Formic acid
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| Names | |||
|---|---|---|---|
| IUPAC name
Methanoic acid
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| Other names
Formic acid
Hydrogen carboxylic acid Formylic acid Aminic acid | |||
| Identifiers | |||
3D model (JSmol)
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| ChEBI | |||
| ChemSpider | |||
| ECHA InfoCard | 100.000.527 | ||
| E number | E236 (preservatives) | ||
PubChem CID
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| RTECS number |
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CompTox Dashboard (EPA)
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| Properties | |||
| CH2O2 | |||
| Molar mass | 46.025 g·mol−1 | ||
| Appearance | Colorless, fuming liquid | ||
| Density | 1.22 g/mL, liquid | ||
| Melting point | 8.4 °C (47.1 °F; 281.5 K) | ||
| Boiling point | 101 °C (214 °F; 374 K) | ||
| Miscible | |||
| Acidity (pKa) | 3.744 | ||
| Viscosity | 1.57 cP at 26 °C | ||
| Structure | |||
| Planar | |||
| 1.41 D(gas) | |||
| Hazards | |||
| Occupational safety and health (OHS/OSH): | |||
Main hazards
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Corrosive; irritant; sensitizer. | ||
| NFPA 704 (fire diamond) | |||
| Flash point | 69 °C (156 °F) | ||
| Related compounds | |||
| Supplementary data page | |||
| Formic acid (data page) | |||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Formic acid (systematically called methanoic acid) is the simplest carboxylic acid. Its formula is HCOOH or HCO2H. It is an important intermediate in chemical synthesis and occurs naturally, most notably in the venom of bee and ant stings. In fact, its name comes from the Latin word for ant, formica, referring to its early isolation by the distillation of ant bodies. Esters, salts and the anion derived from formic acid are referred to as formate.
Properties
Formic acid is miscible with water and most polar organic solvents, and somewhat soluble in hydrocarbons.
Reactions
Formic acid shares most of the chemical properties of other carboxylic acids. Reflecting its high acidity, its solutions in alcohols form esters spontaneously. Formic acid shares some of the reducing properties of aldehydes, reducing solutions of gold, silver, and platinum to the metals.
Decomposition
Heat and especially acids cause formic acid to decompose to carbon monoxide and water. Treatment of formic acid with sulfuric acid is a convenient laboratory source of CO.
In the presence of platinum, it decomposes with release of hydrogen and carbon monoxide. Soluble ruthenium catalysts are also effective.[1].[2]. Carbon monoxide free hydrogen has been generated in a very wide pressure range (1-600 bar). Formic acid has even been considered as a material for hydrogen storage.[3] The co-product of this decomposition, carbon dioxide, can be rehydrogenated back to formic acid in a second step. Formic acid contains 53 g L−1 hydrogen at room temperature and atmospheric pressure, which is twice as much as compressed hydrogen gas can attain at 350 bar pressure. Pure formic acid is a liquid with a flash point - ignition temperature of + 69 °C, much higher than that of gasoline (– 40 °C) or ethanol (+ 13 °C).
Addition to alkenes
Formic acid is unique among the carboxylic acids in its ability to participate in addition reactions with alkenes. Formic acids and alkenes readily react to form formate esters. In the presence of certain acids, including sulfuric and hydrofluoric acids, however, a variant of the Koch reaction occurs instead, and formic acid adds to the alkene to produce a larger carboxylic acid.
Formic acid anhydride
The reaction of formyl fluoride with sodium formate affords formic anhydride. The other three methods involved the use of dehydrating agents. The formic anhydride, however, is only stable in ethereal solution and decomposes upon attempted distillation.[4]
Production
From methyl formate and formamide
When methanol and carbon monoxide are combined in the presence of a strong base, the formic acid derivative methyl formate results, according to the chemical equation:[5]
- CH3OH + CO → HCO2CH3
In industry, this reaction is performed in the liquid phase at elevated pressure. Typical reaction conditions are 80 °C and 40 atm. The most widely-used base is sodium methoxide. Hydrolysis of the methyl formate produces formic acid:
- HCO2CH3 + H2O → HCO2H + CH3OH
Efficient hydrolysis of methyl formate requires a large excess of water. Some routes proceed indirectly by first treating the methyl formate with ammonia to give formamide, which is then hydrolyzed with sulfuric acid:
- HCO2CH3 + NH3 → HC(O)NH2 + CH3OH
- 2 HC(O)NH2 + 2 H2O + H2SO4 → 2HCO2H + (NH4)2SO4
This approach suffers from the need to dispose of the ammonium sulfate byproduct. This problem has led some manufacturers to develop energy efficient means for separating formic acid from the large excess amount of water used in direct hydrolysis. In one of these processes (used by BASF) the formic acid is removed from the water via liquid-liquid extraction with an organic base.
By-product of acetic acid production
A significant amount of formic acid is produced as a byproduct in the manufacture of other chemicals. Acetic acid once was produced on a large scale by oxidation of alkanes, via a process that cogenerates significant formic acid. This oxidative route to acetic acid is declining in importance, so that the aforementioned dedicated routes to formic acid have become more important.
Hydrogenation of carbon dioxide
The catalytic hydrogenation of CO2 has long been studied. This reaction can be conducted homogeneously.[6][7]
Laboratory methods
In the laboratory, formic acid can be obtained by heating oxalic acid in anhydrous glycerol and extraction by steam distillation. Another preparation (which must be performed under a fume hood) is the acid hydrolysis of ethyl isonitrile (C2H5NC) using HCl solution.[8]
- C2H5NC + 2 H2O → C2H5NH2 + HCO2H
The isonitrile can be obtained by reacting ethyl amine with chloroform (note that the fume hood is required because of the overpoweringly objectionable odor of the isonitrile).
Uses
The principal use of formic acid is as a preservative and antibacterial agent in livestock feed. When sprayed on fresh hay or other silage, it arrests certain decay processes and causes the feed to retain its nutritive value longer, and so it is widely used to preserve winter feed for cattle. In the poultry industry, it is sometimes added to feed to kill salmonella bacteria.
Reagent in organic chemistry
Formic acid is a source for a formyl group for example in the formylation of methylaniline to N-methylformanilide in toluene.[9] In synthetic organic chemistry, formic acid is often used as a source of hydride ion. The Eschweiler-Clarke reaction and the Leuckart-Wallach reaction are examples of this application. It or more commonly its azeotrope with triethylamine, is also used as a source of hydrogen in transfer hydrogenation.
Other uses
- It is used to process organic latex (sap) into raw rubber.
- Beekeepers use formic acid as a miticide against the Tracheal (Acarapis woodi) mite and the Varroa mite.
- It is of minor importance in the textile industry and for the tanning of leather.
- Some formate esters are artificial flavorings or perfumes.
- It is the active ingredient in some brands of household limescale remover.
- It is used in laboratories as a solvent modifier for HPLC and CE separations of proteins and peptides, especially when the sample is being prepared for mass spectrometry analysis.
- Fuel cells that use modified formic acid have been reported.
History
Some alchemists and naturalists were aware that ant hills gave off an acidic vapor as early as the 15th century. The first person to describe the isolation of this substance (by the distillation of large numbers of ants) was the English naturalist John Ray, in 1671. Ants secrete the formic acid for attack and defense purposes. Formic acid was first synthesized from hydrocyanic acid by the French chemist Joseph Gay-Lussac. In 1855 another French chemist, Marcellin Berthelot, developed a synthesis from carbon monoxide that is similar to that used today.
Formic acid was long considered a chemical compound of only minor industrial interest in the chemical industry. In the late 1960s, however, significant quantities of it became available as a byproduct of acetic acid production. It now finds increasing use as a preservative and antibacterial in livestock feed.
In nature
In nature, it is found in the stings and bites of many insects of the order Hymenoptera, mainly ants and is also present in stinging nettles[citation needed]. It is also a significant combustion product resulting from alternative fueled vehicles burning methanol (and ethanol, if contaminated with water) when mixed with gasoline.[citation needed]
Safety
85 % formic acid is not inflammable and diluted formic acid is on the US Food and Drug Administration list of food additives[10]. The principal danger from formic acid is from skin or eye contact with the concentrated liquid or vapors. The US OSHA Permissible Exposure Level (PEL) of formic acid vapor in the work environment is 5 parts per million parts of air (ppm).
Formic acid is readily metabolized and eliminated by the body. Nonetheless, it has specific toxic effects; the formic acid and formaldehyde produced as metabolites of methanol are responsible for the optic nerve damage causing blindness seen in methanol poisoning.[11] Some chronic effects of formic acid exposure have been documented. Some animal experiments have demonstrated it to be a mutagen[citation needed], and chronic exposure may cause liver or kidney damage. Another affect of chronic exposure is development of a skin allergy that manifests upon re-exposure to the chemical.
Concentrated formic acid slowly decomposes to carbon monoxide and water, leading to pressure buildup in the container it is kept in. For this reason 98 % formic acid is shipped in plastic bottles with self-venting caps.
The hazards of solutions of formic acid depend on the concentration. The following table lists the EU classification of formic acid solutions:
| Concentration (weight percent) | Classification | R-Phrases |
|---|---|---|
| 2%–10% | Irritant (Xi) | Template:R36/38 |
| 10%–90% | Corrosive (C) | Template:R34 |
| >90% | Corrosive (C) | Template:R35 |
An assay for formic acid in body fluids, designed for determination of formate after methanol poisoning, is based on the reaction of formate with bacterial formate dehydrogenase.[12]
References
- ^ C. Fellay, P. J. Dyson, G. Laurenczy, A Viable Hydrogen-Storage System Based On Selective F SEX SEX SEX SEX SEX SEX SEXormic Acid Decomposition with a Ruthenium Catalyst, Angew. Chem. Int. Ed., 2008, 47, 3966–3970.
- ^ G. Laurenczy, C. Fellay, P. J. Dyson, Hydrogen production from formic acid. PCT Int. Appl. (2008), 36pp. CODEN: PIXXD2 WO 2008047312 A1 20080424 AN 2008:502691
- ^ F. Joó, Breakthroughs in Hydrogen Storage – Formic Acid as a Sustainable Storage Material for Hydrogen, ChemSusChem 2008, 1, 805–808.
- ^ George A. Olah; Yashwant D. Vankar; Massoud Arvanaghi; Jean Sommer (1979). "Formic Anhydride". Angew. Chem. Int. Ed. Engl. 18 (8): 614. doi:10.1002/anie.197906141.
{{cite journal}}: CS1 maint: multiple names: authors list (link) - ^ Cite error: The named reference
Ullmann_2009was invoked but never defined (see the help page). - ^ P. G. Jessop, in Handbook of Homogeneous Hydrogenation (Eds.: J. G. de Vries, C. J. Elsevier), Wiley-VCH, Weinheim, Germany, 2007, pp. 489–511.
- ^ P. G. Jessop, F. Joó, C.-C. Tai, Recent advances in the homogeneous hydrogenation of carbon dioxide, Coord. Chem. Rev., 2004, 248, 2425–2442.doi:10.1016/j.ccr.2004.05.019
- ^ Cohen, Julius B.: Practical Organic Chemistry MacMillan 1930
- ^ L. F. Fieser and J. E. Jones (1955). Organic Syntheses http://www.orgsyn.org/demo.aspx?prep=cv3p0590
{{cite journal}}: Missing or empty|title=(help); Collected Volumes, vol. 3, p. 590. - ^ US Code of Federal Regulations: 21 CFR 186.1316, 21 CFR 172.515
- ^ "Methanol and Blindness". Ask A Scientist, Chemistry Archive. Retrieved 22 May 2007.
{{cite web}}: Unknown parameter|dateformat=ignored (help) - ^ Makar AB, McMartin KE, Palese M, Tephly TR (1975). "Formate assay in body fluids: application in methanol poisoning". Biochem Med. 13 (2): 117–26. doi:10.1016/0006-2944(75)90147-7. PMID 1.
{{cite journal}}: CS1 maint: multiple names: authors list (link)
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