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Chlorous acid HOClO is even more unstable and cannot be isolated or concentrated without decomposition: Some types of organochlorides, though not all, have The Wizard of Oz – Fantastische Gewinnchancen in einer fantastischen Welt toxicity to plants or animals, including humans. This casino spiel gratis is conducted in the oxidising solvent arsenic pentafluoride. Dichlorine monoxide Cl 2 O living active deluxe 2 test a brownish-yellow gas red-brown when solid or liquid which may be obtained by reacting chlorine gas with yellow mercury II oxide. Wikimedia Commons has media related to Chlorine. Like the other carbon—halogen bonds, the C—Cl bond is a common functional group that forms part of core organic chemistry. It may be dissolved in water to regenerate perchloric acid or in aqueous alkalis to regenerate perchlorates. 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Its properties are thus similar to fluorine , bromine , and iodine , and are largely intermediate between those of the first two. Chlorine has the electron configuration [Ne]3s 2 3p 5 , with the seven electrons in the third and outermost shell acting as its valence electrons.
Like all halogens, it is thus one electron short of a full octet, and is hence a strong oxidising agent, reacting with many elements in order to complete its outer shell.
It is also a weaker oxidising agent than fluorine, but a stronger one than bromine. Conversely, the chloride ion is a weaker reducing agent than bromide, but a stronger one than fluoride.
Fluorine is anomalous due to its small size. All four stable halogens experience intermolecular van der Waals forces of attraction, and their strength increases together with the number of electrons among all homonuclear diatomic halogen molecules.
Thus, the melting and boiling points of chlorine are intermediate between those of fluorine and bromine: As a result of the increasing molecular weight of the halogens down the group, the density and heats of fusion and vaporisation of chlorine are again intermediate between those of bromine and fluorine, although all their heats of vaporisation are fairly low leading to high volatility thanks to their diatomic molecular structure.
This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group. Like solid bromine and iodine, solid chlorine crystallises in the orthorhombic crystal system , in a layered lattice of Cl 2 molecules.
This structure means that chlorine is a very poor conductor of electricity, and indeed its conductivity is so low as to be practically unmeasurable.
Chlorine has two stable isotopes, 35 Cl and 37 Cl. Both are synthesised in stars in the oxygen-burning and silicon-burning processes.
The other chlorine isotopes are all radioactive, with half-lives too short to occur in nature primordially.
The most stable chlorine radioisotope is 36 Cl. The primary decay mode of isotopes lighter than 35 Cl is electron capture to isotopes of sulfur ; that of isotopes heavier than 37 Cl is beta decay to isotopes of argon ; and 36 Cl may decay by either mode to stable 36 S or 36 Ar.
In the top meter of the lithosphere, 36 Cl is generated primarily by thermal neutron activation of 35 Cl and spallation of 39 K and 40 Ca. In the subsurface environment, muon capture by 40 Ca becomes more important as a way to generate 36 Cl.
Chlorine is intermediate in reactivity between fluorine and bromine, and is one of the most reactive elements.
Chlorine is a weaker oxidising agent than fluorine but a stronger one than bromine or iodine. However, this trend is not shown in the bond energies because fluorine is singular due to its small size, low polarisability, and lack of low-lying d-orbitals available for bonding which chlorine has.
As another difference, chlorine has a significant chemistry in positive oxidation states while fluorine does not. Chlorination often leads to higher oxidation states than bromination or iodination but lower oxidation states to fluorination.
However, the kinetics of this reaction are unfavorable, and there is also a bubble overpotential effect to consider, so that electrolysis of aqueous chloride solutions evolves chlorine gas and not oxygen gas, a fact that is very useful for the industrial production of chlorine.
The simplest chlorine compound is hydrogen chloride , HCl, a major chemical in industry as well as in the laboratory, both as a gas and dissolved in water as hydrochloric acid.
It is often produced by burning hydrogen gas in chlorine gas, or as a byproduct of chlorinating hydrocarbons. Another approach is to treat sodium chloride with concentrated sulfuric acid to produce hydrochloric acid, also known as the "salt-cake" process: In the laboratory, hydrogen chloride gas may be made by drying the acid with concentrated sulfuric acid.
Deuterium chloride , DCl, may be produced by reacting benzoyl chloride with heavy water D 2 O. At room temperature, hydrogen chloride is a colourless gas, like all the hydrogen halides apart from hydrogen fluoride , since hydrogen cannot form strong hydrogen bonds to the larger electronegative chlorine atom; however, weak hydrogen bonding is present in solid crystalline hydrogen chloride at low temperatures, similar to the hydrogen fluoride structure, before disorder begins to prevail as the temperature is raised.
Hydrochloric acid forms an azeotrope with boiling point Anhydrous hydrogen chloride is a poor solvent, only able to dissolve small molecular compounds such as nitrosyl chloride and phenol , or salts with very low lattice energies such as tetraalkylammonium halides.
Solvolysis , ligand replacement reactions, and oxidations are well-characterised in hydrogen chloride solution: Nearly all elements in the periodic table form binary chlorides.
The exceptions are decidedly in the minority and stem in each case from one of three causes: Chlorination of metals with Cl 2 usually leads to a higher oxidation state than bromination with Br 2 when multiple oxidation states are available, such as in MoCl 5 and MoBr 3.
Chlorides can be made by reaction of an element or its oxide, hydroxide, or carbonate with hydrochloric acid, and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen chloride gas.
These methods work best when the chloride product is stable to hydrolysis; otherwise, the possibilities include high-temperature oxidative chlorination of the element with chlorine or hydrogen chloride, high-temperature chlorination of a metal oxide or other halide by chlorine, a volatile metal chloride, carbon tetrachloride , or an organic chloride.
For instance, zirconium dioxide reacts with chlorine at standard conditions to produce zirconium tetrachloride , and uranium trioxide reacts with hexachloropropene when heated under reflux to give uranium tetrachloride.
The second example also involves a reduction in oxidation state , which can also be achieved by reducing a higher chloride using hydrogen or a metal as a reducing agent.
This may also be achieved by thermal decomposition or disproportionation as follows: Silver chloride is very insoluble in water and is thus often used as a qualitative test for chlorine.
This is very unstable and has only been characterised by its electronic band spectrum when produced in a low-pressure discharge tube.
This reaction is conducted in the oxidising solvent arsenic pentafluoride. The three fluorides of chlorine form a subset of the interhalogen compounds, all of which are diamagnetic.
Chlorine monofluoride ClF is extremely thermally stable, and is sold commercially in gram steel lecture bottles. Its properties are mostly intermediate between those of chlorine and fluorine.
It will react with many metals and nonmetals from room temperature and above, fluorinating them and liberating chlorine.
It will also act as a chlorofluorinating agent, adding chlorine and fluorine across a multiple bond or by oxidation: It will also react exothermically and violently with compounds containing —OH and —NH groups, such as water: It is one of the most reactive known chemical compounds, reacting with many substances which in ordinary circumstances would be considered chemically inert, such as asbestos , concrete, and sand.
It explodes on contact with water and most organic substances. The list of elements it sets on fire is diverse, containing hydrogen , potassium , phosphorus , arsenic , antimony , sulfur , selenium , tellurium , bromine , iodine , and powdered molybdenum , tungsten , rhodium , iridium , and iron.
An impermeable fluoride layer is formed by sodium , magnesium , aluminium , zinc , tin , and silver , which may be removed by heating. When heated, even such noble metals as palladium , platinum , and gold are attacked and even the noble gases xenon and radon do not escape fluorination.
Nickel containers are usually used due to that metal's great resistance to attack by chlorine trifluoride, stemming from the formation of an unreactive nickel fluoride layer.
Its reaction with hydrazine to form hydrogen fluoride, nitrogen, and chlorine gases was used in experimental rocket motors, but has problems largely stemming from its extreme hypergolicity resulting in ignition without any measurable delay.
For these reasons, it was used in bomb attacks during the Second World War by the Nazis. Today, it is mostly used in nuclear fuel processing, to oxidise uranium to uranium hexafluoride for its enriching and to separate it from plutonium.
It is a very strong fluorinating agent, although it is still not as effective as chlorine trifluoride. Only a few specific stoichiometric reactions have been characterised.
The product, chloryl fluoride , is one of the five known chlorine oxide fluorides. All five behave similarly to the chlorine fluorides, both structurally and chemically, and may act as Lewis acids or bases by gaining or losing fluoride ions respectively or as very strong oxidising and fluorinating agents.
The chlorine oxides are well-studied in spite of their instability all of them are endothermic compounds. They are important because they are produced when chlorofluorocarbons undergo photolysis in the upper atmosphere and cause the destruction of the ozone layer.
None of them can be made from directly reacting the elements. Dichlorine monoxide Cl 2 O is a brownish-yellow gas red-brown when solid or liquid which may be obtained by reacting chlorine gas with yellow mercury II oxide.
It is very soluble in water, in which it is in equilibrium with hypochlorous acid HOCl , which it is the anhydride of. It is thus an effective bleach and is mostly used to make hypochlorites.
It explodes on heating or sparking or in the presence of ammonia gas. Chlorine dioxide ClO 2 was the first chlorine oxide to be discovered in by Humphry Davy.
It is a yellow paramagnetic gas deep-red as a solid or liquid , as expected from its having an odd number of electrons: It is usually prepared by reducing a chlorate as follows: Its production is thus intimately linked to the redox reactions of the chlorine oxoacids.
It is a strong oxidising agent, reacting with sulfur , phosphorus , phosphorus halides, and potassium borohydride. It dissolves exothermically in water to form dark-green solutions that very slowly decompose in the dark.
However, in the presence of light, these solutions rapidly photodecompose to form a mixture of chloric and hydrochloric acids. The ClO radical leads to the depletion of atmospheric ozone and is thus environmentally important as follows: Chlorine perchlorate ClOClO 3 is a pale yellow liquid that is less stable than ClO 2 and decomposes at room temperature to form chlorine, oxygen, and dichlorine hexoxide Cl 2 O 6.
It hydrolyses in water to give a mixture of chloric and perchloric acids: It is a shock-sensitive, colourless oily liquid. It is the least reactive of the chlorine oxides, being the only one to not set organic materials on fire at room temperature.
It may be dissolved in water to regenerate perchloric acid or in aqueous alkalis to regenerate perchlorates. However, it thermally decomposes explosively by breaking one of the central Cl—O bonds, producing the radicals ClO 3 and ClO 4 which immediately decompose to the elements through intermediate oxides.
Chlorine forms four oxoacids: As can be seen from the redox potentials given in the adjacent table, chlorine is much more stable towards disproportionation in acidic solutions than in alkaline solutions: The rates of reaction for the chlorine oxyanions increases as the oxidation state of chlorine decreases.
The strengths of the chlorine oxyacids increase very quickly as the oxidation state of chlorine increases due to the increasing delocalisation of charge over more and more oxygen atoms in their conjugate bases.
Most of the chlorine oxoacids may be produced by exploiting these disproportionation reactions. Hypochlorous acid HOCl is highly reactive and quite unstable; its salts are mostly used for their bleaching and sterilising abilities.
They are very strong oxidising agents, transferring an oxygen atom to most inorganic species. Chlorous acid HOClO is even more unstable and cannot be isolated or concentrated without decomposition: However, sodium chlorite is a stable salt and is useful for bleaching and stripping textiles, as an oxidising agent, and as a source of chlorine dioxide.
Its most important salt is sodium chlorate , mostly used to make chlorine dioxide to bleach paper pulp.
The decomposition of chlorate to chloride and oxygen is a common way to produce oxygen in the laboratory on a small scale.
Chloride and chlorate may comproportionate to form chlorine as follows: Perchloric acid and aqueous perchlorates are vigorous and sometimes violent oxidising agents when heated, in stark contrast to their mostly inactive nature at room temperature due to the high activation energies for these reactions for kinetic reasons.
Perchlorates are made by electrolytically oxidising sodium chlorate, and perchloric acid is made by reacting anhydrous sodium perchlorate or barium perchlorate with concentrated hydrochloric acid, filtering away the chloride precipitated and distilling the filtrate to concentrate it.
Anhydrous perchloric acid is a colourless mobile liquid that is sensitive to shock that explodes on contact with most organic compounds, sets hydrogen iodide and thionyl chloride on fire and even oxidises silver and gold.
Like the other carbon—halogen bonds, the C—Cl bond is a common functional group that forms part of core organic chemistry. Formally, compounds with this functional group may be considered organic derivatives of the chloride anion.
Due to the difference of electronegativity between chlorine 3. Chlorination modifies the physical properties of hydrocarbons in several ways: Alkanes and aryl alkanes may be chlorinated under free radical conditions, with UV light.
However, the extent of chlorination is difficult to control: Aryl chlorides may be prepared by the Friedel-Crafts halogenation , using chlorine and a Lewis acid catalyst.
Chlorine adds to the multiple bonds on alkenes and alkynes as well, giving di- or tetra-chloro compounds. However, due to the expense and reactivity of chlorine, organochlorine compounds are more commonly produced by using hydrogen chloride, or with chlorinating agents such as phosphorus pentachloride PCl 5 or thionyl chloride SOCl 2.
The last is very convenient in the laboratory because all side products are gaseous and do not have to be distilled out. Many organochlorine compounds have been isolated from natural sources ranging from bacteria to humans.
Some types of organochlorides, though not all, have significant toxicity to plants or animals, including humans. Dioxins, produced when organic matter is burned in the presence of chlorine, and some insecticides, such as DDT , are persistent organic pollutants which pose dangers when they are released into the environment.
For example, DDT, which was widely used to control insects in the mid 20th century, also accumulates in food chains, and causes reproductive problems e.
Chlorine is too reactive to occur as the free element in nature but is very abundant in the form of its chloride salts. All of these pale in comparison to the reserves of chloride ions in seawater: Small batches of chlorine gas are prepared in the laboratory by combining hydrochloric acid and manganese dioxide , but the need rarely arises due to its ready availability.
In industry, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water.
This method, the chloralkali process industrialized in , now provides most industrial chlorine gas.
The process proceeds according to the following chemical equation: In diaphragm cell electrolysis, an asbestos or polymer-fiber diaphragm separates a cathode and an anode , preventing the chlorine forming at the anode from re-mixing with the sodium hydroxide and the hydrogen formed at the cathode.
Diaphragm methods produce dilute and slightly impure alkali, but they are not burdened with the problem of mercury disposal and they are more energy efficient.
Membrane cell electrolysis employs permeable membrane as an ion exchanger. Saturated sodium or potassium chloride solution is passed through the anode compartment, leaving at a lower concentration.
This method also produces very pure sodium or potassium hydroxide but has the disadvantage of requiring very pure brine at high concentrations.
In the Deacon process , hydrogen chloride recovered from the production of organochlorine compounds is recovered as chlorine.
The process relies on oxidation using oxygen:. The reaction requires a catalyst. As introduced by Deacon, early catalysts were based on copper.
Commercial processes, such as the Mitsui MT-Chlorine Process, have switched to chromium and ruthenium-based catalysts. Sodium chloride is by a huge margin the most common chlorine compound, and it is the main source of chlorine and hydrochloric acid for the enormous chlorine-chemicals industry today.
About chlorine-containing compounds are commercially traded, including such diverse compounds as chlorinated methanes and ethanes , vinyl chloride and its polymer polyvinyl chloride PVC , aluminium trichloride for catalysis , the chlorides of magnesium , titanium , zirconium , and hafnium which are the precursors for producing the pure elements, and so on.
Other particularly important organochlorines are methyl chloride , methylene chloride , chloroform , vinylidene chloride , trichloroethylene , perchloroethylene , allyl chloride , epichlorohydrin , chlorobenzene , dichlorobenzenes , and trichlorobenzenes.
In France as elsewhere , animal intestines were processed to make musical instrument strings, Goldbeater's skin and other products.
This was done in "gut factories" boyauderies , and it was an odiferous and unhealthy process. Labarraque's research resulted in the use of chlorides and hypochlorites of lime calcium hypochlorite and of sodium sodium hypochlorite in the boyauderies.
The same chemicals were found to be useful in the routine disinfection and deodorization of latrines , sewers , markets, abattoirs , anatomical theatres , and morgues.
Labarraque's chlorinated lime and soda solutions have been advocated since to prevent infection called "contagious infection", presumed to be transmitted by " miasmas " , and to treat putrefaction of existing wounds, including septic wounds.
In , the contagion of infections was well known, even though the agency of the microbe was not discovered until more than half a century later. During the Paris cholera outbreak of , large quantities of so-called chloride of lime were used to disinfect the capital.
This was not simply modern calcium chloride , but chlorine gas dissolved in lime-water dilute calcium hydroxide to form calcium hypochlorite chlorinated lime.
Labarraque's discovery helped to remove the terrible stench of decay from hospitals and dissecting rooms, and by doing so, effectively deodorised the Latin Quarter of Paris.
Chloride of lime was used for destroying odors and "putrid matter". One source claims chloride of lime was used by Dr.
John Snow to disinfect water from the cholera-contaminated well that was feeding the Broad Street pump in London,  though three other reputable sources that describe that famous cholera epidemic do not mention the incident.
Perhaps the most famous application of Labarraque's chlorine and chemical base solutions was in , when Ignaz Semmelweis used chlorine-water chlorine dissolved in pure water, which was cheaper than chlorinated lime solutions to disinfect the hands of Austrian doctors, which Semmelweis noticed still carried the stench of decomposition from the dissection rooms to the patient examination rooms.
Long before the germ theory of disease, Semmelweis theorized that "cadaveric particles" were transmitting decay from fresh medical cadavers to living patients, and he used the well-known "Labarraque's solutions" as the only known method to remove the smell of decay and tissue decomposition which he found that soap did not.
The solutions proved to be far more effective antiseptics than soap Semmelweis was also aware of their greater efficacy, but not the reason , and this resulted in Semmelweis's celebrated success in stopping the transmission of childbed fever "puerperal fever" in the maternity wards of Vienna General Hospital in Austria in Much later, during World War I in , a standardized and diluted modification of Labarraque's solution containing hypochlorite 0.
Called Dakin's solution , the method of wound irrigation with chlorinated solutions allowed antiseptic treatment of a wide variety of open wounds, long before the modern antibiotic era.
A modified version of this solution continues to be employed in wound irrigation in modern times, where it remains effective against bacteria that are resistant to multiple antibiotics see Century Pharmaceuticals.
By , the US Department of Treasury called for all drinking water to be disinfected with chlorine.
Chlorine is presently an important chemical for water purification such as in water treatment plants , in disinfectants , and in bleach.
As a disinfectant in water, chlorine is more than three times as effective against Escherichia coli as bromine , and more than six times as effective as iodine.
Chlorine is usually used in the form of hypochlorous acid to kill bacteria and other microbes in drinking water supplies and public swimming pools.
In most private swimming pools, chlorine itself is not used, but rather sodium hypochlorite , formed from chlorine and sodium hydroxide , or solid tablets of chlorinated isocyanurates.
The drawback of using chlorine in swimming pools is that the chlorine reacts with the proteins in human hair and skin. The distinctive 'chlorine aroma' associated with swimming pools is not the result of chlorine itself, but of chloramine , a chemical compound produced by the reaction of free dissolved chlorine with amines in organic substances.
Even small water supplies are now routinely chlorinated. It is often impractical to store and use poisonous chlorine gas for water treatment, so alternative methods of adding chlorine are used.
These include hypochlorite solutions, which gradually release chlorine into the water, and compounds like sodium dichloro-s-triazinetrione dihydrate or anhydrous , sometimes referred to as "dichlor", and trichloro-s-triazinetrione , sometimes referred to as "trichlor".
These compounds are stable while solid and may be used in powdered, granular, or tablet form. When added in small amounts to pool water or industrial water systems, the chlorine atoms hydrolyze from the rest of the molecule forming hypochlorous acid HOCl , which acts as a general biocide, killing germs, micro-organisms, algae, and so on.
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