|Preferred IUPAC name
3D model (JSmol)
CompTox Dashboard (EPA)
|Odor||Pungent, irritating, floral, cucumber like|
|Density||0.7845 g/cm3 (25 °C)|
|Melting point||-94.7 °C (-138.5 °F; 178.5 K)|
|Boiling point||56.05 °C (132.89 °F; 329.20 K)|
|Solubility||Miscible in benzene, diethyl ether, methanol, chloroform, ethanol|
Refractive index (nD)
|1.3588 (VD = 54.46)|
|Viscosity||0.295mPa·s (25 °C)|
|Trigonal planar at C2|
|Dihedral at C2|
Heat capacity (C)
Std enthalpy of
Std enthalpy of
|Safety data sheet||See: data page|
|GHS Signal word||Danger|
|H225, H319, H336, H373|
|P210, P235, P260, P305+351+338|
|NFPA 704 (fire diamond)|
|Flash point||-20 °C (-4 °F; 253 K)|
|465 °C (869 °F; 738 K)|
Threshold limit value (TLV)
|1185mg/m3 (TWA), 2375mg/m3 (STEL)|
|Lethal dose or concentration (LD, LC):|
LD50 (median dose)
LC50 (median concentration)
|20,702ppm (rat, 8 h)|
LCLo (lowest published)
|45,455ppm (mouse, 1 h)|
|NIOSH (US health exposure limits):|
|TWA 250ppm (590mg/m3)|
IDLH (Immediate danger)
|Supplementary data page|
|Refractive index (n),|
Dielectric constant (?r), etc.
|UV, IR, NMR, MS|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Acetone is miscible with water and serves as an important solvent in its own right, in industry, home, and laboratory. About 6.7 million tonnes were produced worldwide in 2010, mainly for use as a solvent and production of methyl methacrylate and bisphenol A. It is a common building block in organic chemistry. Familiar household uses of acetone are as the active ingredient in nail polish remover and as paint thinner. It has volatile organic compound (VOC) exempt status in the USA.
Acetone is produced and disposed of in the human body through normal metabolic processes. It is normally present in blood and urine. People with diabetes produce it in larger amounts. Reproductive toxicity tests show that it has low potential to cause reproductive problems. Ketogenic diets that increase ketone bodies (acetone, ?-hydroxybutyric acid and acetoacetic acid) in the blood are used to counter epileptic attacks in infants and children who suffer from refractory epilepsy.
Acetone was first produced by alchemists during the late Middle Ages via the dry distillation of metal acetates (e.g., lead acetate, which produced "spirit of Saturn" (since the alchemical symbol for lead was also the astrological symbol for the planet Saturn)).
In 1832, French chemist Jean-Baptiste Dumas and German chemist Justus von Liebig determined the empirical formula for acetone. In 1833, the French chemist Antoine Bussy named acetone by adding the suffix -one to the stem of the corresponding acid (viz, acetic acid). By 1852, English chemist Alexander William Williamson realized that acetone was methyl acetyl; the following year, the French chemist Charles Frédéric Gerhardt concurred. In 1865, the German chemist August Kekulé published the modern structural formula for acetone. Johann Josef Loschmidt had presented the structure of acetone in 1861, but his privately published booklet received little attention. During World War I, Chaim Weizmann developed the process for industrial production of acetone (Weizmann Process).
In 2010, the worldwide production capacity for acetone was estimated at 6.7 million tonnes per year. With 1.56 million tonnes per year, the United States had the highest production capacity, followed by Taiwan and mainland China. The largest producer of acetone is INEOS Phenol, owning 17% of the world's capacity, with also significant capacity (7-8%) by Mitsui, Sunoco and Shell in 2010. INEOS Phenol also owns the world's largest production site (420,000 tonnes/annum) in Beveren (Belgium). Spot price of acetone in summer 2011 was 1100-1250 USD/tonne in the United States.
Acetone is produced directly or indirectly from propylene. Approximately 83% of acetone is produced via the cumene process; as a result, acetone production is tied to phenol production. In the cumene process, benzene is alkylated with propylene to produce cumene, which is oxidized by air to produce phenol and acetone:
After that time, during World War I, acetone was produced using acetone-butanol-ethanol fermentation with Clostridium acetobutylicum bacteria, which was developed by Chaim Weizmann (later the first president of Israel) in order to help the British war effort, in the preparation of Cordite. This acetone-butanol-ethanol fermentation was eventually abandoned when newer methods with better yields were found.
Like most ketones, acetone exhibits the keto-enol tautomerism in which the nominal "keto" structure is in equilibrium with the "enol" structure . In acetone vapor at ambient temperature, only 0.00000024% of the molecules are in the enol form. Yet the enol form is chemically important in some chemical reactions.
In the presence of suitable catalysts, two acetone molecules also combine to form the compound diacetone alcohol , which on dehydration gives mesityl oxide . This product can further combine with another acetone molecule, with loss of another molecule of water, yielding phorone and other compouds.
One might expect acetone to also form polymers and (possibly cyclic) oligomers of two types. In one type, units could be acetone molecules linked by ether bridges derived by from the opening of the double bond, to give a polyketal-like (PKA) chain [--]n. The other type could be obtained through repeated aldol condensation, with one molecule of water removed at each step, yielding a poly(methylacetylene) (PMA) chain [--]n.
The conversion of acetone to a polyketal (PKA) would be analogous to the formation of paraformaldehyde from formol, and of trithioacetone from thioacetone. In 1960, Kargin, Kabanov and others observed that the thermodynamics of this process in unfavorable for liquid acetone, so that it (unlike thioacetone and formol) is not expected to polymerize spontaneously, even with catalysts. However, they observed that the thermodynamics became favorable for crystalline solid acetone at the melting point (-96 °C). They claimed to have obtained such a polymer (a white elastic solid, soluble in acetone, stable for several hours at room temperature) by depositing vapor of acetone, with some magnesium as a catalyst, onto a very cold surface.
In 1962, Wasaburo Kawai reported the synthesis of a similar product, from liquid acetone cooled to -70 to -78 °C, using n-butyl lithium or triethylaluminium as catalysts. He claimed that the infrared absorption spectrum showed the presence of linkages but no groups. However, conflicting results were obtained later by other investigators.
Small amounts of acetone are produced in the body by the decarboxylation of ketone bodies. Certain dietary patterns, including prolonged fasting and high-fat low-carbohydrate dieting, can produce ketosis, in which acetone is formed in body tissue. Certain health conditions, such as alcoholism and diabetes, can produce ketoacidosis, uncontrollable ketosis that leads to a sharp, and potentially fatal, increase in the acidity of the blood. Since it is a byproduct of fermentation, acetone is a byproduct of the distillery industry.
Although some biochemistry textbooks and current research publications indicate that acetone cannot be metabolized, there is evidence to the contrary, some dating back thirty years. It can then be metabolized either by CYP2E1 via methylglyoxal to D-lactate and pyruvate, and ultimately glucose/energy, or by a different pathway via propylene glycol to pyruvate, lactate, acetate (usable for energy) and propionaldehyde.
Acetone is a good solvent for many plastics and some synthetic fibers. It is used for thinning polyester resin, cleaning tools used with it, and dissolving two-part epoxies and superglue before they harden. It is used as one of the volatile components of some paints and varnishes. As a heavy-duty degreaser, it is useful in the preparation of metal prior to painting or soldering, and to remove rosin flux after soldering, which helps to prevent the rusty bolt effect.
Although itself flammable, acetone is used extensively as a solvent for the safe transportation and storage of acetylene, which cannot be safely pressurized as a pure compound. Vessels containing a porous material are first filled with acetone followed by acetylene, which dissolves into the acetone. One liter of acetone can dissolve around 250 liters of acetylene at a pressure of 10 bar.
The third major use of acetone (about 20%) is synthesizing bisphenol A. Bisphenol A is a component of many polymers such as polycarbonates, polyurethanes, and epoxy resins. The synthesis involves the condensation of acetone with phenol:
Many millions of kilograms of acetone are consumed in the production of the solvents methyl isobutyl alcohol and methyl isobutyl ketone. These products arise via an initial aldol condensation to give diacetone alcohol.
In the laboratory, acetone is used as a polar, aprotic solvent in a variety of organic reactions, such as SN2 reactions. The use of acetone solvent is critical for the Jones oxidation. It does not form an azeotrope with water (see azeotrope tables). It is a common solvent for rinsing laboratory glassware because of its low cost and volatility. Despite its common use as a supposed drying agent, it is not effective except by bulk displacement and dilution. Acetone can be cooled with dry ice to -78 °C without freezing; acetone/dry ice baths are commonly used to conduct reactions at low temperatures. Acetone is fluorescent under ultraviolet light, and its vapor can be used as a fluorescent tracer in fluid flow experiments.
Low-grade acetone is also commonly used in academic laboratory settings as a glassware rinsing agent for removing residue and solids before a final wash. Acetone leaves a small amount of residue on a surface when dried that is harmful to surface samples.
Acetone is used in the field of pathology to find lymph nodes in fatty tissues for tumor staging (such as looking for lymph nodes in the fat surrounding the intestines). This helps dissolve the fat, and hardens the nodes, making finding them easier.
Dermatologists use acetone with alcohol for acne treatments to chemically peel dry skin. Common agents used today for chemical peeling are salicylic acid, glycolic acid, 30% salicylic acid in ethanol, and trichloroacetic acid (TCA). Prior to chemexfoliation, the skin is cleaned and excess fat removed in a process called defatting. Acetone, Septisol, or a combination of these agents is commonly used in this process.
Acetone has been shown to have anticonvulsant effects in animal models of epilepsy, in the absence of toxicity, when administered in millimolar concentrations. It has been hypothesized that the high-fat low-carbohydrate ketogenic diet used clinically to control drug-resistant epilepsy in children works by elevating acetone in the brain. Because of their higher energy requirements, children have higher acetone production than most adults - and the younger the child, the higher the expected production. This indicates that children are not uniquely susceptible to acetone exposure. External exposures are small compared to the exposures associated with the ketogenic diet.
Make-up artists use acetone to remove skin adhesive from the netting of wigs and mustaches by immersing the item in an acetone bath, then removing the softened glue residue with a stiff brush.
Acetone is often used for vapor polishing of printing artifacts on 3D-printed models printed with ABS plastic. The technique, called acetone vapor bath smoothing, involves placing the printed part in a sealed chamber containing a small amount of acetone, and heating to around 80 degrees Celsius for 10 minutes. This creates a vapor of acetone in the container. The acetone condenses evenly all over the part, causing the surface to soften and liquefy. Surface tension then smooths the semi-liquid plastic. When the part is removed from the chamber, the acetone component evaporates leaving a glassy-smooth part free of striation, patterning, and visible layer edges, common features in untreated 3D printed parts.
The most hazardous property of acetone is its extreme flammability. At temperatures greater than acetone's flash point of -20 °C (-4 °F), air mixtures of between 2.5% and 12.8% acetone, by volume, may explode or cause a flash fire. Vapors can flow along surfaces to distant ignition sources and flash back. Static discharge may also ignite acetone vapors, though acetone has a very high ignition initiation energy point and therefore accidental ignition is rare. Even pouring or spraying acetone over red-glowing coal will not ignite it, due to the high concentration of vapour and the cooling effect of evaporation of the liquid. It auto-ignites at 465 °C (869 °F). Auto-ignition temperature is also dependent upon the exposure time, thus at some tests it is quoted as 525 °C. Also, industrial acetone is likely to contain a small amount of water which also inhibits ignition.
When oxidized, acetone forms acetone peroxide as a byproduct, which is a highly unstable, primary high explosive compound. It may be formed accidentally, e.g. when waste hydrogen peroxide is poured into waste solvent containing acetone. Due to its instability, it is rarely used, despite its simple chemical synthesis.
Acetone has been studied extensively and is believed to exhibit only slight toxicity in normal use. There is no strong evidence of chronic health effects if basic precautions are followed. It is generally recognized to have low acute and chronic toxicity if ingested and/or inhaled. Acetone is not currently regarded as a carcinogen, a mutagenic chemical nor a concern for chronic neurotoxicity effects.
Acetone can be found as an ingredient in a variety of consumer products ranging from cosmetics to processed and unprocessed foods. Acetone has been rated as a generally recognized as safe (GRAS) substance when present in beverages, baked foods, desserts, and preserves at concentrations ranging from 5 to 8 mg/L.
Acetone is however an irritant, causing mild skin irritation and moderate to severe eye irritation. At high vapor concentrations, it may depress the central nervous system like many other solvents. In one documented case, ingestion of a substantial amount of acetone led to systemic toxicity, although the patient eventually fully recovered. Some sources estimate LD50 for human ingestion at 0.621 g/kg; Acute toxicity for mice by ingestion (LD50) is 3 g/kg, and by inhalation (LC50) is 44 g/m3 over 4 hours.
In 1995, the United States Environmental Protection Agency (EPA) removed acetone from the list of "toxic chemicals" maintained under Section 313 of the Emergency Planning and Community Right to Know Act (EPCRA). In making that decision, EPA conducted an extensive review of the available toxicity data on acetone and found that acetone "exhibits acute toxicity only at levels that greatly exceed releases and resultant exposures", and further that acetone "exhibits low toxicity in chronic studies".
Although acetone occurs naturally in the environment in plants, trees, volcanic gases, forest fires, and as a product of the breakdown of body fat, the majority of the acetone released into the environment is of industrial origin. Acetone evaporates rapidly, even from water and soil. Once in the atmosphere, it has a 22-day half-life and is degraded by UV light via photolysis (primarily into methane and ethane.) Consumption by microorganisms contributes to the dissipation of acetone in soil, animals, or waterways.
The LD50 of acetone for fish is 8.3 g/L of water (or about 1%) over 96 hours, and its environmental half-life in water is about 1 to 10 days. Acetone may pose a significant risk of oxygen depletion in aquatic systems due to the microbial consumption.
On 30 July 2015, scientists reported that upon the first touchdown of the Philae lander on comet 67P surface, measurements by the COSAC and Ptolemy instruments revealed sixteen organic compounds, four of which were seen for the first time on a comet, including acetamide, acetone, methyl isocyanate, and propionaldehyde.