| Discovery date | 1774 |
| Discovered by | Carl Wilhelm Scheele |
| Origin of the name | The name is derived from the Greek 'chloros', meaning greenish yellow. |
| Allotropes | Cl2 |
Cl
Chlorine
17
35.45
| Group | 17 | Melting point | −101.5°C, −150.7°F, 171.7 K |
| Period | 3 | Boiling point | −34.04°C, −29.27°F, 239.11 K |
| Block | p | Density (g cm−3) | 0.002898 |
| Atomic number | 17 | Relative atomic mass | 35.45 |
| State at 20°C | Gas | Key isotopes | 35Cl, 37Cl |
| Electron configuration | [Ne] 3s23p5 | CAS number | 7782-50-5 |
| ChemSpider ID | 4514529 | ChemSpider is a free chemical structure database | |
Image explanation
The symbol shows a gas mask. This is because chlorine is a toxic gas, and has been used as a chemical weapon. Chlorine is yellowy-green in colour, as is the image.
Appearance
A yellowy-green dense gas with a choking smell.
Uses
Chlorine kills bacteria – it is a disinfectant. It is used to treat drinking water and swimming pool water. It is also used to make hundreds of consumer products from paper to paints, and from textiles to insecticides.
About 20% of chlorine produced is used to make PVC. This is a very versatile plastic used in window frames, car interiors, electrical wiring insulation, water pipes, blood bags and vinyl flooring.
Another major use for chlorine is in organic chemistry. It is used as an oxidising agent and in substitution reactions. 85% of pharmaceuticals use chlorine or its compounds at some stage in their manufacture.
In the past chlorine was commonly used to make chloroform (an anaesthetic) and carbon tetrachloride (a dry-cleaning solvent). However, both of these chemicals are now strictly controlled as they can cause liver damage.
Chlorine gas is itself very poisonous, and was used as a chemical weapon during the First World War.
Biological role
The chloride ion is essential to life. It is mostly present in cell fluid as a negative ion to balance the positive (mainly potassium) ions. It is also present in extra-cellular fluid (eg blood) to balance the positive (mainly sodium) ions.
We get most of the chloride we need from salt. Typical daily salt intake is about 6 grams, but we could manage with half this amount.
Natural abundance
Chlorine is not found uncombined in nature. Halite (sodium chloride or ‘common salt’) is the main mineral that is mined for chlorine. Sodium chloride is a very soluble salt that has been leached into the oceans over the lifetime of the Earth. Several salt beds, or ‘lakes’ are found where ancient seas have evaporated, and these can be mined for chloride.
Chlorine is also found in the minerals carnallite (magnesium potassium chloride) and sylvite (potassium chloride).
40 million tonnes of chlorine gas are made each year from the electrolysis of brine (sodium chloride solution). This process also produces useful sodium hydroxide.
Hydrochloric acid (HCl) was known to the alchemists. The gaseous element itself was first produced in 1774 by Carl Wilhelm Scheele at Uppsala, Sweden, by heating hydrochloric acid with the mineral pyrolusite which is naturally occuring manganese dioxide, MnO2. A dense, greenish-yellow gas was evolved which he recorded as having a choking smell and which dissolved in water to give an acid solution. He noted that it bleached litmus paper, and decolourised leaves and flowers.
Humphry Davy investigated it in 1807 and eventually concluded not only that it was a simple substance, but that it was truly an element. He announced this in 1810 and yet it took another ten years for some chemists finally to accept that chlorine really was an element.
| Atomic radius, non-bonded (Å) | 1.75 | Covalent radius (Å) | 1.00 |
| Electron affinity (kJ mol−1) | 348.575 |
Electronegativity (Pauling scale) |
3.16 |
|
Ionisation energies (kJ mol−1) |
1st
1251.186
2nd
2297.663
3rd
3821.78
4th
5158.608
5th
6541.7
6th
9361.97
7th
11018.221
8th
33603.91
|
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|
| Common oxidation states | 7, 5, 3, 1, -1 | ||||
| Isotopes | Isotope | Atomic mass | Natural abundance (%) | Half life | Mode of decay |
| 35Cl | 34.969 | 75.76 | - | - | |
| 37Cl | 36.966 | 24.24 | - | - | |
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Specific heat capacity (J kg−1 K−1) |
479 | Young's modulus (GPa) | Unknown | |||||||||||
| Shear modulus (GPa) | Unknown | Bulk modulus (GPa) | 1.1 (liquid) | |||||||||||
| Vapour pressure | ||||||||||||||
| Temperature (K) |
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| Pressure (Pa) |
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| Listen to Chlorine Podcast |
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Transcript :
Chemistry in its element: chlorine(Promo) You're listening to Chemistry in its element brought to you by Chemistry World, the magazine of the Royal Society of Chemistry. (End promo) Chris Smith Hello. What's got three isotopes, keeps swimming pools clean, damages the ozone layer and is used in more chemical synthesis reactions than you can shake a benzene ring at. Well the man with the answer is Tim Harrison. Tim Harrison Chlorine is what you might describe as a Jekyll and Hyde element; it is the friend of the synthetic chemist and has found a use in a number of 'nice' applications such as the disinfecting of drinking water and keeping our swimming pools clean. It also has an unpleasant side, being the first chemical warfare agent and taking some of the blame in the depletion of the Earth's ozone layer. Elemental chlorine is a pale, yellowy green gas at room temperature. It was the Greek word khlôros meaning 'yellowish-green' that was used as inspiration by Sir Humphrey Davy when he named this element in the 19th century. This element was first isolated in 1774 by the Swiss-German chemist Carl Wilhelm Scheele, by reacting hydrochloric acid with manganese (IV) oxide. But he failed to realise his achievement, mistakenly believing it also contained oxygen. It was Davy in 1810 who finally concluded that Scheele had made elemental chlorine. Chlorine is in group 17 of periodic table, also called the halogens, and is not found as the element in nature - only as a compound. The most common of these being salt, or sodium chloride, and the potassium compounds sylvite (or potassium chloride) and carnallite (potassium magnesium chloride hexahydrate). It is also estimated that there are around two thousand organic chlorine compounds. Chlorine has two stable isotopes chlorine-35 and chlorine-37with Chlorine-35 accounting for roughly 3 out of every 4 naturally occurring chlorine atoms. Chlorine-36 is also known naturally and is a radioactive isotope with a half life of about 30,000 years. Chlorine has a major role to play in synthetic organic chemistry, taking part in three of the most common reaction mechanisms. In the first of these, the photochemical substitution reaction, chlorine reacts with an alkane by replacing one of the hydrogen atoms attached to a carbon forming a chloroalkane. This radical reaction is initiated by the use of sunlight or ultraviolet light to split diatomic chlorine into two radicals. Chlorine can also react with alkenes via the electrophilic addition mechanism. This time two chlorine atoms add to a molecule across the electron-rich carbon-carbon double bond. This reaction has to be carried out in the dark to avoid complications with competing free radical substitutions. A third common mechanism is electrophilic substitution, which occurs when chlorine reacts with a benzene ring by replacing a hydrogen atom forming chlorobenzene and hydrogen chloride. This reaction is most commonly known as the Friedal-Crafts reaction. Chlorine also has a multitude of industrial uses. Including making bulk materials like bleached paper products, plastics such as PVC and the solvents tetrachloromethane, chloroform and dichloromethane. It is also used to make dyes, textiles, medicines, antiseptics, insecticides and paints. It's best known uses however are probably in making bleaches such as 'Domestos' and in treating drinking and swimming pool waters to make them safe to use and of course its role as a chemical warfare agent. The treatment of water with chlorine began in London after a cholera outbreak in 1850 when the physician and pioneering hygienist John Snow identified a well in Soho as the source of the outbreak. Chlorine is still used in most sewage treatment works today. Snow also used a compound of chlorine - chloroform with the formula CHCl3 - as an anesthetic to aid the childbirth of two of Queen Victoria's children. The use of chlorine gas as a chemical weapon was pioneered by German chemist Fritz Haber, who is better known for his work with ammonia. It was first used against the Allied soldiers in the battle of Ypres during the first world war. While it was quickly replaced by the more deadly phosgene and mustard gases, chlorine gas has been used as a weapon as recently as 2007 in Iraq during the second gulf war. Chlorine was also once used to make a series of aerosol solvents and refrigerants called chlorofluorocarbons or CFCs. However their use was stopped once it became apparent that when in the atmosphere these compounds absorb ultraviolet light and cause homolytic bond fission producing a chlorine free radical which in turn reacts with ozone. This has led to a reduction in the concentration of ozone in the so-called ozone layer, and therefore a reduction in the protection for those of us on the surface of the planet making us more susceptible to skin cancers. So, that's chlorine - a Jeckll and Hyde element with an extremely wide range of applications. Chris Smith So slap on your sun screen. Tim Harrison was telling the tale of Element number 17, and that's chlorine. Tim's based at the University of Bristol's ChemLabs. Next week, the stuff that gives itself an x-ray. Brian Clegg This grey metallic element gives off beta particles as it decays. These can cause radioactive damage in their own right, but prometheum is probably most dangerous because those beta particles generate X-rays when they hit heavy nuclei, making a sample of promethium bathe its surroundings in a constant low dosage x-ray beam. It was initially used to replace radium in luminous dials. Promethium chloride was mixed with phosphors that glow yellowy-green or blue when radiation hits them. However, as the dangers of the element's radioactive properties became apparent, this too was dropped from the domestic glow-in-the-dark market, only employed now in specialist applications. Chris Smith And you can hear what some of those applications are when Brian Clegg looks at the story of promethium in next week's Chemistry in its Element. In the meantime more elements are available from the Chemistry in its Element podcast, that's on iTunes or on the web at chemistryworld.org/elements. I'm Chris Smith, thank you very much for listening and goodbye. (Promo) Chemistry in its element is brought to you by the Royal Society of Chemistry and produced by thenakedscientists.com. There's more information and other episodes of Chemistry in its element on our website at chemistryworld.org/elements. (End promo)
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Visual Elements images and videos
© Murray Robertson 1998-2017.
W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J. J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions (version 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.
John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.
Derived in part from material provided by the British Geological Survey © NERC.
Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.
Produced by The Naked Scientists.
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
© Murray Robertson 1998-2017.
Data
W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J. J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions (version 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.
Uses and properties
John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.
Supply risk data
Derived in part from material provided by the British Geological Survey © NERC.
History text
Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.
Podcasts
Produced by The Naked Scientists.
Periodic Table of Videos
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
