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A solvent (from the Latin solvō, "I loosen, untie, I solve") is a substance that dissolves a solute (a chemically different liquid, solid or gas), resulting in a solution. A solvent is usually a liquid but can also be a solid or a gas. The maximum quantity of solute that can dissolve in a specific volume of solvent varies with 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), nail polish and in chemical synthesis. The use of inorganic solvents (other than water) is typically limited to research chemistry and some technological processes.
The global solvent market is expected to earn revenues of about US$33 billion in 2019. The dynamic economic development in emerging markets like China, India, Brazil, or Russia will especially continue to boost demand for solvents. Specialists expect the worldwide solvent consumption to increase at an average annual rate of 2.5% over the subsequent years. Accordingly, the growth rate seen during the past eight years will be surpassed.
When one substance is dissolved into another, a solution is formed. This is opposed to the situation when the compounds are insoluble like sand in water. In solution, all of the ingredients are uniformly distributed at a molecular level and no residue remains. The mixture consists of a single phase with all solute molecules occurring as solvates (solvent-solute complexes), as opposed to separate continuous phases as in suspensions, emulsions or other types of mixtures. 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 interact with each other. When something is dissolved, molecules of the solvent arrange 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.
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 0 °C, by a dielectric constant of 88. Solvents with a dielectric constant of less than 15 are generally considered to be nonpolar. 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. In layperson's terms, dielectric constant of a solvent can be thought of as its ability to reduce the solute's effective internal charge.
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.
The Hildebrand parameter is the square root of cohesive energy density. It can be used with nonpolar compounds, but cannot accommodate complex chemistry.
Reichardt's dye, a solvatochromic dye that changes color in response to polarity, gives a scale of ET(30) values. ET is the transition energy between the ground state and the lowest excited state in kcal/mol, and (30) identifies the dye. Another, roughly correlated scale (ET(33)) can be defined with Nile red.
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. Generally, 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.
Polarity can be separated to different contributions. For example, the Kamlet-Taft parameters are dipolarity/polarizability (π*), hydrogen-bonding acidity (α) and hydrogen-bonding basicity (β). These can be calculated from the wavelength shifts of 3–6 different solvatochromic dyes in the solvent, usually Reichardt's dye, nitroaniline and diethylnitroaniline. Another option, Hansen's parameters, separate the cohesive energy density into dispersion, polar and hydrogen bonding contributions.
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. In chemical reactions the use of polar protic solvents favors the SN1 reaction mechanism, while polar aprotic solvents favor the SN2 reaction mechanism.
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||Dielectric constant||Density||Dipole moment|
|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||CH3-CH2-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|
|Propylene carbonate||C4H6O3||240 °C||64.0||1.205 g/ml||4.9 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||24.55||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|
|Nitromethane||CH3-NO2||100–103 °C||35.87||1.1371 g/ml||3.56 D|
|Water||H-O-H||100 °C||80||1.000 g/ml||1.85 D|
The Hansen solubility parameter values are based on dispersion bonds (δD), polar bonds (δP) and hydrogen bonds (δH). These contain information about the inter-molecular interactions with other solvents and also with polymers, pigments, nanoparticles, etc. This allows for rational formulations knowing, for example, that there is a good HSP match between a solvent and a polymer. Rational substitutions can also be made for "good" solvents (effective at dissolving the solute) that are "bad" (expensive or hazardous to health or the environment). 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. For example, acetonitrile is much more polar than acetone but exhibits slightly less hydrogen bonding.
|Solvent||Chemical formula||δD Dispersion||δP Polar||δH Hydrogen bonding|
|Polar aprotic solvents|
|Dimethyl sulfoxide (DMSO)||CH3-S(=O)-CH3||18.4||16.4||10.2|
|Polar protic solvents|
If, for environmental or other reasons, a solvent or solvent blend is required to replace another of equivalent solvency, the substitution can be made on the basis of the Hansen solubility parameters of each. The values for mixtures are taken as the weighted averages of the values for the neat solvents. This can be calculated by trial-and-error, a spreadsheet of values, or HSP software. A 1:1 mixture of toluene and 1,4 dioxane has δD, δP and δH values of 17.8, 1.6 and 5.5, comparable to those of chloroform at 17.8, 3.1 and 5.7 respectively. Because of the health hazards associated with toluene itself, other mixtures of solvents may be found using a full HSP dataset.
|Solvent||Boiling point (°C)|
|methyl isobutyl ketone||116.5|
An important property of solvents is the boiling point. This also determines the speed of evaporation. Small amounts of low-boiling-point solvents like diethyl ether, dichloromethane, or acetone will evaporate in seconds at room temperature, while high-boiling-point solvents like water or dimethyl sulfoxide need higher temperatures, an air flow, or the application of vacuum for fast evaporation.
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: most of the 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.
|Tert-butyl methyl ether||0.741|
|Methyl isobutyl ketone||0.798|
|Methyl ethyl ketone||0.805|
|Diethylene glycol dimethyl ether||0.943|
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 objects such as steam pipes, light bulbs, hotplates and recently extinguished bunsen burners are able to ignite its vapours.
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 are removed as well.
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 also dangerous because it burns with an invisible flame.
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.
Chronic exposure to organic solvents in the work environment can produce a range of adverse neuropsychiatric effects. For example, occupational exposure to organic solvents has been associated with higher numbers of painters suffering from alcoholism.
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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.
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