# Solvent

(Redirected from Polar solvents)
A bottle of acetic acid, a liquid solvent

A solvent (from the Latin solvō, "loosen, untie, solve") is a substance that dissolves a solute (a chemically distinct liquid, solid or gas), resulting in a solution. A solvent is usually a liquid but can also be a solid, a gas, or a supercritical fluid. The 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) and in perfumes (ethanol). Water is a solvent for polar molecules and the most common solvent used by living things; all the ions and proteins in a cell are dissolved in water within a cell. Solvents find various applications in chemical, pharmaceutical, oil, and gas industries, including in chemical syntheses and purification processes.

## Solutions and solvation

When one substance is dissolved into another, a solution is formed.[1] This is opposed to the situation when the compounds are insoluble like sand in water. In a solution, all of the ingredients are uniformly distributed at a molecular level and no residue remains. A solvent-solute 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 and other types of non-solution mixtures. The ability of one compound to be dissolved in another is known as solubility; if this occurs in all proportions, it is called miscible.

In addition to mixing, the substances in a solution interact with each other at the molecular level. When something is dissolved, molecules of the solvent arrange around molecules of the solute. Heat transfer is involved and entropy is increased making the solution more thermodynamically stable than the solute and solvent separately. This arrangement is mediated by the respective chemical properties of the solvent and solute, such as hydrogen bonding, dipole moment and polarizability.[2] Solvation does not cause a chemical reaction or chemical configuration changes in the solute. However, solvation resembles a coordination complex formation reaction, often with considerable energetics (heat of solvation and entropy of solvation) and is thus far from a neutral process.

## Solvent classifications

Solvents can be broadly classified into two categories: polar and non-polar. A special case is mercury, whose solutions are known as amalgams; also, other metal solutions exist which are liquid at room temperature. Generally, the dielectric constant of the solvent provides a rough measure of a solvent's polarity. The strong polarity of water is indicated by its high dielectric constant of 88 (at 0 °C).[3] Solvents with a dielectric constant of less than 15 are generally considered to be nonpolar.[4] The dielectric constant measures the solvent's tendency to partly cancel the field strength of the electric field of a charged particle immersed in it. This reduction is then compared to the field strength of the charged particle in a vacuum.[4] Heuristically, the dielectric constant of a solvent can be thought of as its ability to reduce the solute's effective internal charge. Generally, the dielectric constant of a solvent is an acceptable predictor of the solvent's ability to dissolve common ionic compounds, such as salts.

### 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. Most of these measures are sensitive to chemical structure.

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.[5]

Donor number and donor acceptor scale measures polarity in terms of how a solvent interacts with specific substances, like a strong Lewis acid or a strong Lewis base.[6]

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 including Reichardt's dye, nitroaniline and diethylnitroaniline. Another option, Hansen's parameters, separate the cohesive energy density into dispersion, polar and hydrogen bonding contributions.

### Polar protic and polar aprotic

Solvents with a dielectric constant (more accurately, relative static permittivity) greater than 15 (i.e. polar or polarizable) 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.[7] 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[8] Dielectric constant[9] 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 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-\
(${\displaystyle {\ce {(C2H4O)2}}}$)
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
Dichloromethane (DCM) CH2Cl2 40 °C 9.1 1.3266 g/ml 1.60 D

#### Polar aprotic solvents

Tetrahydrofuran (THF)
/-CH2-CH2-O-CH2-CH2-\
(${\displaystyle {\ce {(CH2)4O}}}$)
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
Nitromethane CH3-NO2 100–103 °C 35.87 1.1371 g/ml 3.56 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
Water
${\displaystyle {\ce {H2O}}}$
100 °C 80 1.000 g/ml 1.85 D

### Hansen solubility parameter values

The Hansen solubility parameter values[10][11] 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

#### 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
1,4-Dioxane /-CH2-CH2-O-CH2-CH2-O-\ 17.5 1.8 9.0

#### Polar aprotic solvents

Ethyl acetate CH3-C(=O)-O-CH2-CH3 15.8 5.3 7.2
Tetrahydrofuran (THF) /-CH2-CH2-O-CH2-CH2-\ 16.8 5.7 8.0
Dichloromethane CH2Cl2 17.0 7.3 7.1
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

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.[10][11] 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.

### Boiling point

Solvent Boiling point (°C)[8]
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 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.

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

### Density

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.

Solvent Specific gravity[12]
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 1.000
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 1.325
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

## Safety

### Fire

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.

In addition some solvents, such as methanol, can burn with a very hot flame which can be nearly invisible under some lighting conditions.[13][14] This can delay or prevent the timely recognition of a dangerous fire, until flames spread to other materials.

### Explosive 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 likely to form such peroxides than diethyl ether. One of the most susceptible solvents is diisopropyl ether, but all ethers are considered to be potential peroxide sources.

The heteroatom (oxygen) stabilizes the formation of a free radical which is formed by the abstraction of a hydrogen atom by another free radical.[clarification needed] The carbon-centred free radical thus formed is able to react with an oxygen molecule to form a peroxide compound. The process of peroxide formation is greatly accelerated by exposure to even low levels of light, but can proceed slowly even in dark conditions.

Unless a desiccant is used which 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, shock-sensitive solid precipitate. Minor mechanical disturbances, such as scraping the inside of a vessel or the dislodging of a deposit, can set off a violent explosion.

In some cases, the peroxides can spontaneously precipitate as concentrated crystalline solids within a storage container for ether. When this solid is formed at the mouth of the bottle, merely twisting the cap may provide sufficient energy for the peroxide to detonate. Peroxide formation is not a significant problem when fresh solvents are used up quickly; they are more of a problem in laboratories which may take years to finish a single bottle. Low-volume users should acquire only small amounts of peroxide-prone solvents, and dispose of old solvents on a regular periodic schedule.

Ethers are best stored in the dark in closed canisters, in the presence of stabilizers like butylated hydroxytoluene (BHT), or over sodium hydroxide desiccant.[citation needed]

A number of tests can be used to detect the presence of a peroxide in an ether; one is to use a combination of iron(II) sulfate and potassium thiocyanate. The peroxide is able to oxidize the Fe2+ ion to an Fe3+ ion, which then forms a deep-red coordination complex with the thiocyanate.

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, and must be disposed of properly. The advantage of using sodium/benzophenone is that moisture and oxygen are removed as well.

## Health effects

General health hazards associated with solvent exposure include toxicity to the nervous system, reproductive damage, liver and kidney damage, respiratory impairment, cancer, and dermatitis.[15]

### Acute exposure

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. Fraudulent substitution of 1,5-pentanediol for the psychoactive 1,4-butanediol by a subcontractor caused the Bindeez product recall.

If ingested, alcohols (other than ethanol) such as methanol, propanol, and ethylene glycol metabolize into toxic aldehydes and acids, which cause potentially fatal metabolic acidosis. Thus, the commonly available alcohol solvent methanol can cause permanent blindness or death if ingested. The solvent 2-butoxyethanol, used in fracking fluids, can cause hypotension and metabolic acidosis.[16]

### Chronic exposure

Some solvents including chloroform and benzene (a common ingredient in gasoline) are known to be carcinogenic, while many others are considered by the World Health Organization to be likely carcinogens. Solvents can damage internal organs like the liver, the kidneys, the nervous system, or the brain. The cumulative effects of long-term or repeated exposure to solvents are called chronic solvent-induced encephalopathy (CSE).

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.[17] 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.

Many solvents are known or suspected to be cataractogenic, greatly increasing the risk of developing cataracts in the lens of the eye.[18] Solvent exposure has also been associated with neurotoxic damage causing hearing loss[19][20] and color vision losses.[21]

### 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; this is particularly a health risk if aquifers are affected. There may be about 5000 sites worldwide that have major subsurface solvent contamination.[citation needed]

## References

1. ^ Tinoco, Ignacio; Sauer, Kenneth and Wang, James C. (2002) Physical Chemistry Prentice Hall p. 134 ISBN 0-13-026607-8
2. ^ Lowery and Richardson, pp. 181–183
3. ^ Malmberg, C. G.; Maryott, A. A. (January 1956). "Dielectric Constant of Water from 0° to 100°C" (PDF). Journal of Research of the National Bureau of Standards. 56 (1): 1. doi:10.6028/jres.056.001. Retrieved 27 June 2014.
4. ^ a b Lowery and Richardson, p. 177.
5. ^ Kosower, E.M. (1969) "An introduction to Physical Organic Chemistry" Wiley: New York, p. 293
6. ^ Gutmann, V. (1976). "Solvent effects on the reactivities of organometallic compounds". Coord. Chem. Rev. 18 (2): 225. doi:10.1016/S0010-8545(00)82045-7.
7. ^ Lowery and Richardson, p. 183.
8. ^ a b Solvent Properties – Boiling Point. Xydatasource.com. Retrieved on 2013-01-26.
9. ^ Dielectric Constant. Macro.lsu.edu. Retrieved on 2013-01-26.
10. ^ a b Abbott, Steven and Hansen, Charles M. (2008) Hansen Hansen Solubility Parameters in Practice, ISBN 0-9551220-2-3
11. ^ a b Hansen, Charles M. (2007) Hansen solubility parameters: a user's handbook CRC Press, ISBN 0-8493-7248-8
12. ^ Selected solvent properties – Specific Gravity. Xydatasource.com. Retrieved on 2013-01-26.
13. ^ Fanick, E. Robert; Smith, Lawrence R.; Baines, Thomas M. (1984-10-01). "Safety Related Additives for Methanol Fuel". Warrendale, PA.
14. ^ Anderson, J. E.; Magyarl, M. W.; Siegl, W. O. (1985-07-01). "Concerning the Luminosity of Methanol-Hydrocarbon Diffusion Flames". Combustion Science and Technology. 43 (3-4): 115–125. ISSN 0010-2202. doi:10.1080/00102208508947000.
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16. ^ Hung, Tawny; Dewitt, Christopher R.; Martz, Walter; Schreiber, William; Holmes, Daniel Thomas (July 2010). "Fomepizole fails to prevent progression of acidosis in 2-Butoxyethanol and ethanol coingestion". Clinical Toxicology. 48 (6): 569–571. PMID 20560787. doi:10.3109/15563650.2010.492350.
17. ^ Lundberg I, Gustavsson A, Högberg M, Nise G (1992). "Diagnoses of alcohol abuse and other neuropsychiatric disorders among house painters compared with house carpenters". Br J Ind Med. 49 (6): 409–15. PMC . PMID 1606027. doi:10.1136/oem.49.6.409.
18. ^ Raitta, C; Husman, K; Tossavainen, A (1976). "Lens changes in car painters exposed to a mixture of organic solvents". Albrecht von Graefes Archiv für klinische und experimentelle Ophthalmologie. 200 (2): 149–56. PMID 1086605. doi:10.1007/bf00414364.
19. ^ Campo, Pierre; Morata, Thais C.; Hong, OiSaeng. "Chemical exposure and hearing loss". Disease-a-Month. 59 (4): 119–138. PMC . PMID 23507352. doi:10.1016/j.disamonth.2013.01.003.
20. ^ Johnson AC and Morata,, TC (2010). "Occupational exposure to chemicals and hearing impairment. The Nordic Expert Group for Criteria Documentation of Health Risks from Chemicals." (PDF). Arbete och Hälsa. 44: 177.
21. ^ Mergler, D; Blain, L; Lagacé, J. P. (1987). "Solvent related colour vision loss: An indicator of neural damage?". International Archives of Occupational and Environmental Health. 59 (4): 313–21. PMID 3497110. doi:10.1007/bf00405275.