Formic acid

Formic acid
General
Systematic name	methanoic acid
Other names	formic acid hydrogen carboxylic acid aminic acid
Molecular formula	H-COOH
Molar mass	46.03 g/mol
Appearance	colourless, fuming liquid
CAS number	[64-18-6]
Properties
Density and phase	1.22 g/cm3, liquid
Solubility in water	miscible
Melting point	8.4 °C (281.5 K)
Boiling point	100.8 °C (373.9 K)
Acidity (pKa)	3.75
Viscosity	1.57 cP at 26 °C
Structure
Molecular shape	planar
Dipole moment	1.41 D(gas)
Hazards
MSDS	External MSDS
Main hazards	corrosive
NFPA 704
Flash point	69 °C
R/S statement	R: 35 S: 1/2, 23, 26, 45
RTECS number	LQ4900000
Supplementary data page
Structure & properties	n, εr, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
Related compounds
Related carboxylic acids	acetic acid propionic acid
Related compounds	formaldehyde methanol
Except where noted otherwise, data are given for materials in their standard state (at 25°C, 100 kPa)

Formic acid (systematically called methanoic acid) is the simplest carboxylic acid. Its formula is CH₂O₂ or HCOOH. Its structure is shown at right.

In nature, it is found in the stings and bites of many insects of the order Hymenoptera, including bees and ants. It is also the principal irritant in the leaves of the stinging nettle. It is also a significant combustion product resulting from alternative fueled vehicles burning methanol (and ethanol, if contaminated with water) when mixed with gasoline. Its name comes from the Latin word for ant, formica, referring to its early isolation by the distillation of ant bodies. A chemical compound such as a salt from the neutralization of formic acid with a base, or an ester derived from formic acid, is referred to as formate (or methanoate). The formate ion has the formula HCOO^-.

History

As early as the 15th century, some alchemists and naturalists were aware that ant hills gave off an acidic vapor. The first person to describe the isolation of this substance (by the distillation of large numbers of dead ants) was the English naturalist John Ray, in 1671. It 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.

In the chemical industry, formic acid was long considered a chemical compound of only minor industrial interest. In the late 1960s, however, significant quantities of it became available as a byproduct of acetic acid production. With its growing use as a preservative and antibacterial in livestock feed, it is now produced for its own sake.

Properties

Formic acid is miscible with water and most polar organic solvents, and somewhat soluble in hydrocarbons. In hydrocarbons and in the vapor phase, it actually consists of hydrogen bonded dimers rather than individual molecules. In the gas phase, this hydrogen bonding results in severe deviations from the ideal gas law. Liquid and solid formic acid consists of an effectively infinite network of hydrogen bonded formic acid molecules.

Formic acid shares most of the chemical properties of other carboxylic acids, although under normal conditions it will not form either an acyl chloride or an acid anhydride. Until very recently, all attempts to form either of these derivatives have resulted in carbon monoxide instead. It has now been shown that the anhydride may be produced by reaction of formyl fluoride with sodium formate at -78^oC, and the chloride by passing HCl into a solution of 1-formimidazole in monochloromethane at -60^oC^[1]. Heat can also cause formic acid to decompose to carbon monoxide and water. Formic acid shares some of the reducing properties of aldehydes.

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 takes place instead, and formic acid adds to the alkene to produce a larger carboxylic acid.

Most simple formate salts are water soluble.

Production

A significant amount of formic acid is produced as a byproduct in the manufacture of other chemicals, especially acetic acid. However, this production is insufficient to meet the present demand for formic acid, and some formic acid must be produced for its own sake.

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

CH₃OH + CO → HCOOCH₃

Industrially, 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:

HCOOCH₃ + H₂O → HCOOH + CH₃OH

However, direct hydrolysis of methanol requires a large excess of water to proceed efficiently, and so some producers perform it by an indirect route, first reacting the methyl formate with ammonia to produce formamide, and then hydrolyzing the formamide with sulfuric acid to produce formic acid:

HCOOCH₃ + NH₃ → HCONH₂ + H₂O

HCONH₂ + H₂O + ½H₂SO₄ → HCOOH + ½ (NH₄)₂SO₄

This technique has problems of its own, particularly disposing of the ammonium sulfate byproduct, and so recently, some manufacturers have developed energy-efficient means of separating formic acid from the large excess of water used in direct hydrolysis. In one of these processes, used by BASF, the formic acid is removed from the water by liquid extraction with an organic base.

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 use to preserve winter feed for cattle. In the poultry industry, it is sometimes added to feed to kill salmonella bacteria.

Some beekeepers use formic acid as a miticide against the Varroa mite.

Formic acid is of minor importance in the textile industry and in the tanning of leather. Some formate esters are artificial flavorings or perfumes.

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 famous examples of this.

Safety

The principle danger from formic acid is from skin or eye contact with liquid formic acid or with the concentrated vapors. Any of these exposure routes can cause severe chemical burns, and eye exposure can result in permanent eye damage. Inhaled vapors may similarly cause irritation or burns in the respiratory tract. Since carbon monoxide may be also be present in formic acid vapors, care should be taken wherever large quantities of formic acid fumes are present. 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, some chronic effects have been documented. Some animal experiments have demonstrated it to be a mutagen, and chronic exposure may cause liver or kidney damage. Another possibility with chronic exposure is development of a skin allergy that manifests upon re-exposure to the chemical.

External links

• Carbon monoxide as reagent in the formylation of aromatic compounds
• MSDS (Material Safety Data Sheet)
• European Chemicals Bureau
• International Chemical Safety Card 0485
• NIOSH Pocket Guide to Chemical Hazards
• Computational Chemistry Wiki