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A solvent is a liquid, solid, or gas that dissolves another solid, liquid, or gaseous solute, resulting in a solution that is soluble in a certain volume of solvent at a specified temperature. Common uses for organic solvents are in dry cleaning (e.g. tetrachloroethylene), as paint thinners (e.g. toluene, turpentine), as nail polish removers and glue solvents (acetone, methyl acetate, ethyl acetate), in spot removers (e.g. hexane, petrol ether), in detergents (citrus terpenes), in perfumes (ethanol), and in chemical synthesis. The use of inorganic solvents (other than water) is typically limited to research chemistry and some technological processes.

In 2005, the worldwide market for solvents had a total volume of around 17.9 million tons, which led to a turnover of about 8 billion Euro.[citation needed]


Solutions and solvation

When one substance is dissolved into another, a solution is formed.[1] This is opposed to a mixture where one compound is added to another and no chemical bond is formed;[citation needed] a way to think of mixtures and solutions is to compare a cup of water with sand mixed in versus a soda where all of the ingredients are uniform to create a new substance. No residue is left in the bottom. The mixing is referred to as miscibility, whereas the ability to dissolve one compound into another is known as solubility. However, in addition to mixing, both substances in the solution can interact with each other in specific ways. Solvation describes these interactions. When something is dissolved, molecules of the solvent arrange themselves around molecules of the solute. Heat is involved and entropy is increased making the solution more thermodynamically stable than the solute alone. This arrangement is mediated by the respective chemical properties of the solvent and solute, such as hydrogen bonding, dipole moment and polarizability.[2]

Solvent classifications

Solvents can be broadly classified into two categories: polar and non-polar. Generally, the dielectric constant of the solvent provides a rough measure of a solvent's polarity. The strong polarity of water is indicated, at 20 °C, by a dielectric constant of 80.10. Solvents with a dielectric constant of less than 15 are generally considered to be nonpolar.[3] Technically, the dielectric constant measures the solvent's ability to reduce the field strength of the electric field surrounding a charged particle immersed in it. This reduction is then compared to the field strength of the charged particle in a vacuum.[3] In laymen's terms, dielectric constant of a solvent can be thought of as its ability to reduce the solute's internal charge.


Other polarity scales

Dielectric constants are not the only measure of polarity. Because solvents are used by chemists to carry out chemical reactions or observe chemical and biological phenomena, more specific measures of polarity are required.

The Grunwald Winstein mY scale measures polarity in terms of solvent influence on buildup of positive charge of a solute during a chemical reaction.

Kosower's Z scale measures polarity in terms of the influence of the solvent on uv absorption maxima of a salt, usually pyridinium iodide or the pyridinium zwitterion.[4]

Donor number and donar acceptor scale measures polarity in terms of how a solvent interacts with specific substances, like a strong Lewis acid or a strong Lewis base.[5]

The polarity, dipole moment, polarizability and hydrogen bonding of a solvent determines what type of compounds it is able to dissolve and with what other solvents or liquid compounds it is miscible. As a rule of thumb, polar solvents dissolve polar compounds best and non-polar solvents dissolve non-polar compounds best: "like dissolves like". Strongly polar compounds like sugars (e.g. sucrose) or ionic compounds, like inorganic salts (e.g. table salt) dissolve only in very polar solvents like water, while strongly non-polar compounds like oils or waxes dissolve only in very non-polar organic solvents like hexane. Similarly, water and hexane (or vinegar and vegetable oil) are not miscible with each other and will quickly separate into two layers even after being shaken well.

Polar protic and polar aprotic

Solvents with a relative static permittivity greater than 15 can be further divided into protic and aprotic. Protic solvents solvate anions (negatively charged solutes) strongly via hydrogen bonding. Water is a protic solvent. Aprotic solvents such as acetone or dichloromethane tend to have large dipole moments (separation of partial positive and partial negative charges within the same molecule) and solvate positively charged species via their negative dipole.[6] In chemical reactions the use of polar protic solvents favors the SN1 reaction mechanism, while polar aprotic solvents favor the SN2 reaction mechanism.

Physical properties of common solvents

Properties table of common solvents

The solvents are grouped into non-polar, polar aprotic, and polar protic solvents and ordered by increasing polarity. The polarity is given as the dielectric constant. The properties of solvents that exceed those of water are bolded.

Solvent Chemical formula Boiling point[7] Dielectric constant[8] Density Dipole moment
Non-polar solvents
Pentane CH3-CH2-CH2-CH2-CH3 36 °C 1.84 0.626 g/ml 0.00 D
Cyclopentane C5H10 40 °C 1.97 0.751 g/ml 0.00 D
Hexane CH3-CH2-CH2-CH2-CH2-CH3 69 °C 1.88 0.655 g/ml 0.00 D
Cyclohexane C6H12 81 °C 2.02 0.779 g/ml 0.00 D
Benzene C6H6 80 °C 2.3 0.879 g/ml 0.00 D
Toluene C6H5-CH3 111 °C 2.38 0.867 g/ml 0.36 D
1,4-Dioxane /-CH2-CH2-O-CH2-CH2-O-\ 101 °C 2.3 1.033 g/ml 0.45 D
Chloroform CHCl3 61 °C 4.81 1.498 g/ml 1.04 D
Diethyl ether CH3CH2-O-CH2-CH3 35 °C 4.3 0.713 g/ml 1.15 D
Polar aprotic solvents
Dichloromethane (DCM) CH2Cl2 40 °C 9.1 1.3266 g/ml 1.60 D
Tetrahydrofuran (THF) /-CH2-CH2-O-CH2-CH2-\ 66 °C 7.5 0.886 g/ml 1.75 D
Ethyl acetate CH3-C(=O)-O-CH2-CH3 77 °C 6.02 0.894 g/ml 1.78 D
Acetone CH3-C(=O)-CH3 56 °C 21 0.786 g/ml 2.88 D
Dimethylformamide (DMF) H-C(=O)N(CH3)2 153 °C 38 0.944 g/ml 3.82 D
Acetonitrile (MeCN) CH3-C≡N 82 °C 37.5 0.786 g/ml 3.92 D
Dimethyl sulfoxide (DMSO) CH3-S(=O)-CH3 189 °C 46.7 1.092 g/ml 3.96 D
Polar protic solvents
Formic acid H-C(=O)OH 101 °C 58 1.21 g/ml 1.41 D
n-Butanol CH3-CH2-CH2-CH2-OH 118 °C 18 0.810 g/ml 1.63 D
Isopropanol (IPA) CH3-CH(-OH)-CH3 82 °C 18 0.785 g/ml 1.66 D
n-Propanol CH3-CH2-CH2-OH 97 °C 20 0.803 g/ml 1.68 D
Ethanol CH3-CH2-OH 79 °C 30 0.789 g/ml 1.69 D
Methanol CH3-OH 65 °C 33 0.791 g/ml 1.70 D
Acetic acid CH3-C(=O)OH 118 °C 6.2 1.049 g/ml 1.74 D
Water H-O-H 100 °C 80 1.000 g/ml 1.85 D

There's another powerful way to look at these same solvents. By knowing their Hansen solubility parameter values (HSPiP)[9][10], which are based on δD=dispersion bonds, δP=polar bonds and δH=hydrogen bonds, you know important things about their inter-molecular interactions with other solvents and also with polymers, pigments, nanoparticles etc. so you can do two things. First, you can create rational formulations knowing, for example, that there is a good HSP match between a solvent and a polymer. Second, you can make rational substitutions for "good" solvents (they dissolve things well) that are "bad" (for the environment, for health, for cost etc.). The following table shows that the intuitions from "non-polar", "polar aprotic" and "polar protic" are put numerically - the "polar" molecules have higher levels of δP and the protic solvents have higher levels of δH. Because numerical values are used, comparisons can be made rationally by comparing numbers. So acetonitrile is much more polar than acetone but slightly less hydrogen bonding.

Solvent Chemical formula δD Dispersion δP Polar δH Hydrogen bonding
Non-polar solvents
Hexane CH3-CH2-CH2-CH2-CH2-CH3 14.9 0.0 0.0
Benzene C6H6 18.4 0.0 2.0
Toluene C6H5-CH3 18.0 1.4 2.0
Diethyl ether CH3CH2-O-CH2-CH3 14.5 2.9 4.6
Chloroform CHCl3 17.8 3.1 5.7
Ethyl acetate CH3-C(=O)-O-CH2-CH3 15.8 5.3 7.2
Dichloromethane CH2Cl2 17.0 7.3 7.1
Polar aprotic solvents
1,4-Dioxane /-CH2-CH2-O-CH2-CH2-O-\ 17.5 1.8 9.0
Tetrahydrofuran (THF) /-CH2-CH2-O-CH2-CH2-\ 16.8 5.7 8.0
Acetone CH3-C(=O)-CH3 15.5 10.4 7.0
Acetonitrile (MeCN) CH3-C≡N 15.3 18.0 6.1
Dimethylformamide (DMF) H-C(=O)N(CH3)2 17.4 13.7 11.3
Dimethyl sulfoxide (DMSO) CH3-S(=O)-CH3 18.4 16.4 10.2
Polar protic solvents
Acetic acid CH3-C(=O)OH 14.5 8.0 13.5
n-Butanol CH3-CH2-CH2-CH2-OH 16.0 5.7 15.8
Isopropanol CH3-CH(-OH)-CH3 15.8 6.1 16.4
n-Propanol CH3-CH2-CH2-OH 16.0 6.8 17.4
Ethanol CH3-CH2-OH 15.8 8.8 19.4
Methanol CH3-OH 14.7 12.3 22.3
Formic acid H-C(=O)OH 14.6 10.0 14.0
Water H-O-H 15.5 16.0 42.3

Consider a simple example of rational substitution. Suppose for environmental reasons we needed to replace the chlorinated solvent, chloroform, with a solvent (blend) of equal solvency using a mixture of two non-chlorinated solvents from this table. Via trial-and-error, a spreadsheet or some software such as HSPiP[9][10] we find that a 50:50 mix of toluene and 1,4 dioxane is a close match. The δD of the mixture is the average of 18.0 and 17.5 = 17.8. The δP of the mixture is the average of 1.4 and 1.8 = 1.6 and the δH of the mixture is the average of 2.0 and 9.0 = 5.5. So the mixture is 17.8, 1.6, 5.5 compared to Chloroform at 17.8, 3.1, 5.7. Because Toluene is itself has many health issues, other mixtures of solvents can be found using a full Hansen solubility parameter dataset.

Boiling point

Solvent Boiling point (°C)[7]
ethylene dichloride 83.48
pyridine 115.25
methyl isobutyl ketone 116.5
methylene chloride 39.75
isooctane 99.24
carbon disulfide 46.3
carbon tetrachloride 76.75
o-xylene 144.42

An important property of solvents is boiling point. This also determines the speed of evaporation. Small amounts of low-boiling solvents like diethyl ether, dichloromethane, or acetone will evaporate in seconds at room temperature, while high-boiling solvents like water or dimethyl sulfoxide need higher temperatures, an air flow, or the application of vacuum for fast evaporation.

  • Low boilers: boiling temperature below 100 °C (boiling point of water)
  • Medium boilers: between 100 °C and 150 °C
  • High boilers: above 150 °C


Most organic solvents have a lower density than water, which means they are lighter and will form a separate layer on top of water. An important exception: many halogenated solvents like dichloromethane or chloroform will sink to the bottom of a container, leaving water as the top layer. This is important to remember when partitioning compounds between solvents and water in a separatory funnel during chemical syntheses.

Often, specific gravity is cited in place of density. Specific gravity is defined as the density of the solvent divided by the density of water at the same temperature. As such, specific gravity is a unitless value. It readily communicates whether a water-insoluble solvent will float (SG < 1.0) or sink (SG > 1.0) when mixed with water.

Solvent Specific gravity[11]
Pentane 0.626
Petroleum ether 0.656
Hexane 0.659
Heptane 0.684
Diethyl amine 0.707
Diethyl ether 0.713
Triethyl amine 0.728
Tert-butyl methyl ether 0.741
Cyclohexane 0.779
Tert-butyl alcohol 0.781
Isopropanol 0.785
Acetonitrile 0.786
Ethanol 0.789
Acetone 0.790
Methanol 0.791
Methyl isobutyl ketone 0.798
Isobutyl alcohol 0.802
1-Propanol 0.803
Methyl ethyl ketone 0.805
2-Butanol 0.808
Isoamyl alcohol 0.809
1-Butanol 0.810
Diethyl ketone 0.814
1-Octanol 0.826
p-Xylene 0.861
m-Xylene 0.864
Toluene 0.867
Dimethoxyethane 0.868
Benzene 0.879
Butyl acetate 0.882
1-Chlorobutane 0.886
Tetrahydrofuran 0.889
Ethyl acetate 0.895
o-Xylene 0.897
Hexamethylphosphorus triamide 0.898
2-Ethoxyethyl ether 0.909
N,N-Dimethylacetamide 0.937
Diethylene glycol dimethyl ether 0.943­
N,N-Dimethylformamide 0.944
2-Methoxyethanol 0.965
Pyridine 0.982
Propanoic acid 0.993
Water 0.998
2-Methoxyethyl acetate 1.009
Benzonitrile 1.01
1-Methyl-2-pyrrolidinone 1.028
Hexamethylphosphoramide 1.03
1,4-Dioxane 1.033
Acetic acid 1.049
Acetic anhydride 1.08
Dimethyl sulfoxide 1.092
Chlorobenzene 1.1066
Deuterium oxide 1.107
Ethylene glycol 1.115
Diethylene glycol 1.118
Propylene carbonate 1.21
Formic acid 1.22
1,2-Dichloroethane 1.245
Glycerin 1.261
Carbon disulfide 1.263
1,2-Dichlorobenzene 1.306
Methylene chloride 1.326
Nitromethane 1.382
2,2,2-Trifluoroethanol 1.393
Chloroform 1.498
1,1,2-Trichlorotrifluoroethane 1.575
Carbon tetrachloride 1.594
Tetrachloroethylene 1.623


Health and safety


Most organic solvents are flammable or highly flammable, depending on their volatility. Exceptions are some chlorinated solvents like dichloromethane and chloroform. Mixtures of solvent vapors and air can explode. Solvent vapors are heavier than air; they will sink to the bottom and can travel large distances nearly undiluted. Solvent vapors can also be found in supposedly empty drums and cans, posing a flash fire hazard; hence empty containers of volatile solvents should be stored open and upside down.

Both diethyl ether and carbon disulfide have exceptionally low autoignition temperatures which increase greatly the fire risk associated with these solvents. The autoignition temperature of carbon disulfide is below 100°C (212°F), so as a result objects such as steam pipes, light bulbs, hotplates and recently extinguished bunsen burners are able to ignite its vapours.

Peroxide formation

Ethers like diethyl ether and tetrahydrofuran (THF) can form highly explosive organic peroxides upon exposure to oxygen and light, THF is normally more able to form such peroxides than diethyl ether. One of the most susceptible solvents is diisopropyl ether.

The heteroatom (oxygen) stabilizes the formation of a free radical which is formed by the abstraction of a hydrogen atom by another free radical. The carbon centred free radical thus formed is able to react with an oxygen molecule to form a peroxide compound. A range of tests can be used to detect the presence of a peroxide in an ether; one is to use a combination of iron sulfate and potassium thiocyanate. The peroxide is able to oxidize the Fe2+ ion to an Fe3+ ion which then form a deep red coordination complex with the thiocyanate. In extreme cases the peroxides can form crystalline solids within the vessel of the ether.

Unless the desiccant used can destroy the peroxides, they will concentrate during distillation due to their higher boiling point. When sufficient peroxides have formed, they can form a crystalline and shock sensitive solid precipitate. When this solid is formed at the mouth of the bottle, turning the cap may provide sufficient energy for the peroxide to detonate. Peroxide formation is not a significant problem when solvents are used up quickly; they are more of a problem for laboratories which take years to finish a single bottle. Ethers have to be stored in the dark in closed canisters in the presence of stabilizers like butylated hydroxytoluene (BHT) or over sodium hydroxide.

Peroxides may be removed by washing with acidic iron(II) sulfate, filtering through alumina, or distilling from sodium/benzophenone. Alumina does not destroy the peroxides; it merely traps them. The advantage of using sodium/benzophenone is that moisture and oxygen is removed as well.

Health effects

Many solvents can lead to a sudden loss of consciousness if inhaled in large amounts. Solvents like diethyl ether and chloroform have been used in medicine as anesthetics, sedatives, and hypnotics for a long time. Ethanol (grain alcohol) is a widely used and abused psychoactive drug. Diethyl ether, chloroform, and many other solvents (e.g. from gasoline or glues) are used recreationally in glue sniffing, often with harmful long term health effects like neurotoxicity or cancer. Methanol can cause permanent blindness and death.

It is interesting to note that ethanol has a synergistic effect when taken in combination with many solvents. For instance a combination of toluene/benzene and ethanol causes greater nausea/vomiting than either substance alone.

Environmental contamination

A major pathway to induce health effects arises from spills or leaks of solvents that reach the underlying soil. Since solvents readily migrate substantial distances, the creation of widespread soil contamination is not uncommon; there may be about 5000 sites worldwide that have major subsurface solvent contamination; this is particularly a health risk if aquifers are affected.

Chronic health effects

Some solvents including chloroform and benzene (an ingredient of gasoline) are carcinogenic. Many others can damage internal organs like the liver, the kidneys, or the brain.

General precautions

  • Avoid being exposed to solvent vapors by working in a fume hood, or with local exhaust ventilation (LEV), or in a well ventilated area
  • Keep the storage containers tightly closed
  • Never use open flames near flammable solvents; use electrical heating instead
  • Never flush flammable solvents down the drain; read safety data sheets for proper disposal information
  • Avoid the inhalation of solvent vapors
  • Avoid contact of the solvent with the skin — many solvents are easily absorbed through the skin. They also tend to dry the skin and may cause sores and wounds.

See also


  1. ^ Tinoco, Sauer, Wang & Puglisi, Physical Chemistry Prentice Hall 2002 p. 134 ISBN 0130266078
  2. ^ Lowery and Richardson, pp. 181-183
  3. ^ a b Lowery and Richardson, p. 177.
  4. ^ Kosower, E.M. "An introduction to Physical Organic Chemistry" Wiley: New York, 1969 p. 293
  5. ^ Gutmann, V. (1976). "Solvent effects on the reactivities of organometallic compounds". Coord. Chem. Rev. 18: 225. doi:10.1016/S0010-8545(00)82045-7. 
  6. ^ Lowery and Richardson, p. 183.
  7. ^ a b Solvent Properties - Boiling Point
  8. ^ Dielectric Constant
  9. ^ a b Steven Abbott and Charles M. Hansen Hansen Solubility Parameters in Practice, ISBN 0955122023 (2008)
  10. ^ a b Charles M. Hansen Hansen solubility parameters: a user's handbook CRC Press, 2007, ISBN 0849372488
  11. ^ Selected solvent properties - Specific Gravity


  • Lowery, T.H. and Richardson, K.S., Mechanism and Theory in Organic Chemistry, Harper Collins Publishers 3rd ed. 1987 0063640449

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