron oxide seems to be governed by many reaction
steps. The reaction proceeded well in the pH range 2.5–3.0, outside of
which, the reaction rate dropped significantly.
Fig. 12. Speciation of various oxalate species at 0.2 M oxalic acid.
the region in which FeC2O4(s) is stable is from pH1.6 to pH3.2. This
diagram also indicates that HC2O4− species is critical for the
dissolution process (Fig. 13).
Fig. 13. Stability diagram for Fe-Oxalate ([Fe] = 0.01 M, [oxalate] =
0.2 M) — shaded area is for solid species.
Fig. 14. Stability diagrams at 0.2 M for ammonium oxalate (a) and
sodium oxalate (b) — shaded area is for solid species.
<ref name=>R. N. Sahoo, P. K. Naik, S. C. Das (2001): "Leaching of manganese from low-grade manganese ore using oxalic acid as reductant in sulphuric acid solution". ''Hydrometallurgy'', volume 62, issue 3, pages 157-163{{doi|0.1016/S0304-386X(01)00196-7}}</ref>
The dissolution of manganese is due to reduction of its dioxide by
oxalic acid. The reduction (Ehrlich, 1980) between MnO2 and oxalic
acid in acid medium may be given as follows.(1) MnO2 + H2C2O4 + 2 H+ =
Mn++ + 2CO2 + H2O
<ref name=ahmed1953>F. R. Ahmed and D. W. J. Cruickshank (1953): "A
refinement of the crystal structure analyses of oxalic acid
dihydrate". ''Acta Crystallographica'' volume 6, pages 385-392.
{{doi|10.1107/S0365110X53001083}}</ref>
C-C 153 pm C-O1 129 pm C-O2 119 pm
<ref name=sabi1969>T. M. Sabine, G. W. Cox and B. M. Craven (1969): "A neutron diffraction study of [alpha]-oxalic acid dihydrate" ''Acta Crystallographica Section B'', volume B25, pages 2437-2441. {{doi|10.1107/S0567740869005905}}</ref>
Space group ''C''<sup>5</sup><sub>2''h''</sub>--''P''2<sub>1</sub>/''n'',
a = 6.119, b = 3.607, c = 12.057, beta = 106°19', Z = 2.
<ref name=hark1972>S. Harkema, J. W. Bats, A. M. Weyenberg and D. Feil (1972) "The crystal structure of urea oxalic acid (2:1)". ''Acta Crystallographica Section B'', volume B28, pages 1646-1648. {{doi|10.1107/S0567740872004789}}</ref>
Evaporation of a solution of urea and oxalic acid in 2:1 molar ratio
yields a solid crystalline compound (H2C2O4).[CO(NH2)2]2, in which the
neutral molecules are held by hydrogen bonds with the oxygen atoms.
<ref name=colm2019>Francisco Colmenero (2019): "Negative area compressibility in oxalic acid dihydrate". ''Materials Letters'', volume 245, pages 25-28. {{doi|10.1016/j.matlet.2019.02.077}}</ref>
Theoretical studies indicate that oxalic acid dihydrate is one of very
few crystalline substances that exhibit [[negative area
compressibility]]. Namely, when subjected to isotropic tension
[[stress (mechanics)|stress]] (negative [[pressure]]), the ''a'' and
''c'' [[lattice parameter]]s increase as the stress decreases from
−1.17 [[gigapascal|GPa]] to −0.12 GPa and from −1.17 GPa to −0.51 GPa,
respectively.
<ref name=niem1992>J. Nieminen, M. Rasanen, and J. Murto (1992): "Matrix-isolation and ab initio studies of oxalic acid". ''Journal of Physical Chemistry'', volume 96, issue 13, pahes 5303–5308. {{doi|10.1021/j100192a024}}</ref>
In matrix isolation, the most stable form has two hydrogen bonds
bridging the two carboxyls, rather than inside each carboxyl.
<ref name=>M. K. Chantooni Jr. and I. M. Kolthoff (1975): "Acid-base equilibriums in methanol, acetonitrile, and dimethyl sulfoxide in acids and salts of oxalic acid and homologs, fumaric and o-phthalic acids. Transfer activity coefficients of acids and ions". ''Journal of Physical Chemistry'', volume 79, issue 12, pages 1176–1182 {{doi|10.1021/j100579a007}}</ref>
Solubility of dicarboxylic acids in water and methanol at 25 C (brackets is num of
CH2 units in chain):
Succinic[2] 0.70 1.28
Glutaric[3] ? 4.58
Adipic[4] 0.136 1.12
Pimelic[5] 0.324 2.76
Suberic[6] 0.152 0.85
Azelaic[7] 0.013 1.48
Sebacic[8] 0.0015 0.51
Fumaric[?] 0.051 0.51
<ref name=apel1987>Alexander Apelblat and Emanuel Manzurola (1987): "Solubility of oxalic, malonic, succinic, adipic, maleic, malic, citric, and tartaric acids in water from 278.15 to 338.15 K". ''The Journal of Chemical Thermodynamics'', volume 19, issue 3, pages 317-320 {{doi|10.1016/0021-9614(87)90139-X}}</ref>
Mole fractions
K oxal malon succin adip male mal citr tartar
278.15 5 0.00924 0.1686 0.00539 0.00123 0.06729 0.1123 0.09084 0.1211
283.15 10 0.01124 0.1791 O.OO680 0.OO143 0.07786 0.1214 0.09975 0.1273
288.15 15 0.01479 0.1903 0.00859 0.00194 0.08688 0.1318 0.1100 0.1326
293.15 20 0.01863 0.2038 0.01092 0.00239 0.09686 0.1454 0.1195 0.1370
298.15 25 0.02301 0.2176 0.01337 0.00307 0.1109 0.1578 0.1321 0.1435
303.15 30 0.02690 0.2309 0.01591 0.00381 0.1237 0.1681 0.1496 0.1503
308.15 35 0.03421 0.2455 0.01928 0.00476 0.1419 0.1810 0.1617 0.1583
313.15 40 0.04133 0.2633 0.02384 0.00619 0.1545 0.1976 0.1687 0.1638
318.15 45 0.04934 0.2767 0.02961 0.00843 0.1745 0.2094 0.1823 0.1720
323.15 50 0.05887 0.2911 0.03550 0.01080 0.1842 0.2313 0.1904 0.1826
328.15 55 0.06956 0.3120 0.04237 0.01518 0.1990 0.2482 0.1991 0.1924
333.15 60 0.07813 0.3335 0.04862 0.02135 0.2177 0.2691 0.2103 0.1989
338.15 65 0.09830 0.3532 0.06023 0.02437 0.2323 0.2939 0.2239 0.2088
grams per liter:
oxal
5 0.00924 46.92
10 0.01124 57.19
15 0.01479 75.53
20 0.01863 95.51
25 0.02301 118.49
30 0.02690 139.08
35 0.03421 178.21
40 0.04133 216.90
45 0.04934 261.12
50 0.05887 314.71
55 0.06956 376.12
60 0.07813 426.39
65 0.09830 548.47
46.9 g/L (5 °C), 57.2 (10 °C), 75.5 (15 °C), 95.5 (20 °C), 118 (25 °C), 139 (30 °C), 178 (35 °C), 217 (40 °C), 261 (45 °C), 315 (50 °C), 376 (55 °C), 426 (60 °C), 548 (65 °C)
gawk '/[0-9]/ { t = $1; f = $2; x = f/(1-f)*90.636/18.0153*1000; printf "%3d %7.5f %6.2f\n", t, f, x}'
<ref name=menc2004>Batella Menczel, Alexander Apelblat, Eli Korin (2004): "The molar enthalpies of solution and solubilities of ammonium, sodium and potassium oxalates in water". ''The Journal of Chemical Thermodynamics'', volume 36, issue 1, pages 41-44. {{doi|10.1016/j.jct.2003.09.012}}</ref>
T/K m/(mol kg^{-1}) T/C g/L
----- ---------------- --- ----------------
NH4 Na K NH4 Na K
----- ---------------- --- ----------------
273.15 0.201 0.210 1.483 0 24.9 28.1 246.5
278.15 0.000 0.219 1.598 5 0.0 29.3 265.6
283.15 0.259 0.000 1.739 10 32.1 0.0 289.1
288.15 0.000 0.243 1.811 15 0.0 32.6 301.0
293.15 0.361 0.263 1.957 20 44.8 35.2 325.3
298.15 0.422 0.270 2.060 25 52.4 36.2 342.4
303.15 0.496 0.284 2.170 30 61.6 38.1 360.7
308.15 0.577 0.301 2.300 35 71.6 40.3 382.3
313.15 0.653 0.316 2.390 40 81.0 42.3 397.3
318.15 0.777 0.326 2.510 45 96.4 43.7 417.2
323.15 0.880 0.345 2.620 50 109.2 46.2 435.5
328.15 1.004 0.355 2.730 55 124.6 47.6 453.8
333.15 1.164 0.367 2.890 60 144.5 49.2 480.4
gawk '/[0-9]/{ tk=$1; a=$2; b=$3; c=$4; tc=tk-273.15; sa=sol(a,124.1); sb = sol(b,134); sc = sol(c,166.22); printf "%6.2f %5.3f %5.3f %5.3f %3.0f %5.1f %5.1f %5.1f\n", tk, a,b,c, tc,sa,sb,sc;} function sol(m,M){ return m*M }'
molar mass (g/mol):
(NH4)2Ox 124.1
Na2Ox 134
K2Ox 166.22
ICSC: NH42Ox 45 g/L at 20 °C;
Merck: NH42Ox 50 g/L at 25 °C? 384 at 100 C
Wikipedia: Na2Ox 26.9 g/L (0 °C) 37.0 g/L (20 °C) 62.5 g/L (100 °C)
Santa Cruz: K2Ox 360 g/L at 20 °C.
<ref name=hill1935>Arthur E. Hill and Edgar F. Distler (1935): "The Solubility of Ammonium Oxalate in Water".''Journal of the American Chemical Society'', volume 57, issue 11, pages 2203–2204. {{doi|10.1021/ja01314a049}}</ref>
T(C) wt% g/L
0.00 2.269 23.2
10.30 3.107 32.1
16.78 3.892 40.5
25.00 4.985 52.5
34.97 6.630 71.0
44.75 8.619 94.3
60.30 12.300 140.3
74.80 16.440 196.7
87.70 20.860 263.6
99.80 25.790 347.5
gawk '/[0-9]/ { tc=$1; wp=$2; so=wp/(100-wp)*1000; printf "%5.2f %6.3f %6.1f\n", tc, wp, so }'
<ref name=roza2009>Mohd Zul Helmi Rozaini, Peter Brimblecombe (2009): "The solubility measurements of sodium dicarboxylate salts; sodium oxalate, malonate, succinate, glutarate, and adipate in water from T = (279.15 to 358.15) K". ''Journal of Chemical Thermodynamics'', volume 41, issue 9, pages 980-983. {{doi|10.1016/j.jct.2009.03.017}}</ref>
Figures 1-s2.0-S0021961409000688-gr*-*.jpg are ln(m) versus 1/T where m = solubility in mol/kg and T = kelvin temp.
At T = 298 K = 25 C, the solubility of sodium oxalate from our study is 2.11 mol · kg−1 = 261.9 g/L
Makes no sense !!! See complaint by <ref name=butt2016/> below.
<ref name=goma2013>Esam A.Gomaa (2013): "Solvation parameters for sodium oxalate in aq. EtOH at 301.15 K" ''Europan Chemical Bullettin volume 2, issue 5, pages 259-261. {{doi|}}</ref>
Ef = mole fraction of ethanol.
S = molar solubility g mol^{-1} (???)
T = 301.15 K = 28 C
Ef S
0.0000 7.91
0.0330 7.65
0.0715 7.16
0.1166 6.45
0.1703 6.750
0.2355 6.61
0.3159 6.493
0.4181 6.250
0.5591 6.150
0.7349 6.051
1.0000 5.97
<ref name=taft1953>Robert Taft and Frank H. Welch (1951): "Physical Properties of Aqueous Solutions of Sodium Oxalate, Sodium Malonate, and Sodium Succinate, I" ''Transactions of the Kansas Academy of Science'', volume 54, issue 2, pages 233-246. {{jstor|3625790}}</ref>
Solubilities of the acids:
n MP S(g/dL)
Oxalic 2 187 10.2
Malonic 3 135 138
Succinic 4 185 6.8
Glutaric 5 97.5 63.9
Adipic 6 151 1.4
Pimelic 7 105 2.5
Suberic 8 142 0.14
Azelaic 9 106 0.2
Sebacic 10 134 0.1
<ref name=butt2016>Lukas G. Buttke, Justin R. Schueller, Christian S. Pearson, and Keith D. Beyer (2016): "Solubility of the Sodium and Ammonium Salts of Oxalic Acid in Water with Ammonium Sulfate". ''Journal of Physical Chemistry A'', volume 120, issue 32, pages 6424–6433. {{doi|10.1021/acs.jpca.6b05208}}</ref>
NaHC2O4: e were able to easily make saturated solutions of this
compound by mixing NaOH and H2C2O4in a1:1 ratio at an elevated
temperature (but below 323 K) andallowing the mixture to cool. We
alsocompletely dried crystals from a saturated solution for IRanalysis
and found good agreement between our spectrum andthat reported in the
literature for NaHC2O4/H2O
The solubility of Na2C2O4 was also recently investigated by Rozaini
and Brimblecomb;32 however, theirvalues are significantly higher than
those reported in theliterature and our DSC data (see Figure S4 in the
Supporting Information.) Rozaini and Brimblecomb did not address
thisdiscrepancy, and we do not have an explanation for how theycould
obtain such high values for the solubility of sodium oxalate.
<ref name=tell1971>Roland Tellgren and Ivar Olovsson (1971): "Hydrogen Bond Studies. XXXXVI. The Crystal Structures of Normal and Deuterated Sodium Hydrogen Oxalate Monohydrate NaHC2O4·H2O and NaDC2O4·D2O". ''Journal of Chemical Physics'', volume 54, issue 1. {{doi|10.1063/1.1674582}}</ref>
The crystal structures of NaHC2O4·H2O and NaDC2O4·D2O have been
determined from threeâ€dimensional single crystal xâ€ray diffractometer
data obtained at room temperature. Two formula units crystallize in a
triclinic unit cell with the dimensions: a = 6.503, b = 6.673,
c = 5.698 Å, α = 85.04, β = 110.00, γ = 105.02° for the hydrogen
compound and a = 6.501, b = 6.671, c = 5.716 Å, α = 84.91 β = 109.93
γ = 105.00° for the deuterated compund. The space group is P1̄. The
hydrogen oxalate ions are linked end to end in infinite chains by
hydrogen bonds (2.571 Ã…). The chains are cross linked to form layers
by both O–H···O bonds from the water molecules (2.808, 2.826 Å) and by
ionic bonds Na+···O. These layers are in turn held together by Na+···O
bonds. The oxalate group is nonplanar with an angle of twist about the
C–C bond of 12.9°. The only significant difference between the bond
distances in the normal and deuterated compound occurs in the shortest
hydrogen bond, which is 0.022 Ã… longer in the deuterated case.
<ref name=ramki2017>C. Ramki, R. Ezhil Vizhi (2017): "Growth, optical, electrical and mechanical properties of sodium hydrogen oxalate hydrate (NaHC2O4·H2O) single crystal for NLO applications". ''Materials Chemistry and Physics'', volume 197, pages 70-78. {{doi|10.1016/j.matchemphys.2017.04.066}}</ref>
A metalorganic NaHC2O4$H2O single crystal was successfully grown by
slow evaporation method. Thegrown crystal was exposed to single
crystal X-ray diffraction which confirms that the crystal belongs
totriclinic crystal system withP1 space group having unit cell
parameters a = 6.53Ã…,b = 6.71Ã…,c = 5.72Ã… and alpha=75, beta = 85.06,
gamma = 70.31
NaHC2O4$H2O single crystal was prepared by dissolving equi-molar ratio
of sodium hydroxide (Sigma Aldrich), oxalic acid (ARgrade)[15]. Boric
acid (AR grade) was then added in 50 ml of theabove said solution. The
solution was stirred continuously for 15 hto obtain a homogeneous
mixture at room temperature. The satu-rated solution was thenfiltered
with high quality whatmanfilterpaper (Cat No 1001125). After attaining
a clearfiltrate, the solutionwas transferred to a beaker with a
perforated lid and the beakerwas left undisturbed for slow evaporation
at room temperature. Atransparent crystal of NaHC2O4$H2O was harvested
in the period of15 days.
<ref name=hami2002>Susan E Hamilton, Paul J Pielage, Robert G Fassett (2002): "Acute renal failure following sodium oxalate ingestion". ''Emergency Medicine'', volume 11, issue 1, pages 35-37. {{doi|10.1046/j.1442-2026.1999.00317.x}}</ref>
A 47-year-old laboratory worker who had a history of depression,
presented to his general practitioner in March 1997 after taking an
overdose of 150 g sodium oxalate and 24 mg flunitrazepam. ... He
ingested the sodium oxalate as a water-based slurry ... Within 10 min
of ingestion the patient developed severe vomiting, and profuse
diarrhoea and had one episode of haematemesis. He was treated with
oral calcium gluconate and charcoal by his general practitioner, ...
[taken to hospital] approximately four hours after ingestion. ... By
day 14 the patient was in the recovery phase of presumed acute tubular
necrosis, producing large volumes of dilute urine. He remained
polyuric for two weeks. At follow up 2 months later, he had returned
to part-time work. He had an elevated creatinine of 150 µmol/L, but
demonstrated no other adverse effects.
Due to its resemblance to magnesium sulphate (Epsom salts) [sodium oxalate]
is an occasional cause of accidental poisoning. In the [20th century] it
was widely used as a domestic cleaner and it accounted for 16% of deaths by suicide.
Oxalic acid is one product of ethylene glycol
(antifreeze) metabolism and this accounts for some of
its toxicity. Ascorbic acid (vitamin C) is also metabol-
ised to oxalate, and there has been a case report of
oxalate cast nephropathy, acute tubular necrosis and
nephrolithiasis from i.v. vitamin C administration.
There has been a case report from Spain of death following sorrel soup
and reports of illness due to the substitution of rhubarb leaves for
spinach in wartime England.
The average lethal adult dose of sodium oxalate is estimated to be
between 15 and 30 g, and death may occur within minutes of ingestion.
However, as little as 5 g may be fatal. In concentrated doses, oxalic
acid has an immediate corrosive effect on the gastrointestinal mucosa,
resulting in severe haemorrhagic gastro- enteritis and hypovolaemic
shock.
Patients who survive the early phase of poisoning, or consume dilute
forms of [sodium oxalate], may experience symptoms of systemic
intoxication. Oxalate complexes with calcium and calcium oxalate are
deposited in the liver, kidney and blood vessels. Calcium levels fall,
producing tetany, muscle spasms, twitching, cramps, seizures and an
impaired conscious state. ... Metabolic acidosis and hepatic necrosis
and failure may also occur.
For first aid and initial manage- ment, give milk (maximum 15 mL/kg in a
child) and/or calcium gluconate 150 mg/kg orally. Emesis should not be
initiated, and activated charcoal with sorbitol (1–2 g/kg) should be
used for gastrointestinal decon- tamination. ... To prevent
hypocalcaemia, 10–20 mL intravenous calcium gluconate (10%) or CaCl (5%)
should be repeated as needed.
If patients survive the initial 2 days, renal complications supervene
and are the main cause of mortality. The mechanism of kidney damage is
contro- versial. Oxalate is excreted by the kidney, and calcium oxalate
crystals precipitate and deposit in the tubules, capillaries and
glomeruli. This is believed to produce vascular stasis, tubular
obstruction, an interstitial inflammatory reaction and ischaemia,
resulting in acute tubular necrosis. Patients who develop oliguria,
anuria and an acute rise in urea and creatinine require dialysis. The
urinary profile demonstrates haematuria, albuminuria, reduced specific
gravity, calcium oxalate crystals and occasional casts. Recovery from
renal failure may take weeks to months.
<ref name=james1978>Lynn F. James (1978): "Oxalate poisoning in livestock". ''Effects of Poisonous Plants on Livestock'', volume 1978, pages 139-145. {{doi|10.1016/B978-0-12-403250-7.50020-0}}</ref>
It now appears that the toxic effects these various oxalate compounds
[from ingested plants] have on animals may be modified by the cation
with which they are associated. Some of these effects are discussed in
this paper.
<ref name=>Anja Verhulst, Marc E. De Broe (2008): "[https://link.springer.com/chapter/10.1007/978-0-387-84843-3_32 Oxalate]" chapter of ''Clinical Nephrotoxins'', pages 749-756. {{doi|10.1007/978-0-387-84843-3_32}} {{isbn|978-0-387-84843-3}}</ref>
Plant oxalate is the main regulator of calcium concentrations in
plant tissues, an important factor in plants defense (against
herbivores), and in heavy metal tolerance [2
Dietary oxalate is absorbed throughout the length of the intestine,
but mainly in the small intestine. ... When calcium is plentiful in
the gut, a greater proportion of oxalate will be complexed to the
cations leaving less free for absorption. Hence patients with
hyperoxaluria should be advised to consume a calcium rich diet
Calcium oxalate is practically insoluble (8.76x10−8 mol/L at 37°.C in
a urine like solution [1])
Oxalate is an unavoidable component of the human diet since it is a
ubiquitous component of plants
Depending on dietary intake, daily oxalate excretion in healthy
volunteers varies from 0.1 to 0.45 mmol.
<ref name=>S. C. Noonan, G. P. Savage (1999): "Oxalate content of foods and its effect on humans". ''Asia Pacific Journal of Clinical Nutrition'', volume 8, issue 1, pages 64-74. {{doi|10.1046/j.1440-6047.1999.00038.x}}</ref>
The mean daily intake of oxalate in English diets has been calculated
to be 70–150 mg, with tea appearing to contribute the greatest
proportion of oxalate in these diets; rhubarb, spinach and beet are
other common high oxalate-content foods. Vegetarians who consume
greater amounts of vegetables will have a higher intake of oxalates,
which may reduce calcium availability.
Oxalic acid forms water-soluble salts with Na, K, and NH4ions; it also
binds with Ca, Fe2+, and Mg2+, rendering these minerals unavailable to
animals. However, Zn2+ appears to be relatively unaffected.
In plants with a cell sap of approxi- mately pH 2, such as some
species of Oxalis and Rumex, oxalate exists as the acid oxalate
(HC2O4–), primarily as acid potassium oxalate.
In plants with a cell sap of approximately pH 6, such as some plants
of the goosefoot family, it exists as the oxalate ion (C2O42–),
usually as soluble sodium oxalate and the insoluble calcium and
magnesium oxalates. Calcium oxalate (Ca(COO)2) is insoluble at a
neutral or alkaline pH, but freely dissolves in acid.
The oxalic acid content is variable within some species; some
cultivars of spinach (Universal, Winter Giant) contain 400 to 600
mg/100 g, while others range from 700 to 900 mg/100 g.
Aspecies of snail (Limicolaria aurora), used as human food in Nigeria, has been
reported to contain 381 mg total oxalate/100 g dry weight (DW).
The mollusc, dogwhelk (Thais cattifera), contains an even higher level of oxalate, 1686 mg/100 g DW.
Fungi such as Aspergillus niger, Penicillium, Mucor,
Boletus sulphurens and Sclerotinia, can synthesize oxalic
acid at a rate of up to 4–5 g/100 g DW in isolated cultivation,
in foodstuffs and on the surface of forages.
The coriander leaf (Coriandrum sativum) contains 1268 mg/100 g,
while horsegram and santhi (Boer- navia diffusa) contain 508 mg/100 g
and 3800 mg/100g, respectively. Sesame seeds have been reported
to contain relatively high quantities of oxalate, ranging from 350 to
1750 mg/100 g FW
The proportion of oxalic acid in the leaves of the goosefoot family
can double during ripening and occasionally accumulate to such a
degree that it makes up more than 15% of the total DW, which
suggests that oxalates are an end product of metabolism and act as a
‘dump’ system.
The absorption of oxalates
from individual foods varies depending on their dietary con-
ditions and source; in general the absorption is relatively lim-
ited. It has been estimated that 2–5% of administered oxalate
is absorbed in humans.
Using radiolabelled oxalate, 6.6%
of the administered dose was absorbed when consumed with
a normal diet, whereas 12% was absorbed when oxalate was
consumed during fasting.
The percentage of oxalate
absorption varied markedly, from 1% for rhubarb and
spinach to 22% for tea, but generally absorption was higher
at low doses.
Cocoa contains theobromine (1500–2500 mg/100 g) and tannic acid
(4000–6000 mg/100 g), both of which are more toxic than the oxalic
acid present (500–700 mg/100 g).
The adverse effect of oxalates [for calcium abdosption] is greater if
the oxalate:calcium ratio exceeds 9:4. The adverse effects of oxalates
must be considered in terms of the oxalate:calcium ratio in a food.
... Foods that have a ratio greater than two and that contain no
utilizable calcium have excess oxalate which can bind calcium in other
foods eaten at the same time. Rhubarb, sorrel, beet and spinach are
not good sources of calcium despite their apparently high levels.
A number of plants contain calcium oxalate crystals. These
are not absorbed into the blood stream and remain largely
undissolved within the digestive tract. Thus, they have no
systemic toxicity but the sharp raphide crystals can penetrate
the tissues of the mouth and tongue, causing considerable
discomfort.
Intakes of oxalate exceeding 180 mg/day lead to a marked increase in
the amount excreted.
2–3 cups/day of black tea would not affect the risk of urinary stone
forma- tion. It appears that tea is a significant source of oxalate
intake in English diets
As the majority of oxalate excreted in the urine is
reported to be synthesized endogenously from ascorbate, gly-
colate, glyoxylate and glycine, excessive intake of these sub-
stances would not be advised. Excessive ascorbic acid intake
may increase urinary levels of oxalate, making it a possible
risk factor for kidney stone formation. Ascorbic acid doses
greater than 500 mg/day were reported to induce a significant
increase in urinary oxalate, and doses of 1000 mg/day would
increase urinary oxalate excretion by 6–13 mg/day.
<ref name=chai2011>Nopsiri Chaiyo, Rangson Muanghlua, Surasak Niemcharoen, Banjong Boonchom Panpailin Seeharaj, and Naratip Vittayakorn (2011): "Non-isothermal kinetics of the thermal decomposition of sodium oxalate Na2C2O4". ''Journal of Thermal Analysis and Calorimetry'', volume 107, issue 3, pages 1023–1029. {{doi|10.1007/s10973-011-1675-6}}</ref>
Decomposition to Na2CO3 + CO occurs at a very narrow range centered at
844 K (571 C).
<ref name=yoshi1978>Takayoshi Yoshimori, Yoshihiro Asano, Yasuo Toriumi. Takashi Shiota (1978): "Investigation on the drying and decomposition of sodium oxalate". ''Talanta'', volume 25, issue 10, October , Pages 603-605. {{doi|10.1016/0039-9140(78)80158-1}}</ref>
Sodium oxalate heated for 2hr above 200° and cooled contains less than
20 ppm of water, and may be used as a standard for titrimetry.
The decomposition of Na2C2O4 begins at 290° and heating between 200°
and 250° is recommended for the dehydration of sodium oxalate. The
decomposition is complete between 750° and 800° within 20 min and the
sodium carbonate obtained begins to decompose at above 810°.
Since 1898, sodium oxalate has been used as a standard reference
material (SRM) for standardization of potassium permanganate and
acids. Although the reaction is rather complex and is sometimes not
recommended, the reagent is still used in many industrial
laboratories
No carbon monoxide could be detected on heating at 280†for 3 hr, but
some decomposition at 290†was observed and this increased markedly at
higher temperatures
<ref name=>G. A. Jeffrey and G. S. Parry (1954): "The Crystal Structure of Sodium Oxalate". ''Journal of the American Chemical Society''. volume 76, issue 21, pages 5283–5286. {{doi|10.1021/ja01650a007}}</ref>
Anhydrous Na2C2O4
Thedifferenceinconfigurationbetweenthenon-planaroxalateionfound
in(NH4)2(C00)2'H203andtheplanaracidmolecule
ina-(COOH)2,4/3-(COOH)2,6(C00H)2-2H20,sraisesthequestion
astowhetherthisisacharacteristicdifferencebetweentheionand
theacidmoleculesoraconsequenceofinterionicorintermolecular
forcesinthecrystalstructures.
SaltgroupNo.ofmolesincellOxalateCelldimensionsbcßDensityObs.Caled.
K2C2O4.H2O C2/c 4 T 9.32 6.17 10.65 110°58' 2.154 2.139
Rb2C204.H20 C2/c 4 T 9.66 6.38 11.20 110°30' 2.763 2.845
Ag2C204 P2_1/a 2 T 9.47 6.16 3.46 104 5.029 5.151
Na2C204 P2_1/a 2 T 10.35 5.26 3.46 92°54' 2.34 2.365
(T = $\bar 1$)
Thecrystalstructureofsilveroxalatehasbeenstudiedinmoredetail.7
Itisnottrulyanionicstructurefortheprincipalfeatureistheexistenceof
chainmoleculesofcomposition(Ag2C204)zwithstrongbondsbetweensilver
andoxygenatomsalongthelengthofthechain,
WithtwoNa2(COO)2molecules
intheunitcell,theNaionslieingeneralpositionsandthecenters
ofsymmetryoftheoxalateionsareat(0,0,0)and(1/2,1/2,0).
Suitablecrystalsforintensitymeasurementswereobtainedby
slowcoolingofaqueoussolutions
Oxalate ion is planar.
Atomiccoordinatesin Atomiccoordinatesin angstroms
fractionsofmonoclinicaxes (Z'is perpendicular to to X and Y)
x y z X Y Z'
Na 0.353 0.053 0.307 3.600 0.279 1.061
O1 0.152 -0.114 0.163 1.544 -0.600 0.563
O2 0.067 0.260 0.228 0.653 1.368 0.788
C 0.064 0.040 0.107 0.643 0.210 0.370
Interatomic distances of oxalate
C-C 1.54
C-O1 1.23
C-O2 1.23
O1-O2 2.18
O1-O2' 2.69
Angles in oxalate
C-C-O1 120.6
C-C-O2 115.3
O1-C-02 124.3
Na+ is coordnated with 6 O
Interionic distances Na-O vary between 2.29 and
2.64 angstroms
Columns of like ions along the c axis,
with each column surrounded by six of opposite
charge (?!)
<ref name=daha1997>N. D. Dahale, K. L. Chawla, N. C. Jayadevan, V. Venugopal (1997): "X-ray, thermal and infrared spectroscopic studies on lithium and sodium oxalate hydrates".''Thermochimica Acta'', volume 293, issues 1–2, pages 163-166. {{doi|10.1016/S0040-6031(97)00015-4}}</ref>
The title should say "lithium URANYL and sodium URANYL oxalate hydrates".
<ref name=buch2003>Richard Buchner, Faradj Samani, Peter M. May, Peter Sturm, Glenn Hefter (2003): "Hydration and Ion Pairing in Aqueous Sodium Oxalate Solutions". ''ChemPhysChem'', volume 2003, issue 4, pages 373-378. {{doi|10.1002/cphc.200390064}}</ref>
No useful info?
<ref name=heft2018>Glenn Hefter, Andrew Tromans, Peter M. May, and Erich Königsberger (2018): "Solubility of Sodium Oxalate in Concentrated Electrolyte Solutions". ''Journal of Chemical Engineering Data'', volume 63, issue 3, pages 542–552. {{doi|10.1021/acs.jced.7b00690}}</ref>
The solubility of solid sodium oxalate (Na2Ox) has been measuredin a
variety of concentrated aqueous electrolyte solutions atT= 298.15,
323.15,and 343.15 K by titration of dissolved oxalate with
permanganate. The electrolytesolutions studied (not necessarily at all
temperatures) were NaCl, NaClO4,NaOH, LiCl, KCl, Me4NCl, and KOH at
concentrations ranging fromapproximately 0.5 mol·kg−1to at least 5
mol·kg−1
Thesolubility of Na2Ox(s) decreased markedly with increasing
concentrations ofNa+(aq), due to the common ion effect. This decrease
was almost independent ofthe electrolyte anion. A number of ternary
mixtures of these electrolytes were alsoinvestigated at constant ionic
strength. Consistent with the binary mixtures, thesolubility of
Na2Ox(s) showed almost no dependence on solution composition
atconstant Na+(aq) concentrations.
Solubilities in non-Na+media, with theexception of Me4NCl, showed
small but regular increases with increasingconcentration of added
electrolyte, probably reflecting activity coefficient variations. The
solubility data in certain Na+-containingmedia could be correlated
accurately at all temperatures and concentrations using a relatively
simple Pitzer model withinteraction parameters for Na2Ox(aq) assumed
to be identical to those available in the literature for Na2SO4(aq).
Solubility of Na2Ox in H2O and p = 100 kPa,
in mol per kg of H2O. Solid phse is alpha.:
T = 298.15 K : 0.263--0.272 (avg 0.269, conf. lit.)
T = 323.15 K : 0.317--0.334
T = 343.15 K : 0.383--0.385
snapshot 17:48
<ref name=balc1980>W. Balcerowiak, J. Wasilewski, and Cz. Latocha (1980): "Thermoanalytical investigation of mixtures containing oxalic acid, sodium hydrogen oxalate and sodium oxalate". ''Journal of Thermal Analysis and Calorimetry'', volume 18, issue 1, pages 57–63. {{doi|10.1007/bf01909453}}</ref>
Considers thermal decomposition of
mixtures of H2C2O4.2H2O, NaHC2O4.H2O, Na2C2O4. Reactions
for pure components:
Dehydration:
H2C2O4.2H2O --> H2C2O4 + 2 H2O
NaHC2O4.H2O --> NaHC2O4 + H2O
Disproportionation of hydrogen oxalate:
2 NaHC2O4 --> Na2C2O4 + H2C2O4
Decomposition od anhydrous oxalate:
Na2C2O4 --> Na2CO3 + CO
Sodium hydrogen oxalate monohydrate was prepared by evaporating the
excess water from an aqueous solution of a stoichiometric mixture of
oxalic acid and sodium oxalate and drying the precipitated salt at
about 350 K till free from moisture.
Melting point of H2C2O4.2H2O = 370 K (96.85 C)
Melting point of H2C2O4 = 460 K (186.85 C)
The establishment by X-ray analysis of a new,
previously-unidentified crystalline phase apart from Na2C204 in the
mixture (f) heated up to 460 K [indicates that there are reactions
between the components]
DTG curve of H2C2O4.2H2O shows (1) gradual dehydration in
300--365 K, then (4) gradual sublimation in 406-460.
DTG curve of NaHC2O4.H2O shows (2) rather sharp dehidration
in 365--385, (6) disproportionation in 460-485 with sublimation of
H2C2O4, then (7) decomp to Na2CO3 + C in 660--770.
Data for mixtures suggest two more solid phases,
3NaHC2O4.H2C2O4.H2O and 3NaHC2O4.H2C2O4
Dehydration of 3NaHC2O4.H2C2O4.H2O (3) occurs in 385--405.
Dissociation of 3NaHC2O4.H2C2O4 (5) with sublimation of H2C2O4
occurs in 485--540
<ref name=pede1967>Berit F. Pedersen (), "Interpretation of the infrared spectra of solid alkali metal oxlates, their hydrates and perhydrates". ''Acta Chemica Scandinavica'', volume 21, issue 3, pages 801--811. {{doi|}}</ref>
Monoperhydrates:
Li2C2O4.H2O2
Na2C2O4.H2O2
K2C2O4.H2O2
RbC2O4.H2O2
CsC2O4.H2O2
Also hydrates:
K2C2O4.H2O
RbC2O4.H2O
CsC2O4.H2O
and fully deuterated versions of the K, Rb, Cs hydrates and perhydrated]s.
Main absorption peaks, cm-1
NaC2O4 ~3400 (weak), ~1630 (strongest), ~1320 (medium), ~770 (md)
K2C2O4.H2O ~3400 (medium), ~1600 (str) ~1305 (str) ~770 (md)
Li2C2O4.H2O2 ~3200 (md) ~2650 (md), ~1640 (str) ~1320 (str) ~810 ms) ~770 (md)
NaC2O4.H2O2 ~3200 (md) ~2650 (md), ~1640 (str), ~1410 (md) ~1315 (str) ~820 (md) ~770 (md)
K2C2O4.H2O2 ~3100 (md) ~2650 (md), 1595 (str) ~1305 (str) 920 (wk), 880 (wh), 735 (str)
RbC2O4.H2O2 3100 (md) 2700 (md) 1595 (str), 1300 (str), 730 (md)
CsC2O4.H2O2 3050 (md) 2700 (str) 1560-1610 (vstr) 1300-1310 (vstr) 940 (md) 870 (md) 745 (vstr)
1658 and 1587: antisymmetric O-C-O stretching freq.
1340 and 1305: symmetric O-C-O stretching freq.
775 and 740: in-plane O-C-O deformation freq.
1400: oxalate Raman band?
3350-3400: O-H stretching of water.
1207-1225: H-O-H bending of water.
710-712: torsional water?
880: O-O stretching of peroxide
Have planar centersymmstric oxalate:
LiC2O4
NaC2O4 Na2C2O4.H2O2
K2C2O4.H2O K2C2O4.H2O2
RbC2O4.H2O RbC2O4.H2O2
In Li2C2O4.H2O2 the H2O2 molecule seems to be planar
with 180 degre H-O-O-H dihedral. In Cs2C2O4.H2O2
however it is like in solid H2O2.
<ref name=Bold2009>Elena V. Boldyreva, Hans Ahsbahs, Vladimir V. Chernyshev, Svetlana N. Ivashevskaya and Artem R. Oganov (2006): "Effect of hydrostatic pressure on the crystal structure of sodium oxalate: X-ray diffraction study and ab initio simulations". ''Zeitschrift für Kristallographie'' (''Crystalline Materials''), volume 221, issue 3, pages 186-197. {{doi|10.1524/zkri.2006.221.3.186}}</ref>
Effect of hydrostatic pressures up to 8 GPa on the crystals of Na2C2O4
(sp. gr. P21/c) was studied in situ in the diamond anvil cells a) in
neon, b) in methanol-ethanol mixture by high-resolution X-ray powder
diffraction (synchrotron radiation, λ = 0.7 Å, MAR345-detector). Below
3.3–3.8 GPa, anisotropic structural distortion was observed, which was
similar to, but not identical with that on cooling. At 3.8 GPa, a
reversible isosymmetric first-order phase transition without
hysteresis occurred. The orientation of the oxalate anions changed at
the transition point by a jump, and so did the coordination of the
sodium cations by oxygen atoms.
Effect of hydrostatic pressures up to 8 GPa on the crystals of Na2C2O4
(sp. gr. P2_1/c) was studied in situ in the diamond anvil cells a) in
neon, b) in methanol-ethanol mixture by high-resolution X-ray powder
diffraction (synchrotron radiation, λ = 0.7 Å, MAR345-detector). Below
3.3–3.8 GPa, anisotropic structural distortion was observed, which was
similar to, but not identical with that on cooling. At 3.8 GPa, a
reversible isosymmetric first-order phase transition without
hysteresis occurred. The orientation of the oxalate anions changed at
the transition point by a jump, and so did the coordination of the
sodium cations by oxygen atoms.
<ref name=guru1976>A. K. Guru, A. V. Mahajan (1976): "Polarographic studies of indium (III) complex with sodium oxalate". ''Current Science'', volume 45, issue 13, pages 492-494. {{doi|}} {{jstor|24079826}}</ref>
At pH < 4 the formation of InC204+ has been shown by Dubovenko
of 1:3 indium (III) oxalate complex with dissociation constant
1.905 x 10-13
<ref name=poulou1989>R. Polou, C. Triche, B. Barbier and G. Pitet (1989): "Sodium Copper Oxalate Dihydrate: Na2Cu (C2O4)2. 2H2O Synthesis, Characterization, Morphology and Optical Properties". ''Powder Diffraction'', volume 4, issue 1, pages 14-18. {{doi|10.1017/S0885715600016237}}</ref>
Crystals of sodium copper oxalate dihydrate [Na2Cu (C2O4)2.2H2,O] were
obtained by the gel method, from solutions of oxalic acid and copper
chloride. The crystals form blue needles with idiomorphic faces of
brilliant luster, permitting goniometric measurements and the
determination of the morphology with the aid of crystallographic
parameters. Optically the crystals are biaxial negative, 2V = 38°,
with a weak dispersion r<v. The orientation of the indicatrix was
determined using a universal stage.
Crystals of sodium copper oxalate dihydrate, Na2Cu (C2O4)2.2H2O, were
apparently first obtained in 1929 by Riley. Gleizes et al. (1980)
undertook a preliminary crystallographic study of crystals obtained by
a different technique from Riley's. In the first case, a solution of
33.5 g/L of sodium oxalate was heated and then poured gradually into a
nearly saturated solution of copper sulfate until slight turbidity
appeared. The turbidity was eliminated and the solution clarified by
the addition of a little more sodium oxalate solution. In this way
Riley obtained a dark blue solution which after filtration yielded
extremely fine sky-blue needle-like crystals, rarely more than 8mm
long.
In the method of Gleizes et al. (1980), copper oxalate was dissolved
in an aqueous solution of sodium oxalate; those authors observed
complete dissolution when the molar ratio of sodium oxalate to copper
oxalate was near 2. By evaporating the solution, they obtained long,
prismatic crystals whose crystallographic constants they determined.
<ref name=ceti2011>Halil Cetişli, Gülbanu Koyundereli Çılgı, and Ramazan Donat (2012) "Thermal and kinetic analysis of uranium salts Part 1. Uranium (VI) oxalate hydrates". ''Journal of Thermal Analysis and Calorimetry'', volume 108, issue 3, pages 1213–1222 . {{doi|10.1007/s10973-011-1826-9}}</ref>
In this study, thermal decomposition of uranyl oxalate UO2C2O4·3H2O compound has been examined in detail from a fresh point of view, and reaction kinetic and thermodynamic parameters were determined.
It is found that uranyl oxalate hydrate's decomposition consists of
four major steps, first two are endothermic and second two are
exothermic. Endothermic reactions are the same in all atmospheres and
correspond to dehydration reactions. Two exothermic reactions occur
consecutively and determining the temperature range of reactions is
very stressful, especially in nitrogen atmosphere.
Loses 2 water molecules at 323--388 K (peak 350.5).
Loses the last one at 388--443 (peak 430)
Anhydrous is stable at 443–590 K
Then decomps to UO2CO3 + CO exhotermically, peak at 602-625.
The UO2CO3 decomposes to UO3 + CO2 at 621-640 (peak 633).
<ref name=john2012>M. Jose John, K. Muraleedharan, V. M. Abdul Mujeeb, M. P. Kannan, and T. Ganga Devi (2012): "Effect of pre-compression on the kinetics of thermal decomposition of pure and doped sodium oxalate under isothermal conditions". ''Reaction Kinetics, Mechanisms and Catalysis'', volume 106, issue 2, pages 355–367. {{doi|10.1007/s11144-012-0436-2}}</ref>
Pure and doped samples of sodium oxalate (Na2C2O4) were subjected to
pre-compression and their thermal decomposition kinetics was studied
at five different temperatures in the range 783–803 K under isothermal
conditions by thermogravimetry. The pre-compressed samples decomposed
in two stages governed by different rate laws; the Prout–Tompkins
model best describes the acceleratory stage of the decomposition while
the decay region is best fitted with the contracting cylinder model as
in the case of uncompressed sodium oxalate samples. The rate constants
k1 and k2 of the acceleratory and deceleratory stages of the thermal
decomposition were dramatically decreased on pre-compression. However,
the activation energies, evaluated by model fitting kinetic method, E1
and E2 for the respective stages of decomposition remained unaltered
by pre-compression. The results favor ionic diffusion mechanism
proposed earlier on the basis of doping studies.
<ref name=john2013>M. Jose John, K. Muraleedharan, M. P. Kannan, T. Ganga Devi (2013): "Kinetic studies on the thermal decomposition of phosphate-doped sodium oxalate". ''Journal of Thermal Analysis and Calorimetry'', volume 111, pages 137–144. {{doi|10.1007/s10973-012-2219-4}}</ref>
The decomposition of oxalates involves the cleavage of the C–C bond,
since the products are CO and CO2 that contain only one carbon atom
each. Gorski and Kranicka [23] proposed that the decomposition in
oxalates begins with the heterolytic dissociation of C–C bond forming
CO2 and CO2 2−. In many cases, the C–C bond cleavage is the rate
determining step [23]. The cleavage may be heterolytic to produce CO2
and CO2 2− [23] or homolytic to produce two CO2 − anions [24]. In
silver oxalate [25], the transfer of an electron from the C2O4 2− to
the cation is the first stage of the decomposition which leads to the
rupture of the C–C bond [11]. A review on the literature of the
thermal behaviour of inorganic oxalates reveals that except yttrium
oxalate, all undergo thermal decomposition before melting and the
decomposition kinetics are not complicated except in the case of a
few.
10 g of sodium oxalate was dissolved in 230 mL of distilled water at
boiling temperature in a 500 mL beaker; 10 mL of a solution containing
the desired quantity of PO4 3− is added to the hot solution so as to
achieve a total volume of 240 mL. The solution, containing desired
concentration of the dopant, was then cooled slowly to room
temperature. The beaker containing the solution was covered using a
clean uniformly perforated paper and kept in an air oven at a
temperature of 333 K over a period of 6–7 days to allow slow
crystallization by evaporation. The resulting crystals were removed,
air dried, powdered in an agate mortar, fixed the particle size in the
range 106–125 μm and kept in a vacuum desiccator. The doped samples
were prepared at five different concentrations, viz. 10−5, 10−4, 10−3,
10−2, 10−1 and 1 mol%.
Thermal decomposition of sodium oxalate was found to be very slow
below 783 K and very fast above 803 K. The decomposition was thus
studied in the range 783–803 K.
<ref name=laks1957>B R Lakshmanan (1957): Infrared absorption spectra of copper potassium oxalate, copper sodium oxalate and copper ammonium oxalate"". ''Journal of the Indian Institute of Science'', volume 39, issue 1, pages 30-33.</ref>
(No details on substances other than the spectra).
<ref name=pede1964>Berit F. Pedersen and Bjørn Pedersen (1964): "The crytal structure of Sodium Oxalate Perhydrate, {{chem|Na|2|C|2|O|4}}·{{cehm|H|2|O|2}}". ''Acta Chemica Scandinavica'', volume 18, pages 1454-1468.</ref>
The compound crystallizes in monoclinic system, space group ''P''2<sub>1</sub>/''c'', with two Na2C2O4.H2O2 groups in a unit cell of dimensions
a = 3.548, b = 8.23, c = 9.01 angstroms and beta = 96.0.
As is well known, both lithium an sodiu oxalates crystallize from
a water solutin with no water of crystallization, contrary to the
remaining alkali oxalates which crystallize as monohydrates.
Both crystallize as monoperhydrates.
Sodium oxalate crystallzes from perhydrol [H2O2 30%] as the monoperhydrate
in the form of thin needles. The crystals are unstable and decompose
over a period of a few days.
<ref name=zakn1995>A. Zaknich; Y. Attikiouzel (1995): "Detection of sodium oxalate needles in optical images using neural network classifiers". ''Proceedings of ICNN'95: International Conference on Neural Networks'' (Perth, Australia), volume 4, pages 1699-1702. {{doi|10.1109/ICNN.1995.488875}} {{isbn|0-7803-2768-3}}. {{inspec|5257000}}</ref>
Alcoa of Australia Limited uses the Bayer process to refine alumina.
This involves the digestion of gibbsite from bauxite in a caustic soda
solution and subsequent reprecipitation of the gbbeite in a continuous
precipitation circuit. The gibbsite is calcined at 1000 'C and shipped
to smelters. Some of the organics that are introduced into the Bayer
liquor via the bauxite form sodium oxalate. The liquor in the
precipitation circuit is saturated with respect to sodium oxalate. In
this saturated state there is a likelihood that oxalate will nucleate
as needle-like crystals and co-precipitate with the gbbsite (hydrate)
particles which is not desirable
A description is given of a PC based system for the automatic
detection, counting and sizing of sodium oxalate needles in optical
microscope images predominated by a background of hydrate particles.
<ref name=>K. Mahendra, N. K. Udayashankar (2020): "Growth and comparative studies on oxalic acid dihydrate, potassium oxalate hydrate and potassium hydrogen oxalate oxalic acid dihydrate single crystals". ''Journal of Physics and Chemistry of Solids'', volume 138, article 109263. {{doi|10.1016/j.jpcs.2019.109263}}</ref>
Single crystals of Oxalic acid dihydrate (OAD) H2C2O4.2H2O, Potassium oxalate
hydrate (KOH) K2C2O4.H2O, Potassium hydrogen oxalate oxalic acid dihydrate
(KHOOD) KHC2O4.H2C2O4.2H2O were grown using solvent evaporation technique.
Potassium oxalate consists of 2 potassium atoms combined with oxalate
molecule to form a stable structure. KHOOD consists of one potassium,
one oxalic acid and one oxalate ion combined with 2 water molecules in
hydrated state.
Slow evaporation technique is employed for the synthesis of oxalic
acid dihydrate (OAD), potassium oxalate (KOH) and potassium hydrogen
oxalate oxalic acid dihydrate (KHOOD). Analytical grade OAD and KOH
are purchased and recrystallized to obtain the required crys-tals. The
KHOOD crystals were synthesized using the equimolar mixture of
potassium dihydrogen phosphate (KDP) and oxalic acid dihydrate was
dissolved in DI water and crystallized using evaporation technique
[34, 35]. The crystals were harvested in the time intervals of 15–20
days and compared (see Fig. 1).
It was noticed that the OAD and KOH crystallizes in monoclinic crystal
system with P2_1/a and A2/a [C2/c?] space group. Whereas, KHOOD crystallizes in
triclinic crystals system of P\bar 1 space group.
The absorption peaks for the OAD, KOH and KHOOD crystals were measured
to be 286 nm, 269 nm and 280 nm. There is only one absorption peak for
the crystal in the UV region, which is due to n → π* electron transfer
Unit cell parameters
OAD dens:1.63 a:11.900 b:3.6040 c:6.1240 alf:90 bet:103.3 gam: 90 V:255.60
KOH dens:2.127 a:10.6840 b:6.1820 c:9.226 alf:90 bet:110.77 gam:90 V:569.76
KHOOD dens:1.856 a:6.3632 b:7.0291 c:10.5984 alf:93.835 bet:101.358 gam:100.174 V:454.90
The pure and doped samples were excited at a wavelength of 280 nm and
Photoluminescence emission spectra were recorded in the range 320
nm–550 nm as shown in Fig. 5. The oxalic acid crystal does not show
any emission but KOH and KHOOD crystals emits in UV and blue region.
Single emission peak observed for the KOH and KHOOD crystals. The
emission peak of 399 nm [350-450] and 430 nm [380-480] is observed for
KOH and KHOOD crystals respectively. The emission intensity is more
for KOH crystal when compared with the KHOOD. And also, KOH crystal
emits in broad region when compared with KHOOD crystal. It is worth
noting that, KOH and KHOOD crystal emission are observed in the UV and
blue region of the spectrum and can be due to delocalization effects
of O-C––O bond present in oxalic acid [50].
Oxalic acid: DTA peaks 93 C and 125 C due to dehidration.
130-200 decomposition; completely decomposes aftr 200 C.
Peaks at 191.
Potassium Oxalate monohydrate: dehydrates peak at 116.
Peak at 394 due to unexplained phase transition in anhydr
oxalate? Decomp at 582 (exo) into K2CO3 and CO. (Does the CO burn?)
KHOOD: Dehydration peaks at 85 and 152. Peak at 194 may be decomp of
oxalic acid leaving anhydr KHC2O4. Peak at 279 conjectured to be
liberation of formic acid and CO2. leaving anhydr K2C2O4. Same phase
trans at 393 and decomp at 578. Conjectured reactions:
2[KHC2O4.H2C2O4.2H2O] --> 2[KHC2O4.H2C2O4] + 2H2O
2[KHC2O4.H2C2O4] --> 2 KHC2O4 + 2 CO2 + 2 H2CO2
2 KHC2O4 --> K2C2O4 + CO2 + H2CO2
K2C2O4 --> K2CO3 + CO
The resistivity of the crystals were obtained as 1.399 10^4, 2.988
10^4 and 1.8314 10^4 Ωm
<ref name=>V. Metler (1934): "The System Zinc Oxalate, Potassium Oxalate, Water. II. At 35°". ''Journal of the American Chemical Society'', volume 56, issue 7, pages 1509–1510. {{doi|10.1021/ja01322a018}}</ref>
It thas been shown 1 that the complex compound K2Zn(C204)2'7H20 exists in stable
equilibrium with solutions of potassium oxalate saturated with zinc oxalate at 25°
The hydrate was observed to be unstable, becoming anhydrous when exposed to
the air. Scholder and Linstrom 2 pre-pared the anhydrous compound.
Other complex compounds of the type K2M2(C2O4)3 have been prepared
by Scholder and co-workers. 3
Well-formed crystals of the anhydrous salt K2Zn(C204)2 form from the
hepta hydrate, or from the compound K2Zn2(C204)3 described below, when
they are shaken at 35° with a saturated solution of potassium and
zinc oxalates. This preparation is preferable to that formed
when the heptahydrate is dehydrated in air since the latter tends
to form colloidal solutions.
Stable Systems. — Solutions of various concentrations of
potassium oxalate were saturated with zinc oxalate
(solutions 12-29), or the hydrated complex salt, K2Zn-(C204)2-7H20 (solutions 1-11).
It may be seen that zinc oxalate is the stable solid phase in
equilibrium with solutions up to 25.28% K2C204 (and saturated with
zinc oxalate). From this concentrati onto saturation with both
potassium oxalate and the oxalato-zincate the solutions are in
equilibrium with anhydrous K2Zn(C204)2.
In the preparation of the hepta-hydrate, thin, hexagonal crystals of
another complex com-pound form. Its composition when dried in air is
K2Zn2-(C204)3 5H20. Solutions of potassium oxalate were saturated with
this compound at 35°. Several solutions with their corre-sponding
solid phases are indicated by squares in Fig.1. All tie-lines cross
the K2Zn2(C204)$-H20 axis. Since the tie-lines lie very nearly
parallel to this axis, the water con-tent of the metastable complex is
indefinite. By algebraic extrapolation the formula of the complex salt
appears to be K2Zn2(C204)3·12H20.