Sodium oxalate

From Wikipedia, the free encyclopedia
  (Redirected from Disodium oxalate)
Jump to: navigation, search
Sodium oxalate
Disodium oxalate
Names
IUPAC name
Sodium ethanedioate
Other names
Oxalic acid, disodium salt
Sodium ethanedioate
Identifiers
3D model (Jmol)
ChEBI
ChemSpider
ECHA InfoCard 100.000.501
EC Number 200-550-3
RTECS number K11750000
Properties
Na2C2O4
Molar mass 133.999 g mol−1
Density 2.34 g cm−3
Melting point 260 °C (500 °F; 533 K) decomposes above 290 °C[2]
2.69 g/100 mL (0 °C)
3.7 g/100 mL (20 °C)
6.25 g/100 mL (100 °C)
Solubility soluble in formic acid
insoluble in alcohol, ether
Structure
monoclinic
Thermochemistry
-1318 kJ/mol
Hazards
Safety data sheet Oxford MSDS
Harmful Xn
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Lethal dose or concentration (LD, LC):
11160 mg/kg (oral, rat)[3]
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YesYN ?)
Infobox references

Sodium oxalate, or disodium oxalate, is the sodium salt of oxalic acid with the formula Na2C2O4. It is a white, crystalline, odorless solid, that decomposes above 290 °C.[2]

Disodium oxalate can act as a reducing agent, and it may be used as a primary standard for standardizing potassium permanganate (KMnO4) solutions.

The mineral form of sodium oxalate is natroxalate. It is only very rarely found and restricted to extremely sodic conditions of ultra-alkaline pegmatites.[4]

Preparation[edit]

Sodium oxalate can be prepared through the neutralization of oxalic acid with sodium hydroxide (NaOH) in a 1:2 acid-to-base molar ratio. Evaporation yields the anhydrous oxalate[5] that can be thoroughly dried by heating to between 200 and 250 °C.[2]

Half-neutralization can be accomplished with NaOH in a 1:1 ratio which produces NaHC2O4, monobasic sodium oxalate or sodium hydrogenoxalate.

Alternatively, it can be produced by decomposing sodium formate by heating it at a temperature exceeding 360 °C.[citation needed]

Reactions[edit]

Sodium oxalate starts to decompose above 290 °C into sodium carbonate and carbon monoxide:[2]

Na
2
C
2
O
4
Na
2
CO
3
+ CO

When heated at between 200 and 525°C with vanadium pentoxide in a 1:2 molar ratio, the above reaction is suppressed, yielding instead a sodium vanadium oxibronze with release of carbon dioxide[6]

x Na
2
C
2
O
4
+ 2 V
2
O
5
→ 2 Na
x
V
2
O
5
+ 2x CO
2

with x increasing up to 1 as the temperature increases.

Sodium oxalate is used to standardize potassium permanganate solutions. It is desirable that the temperature of the titration mixture is greater than 60 °C to ensure that all the permanganate added reacts quickly. The kinetics of the reaction is complex, and the manganese(II) ions formed catalyze the further reaction between permanganate and oxalic acid (formed in situ by the addition of excess sulfuric acid). The final equation is as follows:[7]

5 Na2C2O4 + 2 KMnO4 + 8 H2SO4 → K2SO4 + 5 Na2SO4 + 2 MnSO4 + 10 CO2 + 8 H2O

Biological activity[edit]

Like several other oxalates, sodium oxalate is toxic to humans. It can cause burning pain in the mouth, throat and stomach, bloody vomiting, headache, muscle cramps, cramps and convulsions, drop in blood pressure, heart failure, shock, coma, and possible death. Mean lethal dose by ingestion of oxalates is 10-15 grams/kilogram of body weight (per MSDS).

Sodium oxalate, like citrates, can also be used to remove calcium ions (Ca2+) from blood plasma. It also prevents blood from clotting. Note that by removing calcium ions from the blood, sodium oxalate can impair brain function, and deposit calcium oxalate in the kidneys.

References[edit]

  1. ^ http://chem.sis.nlm.nih.gov/chemidplus/rn/62-76-0
  2. ^ a b c d Yoshimori T1, Asano Y, Toriumi Y, Shiota T. (1978) "Investigation on the drying and decomposition of sodium oxalate". Talanta, volume 25, issue 10, pages 603-605. doi:10.1016/0039-9140(78)80158-1
  3. ^ http://chem.sis.nlm.nih.gov/chemidplus/rn/62-76-0
  4. ^ http://rruff.geo.arizona.edu/doclib/hom/natroxalate.pdf Handbook of Mineralogy
  5. ^ H. W. Foote and John E. Vance (1933), "The system; sodium iodate, sodium oxalate, water". American Journal of Science, series 5, volume 26, issue 151, pages 16-18. doi:10.2475/ajs.s5-26.151.16
  6. ^ D. Ballivet-Tkatchenko, J. Galy, -M. Savariault (1994): "Thermal decomposition of sodium oxalate in the presence of V2O5: Mechanistic approach of sodium oxibronzes formation". Thermochimica Acta, volume 232, issue 2, pages 215-223. doi:10.1016/0040-6031(94)80061-8
  7. ^ Mcbride, R. S. (1912). "The standardization of potassium permanganate solution by sodium oxalate". J. Am. Chem. Soc. 34: 393. doi:10.1021/ja02205a009.