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The key question is why the use of gas chlorination as a treatment method has been targeted with such strong emotional arguments, usually not anchored in facts.
Use of chlorine gas, or gas chlorination, in a public water supply was first introduced early in the 20th century in Union City, NJ. The equipment was quite complicated, but it made an immediate difference in reducing the number of gastrointestinal outbreaks, and in eliminating most of the diseases mentioned above in a relatively short span of time. Through the 20th century, chlorination, and especially gas chlorination, became a common method of disinfecting water supplies and, subsequently, wastewater effluents.
Many of those early gas chlorination systems were massive pieces of equipment, prone to chlorine leaks, and highly susceptible to the corrosive effects of chlorine. By the end of the 20th century, deaths or serious injuries due to the use of gas chlorine in water and wastewater treatment plants, both municipal and industrial, were virtually non-existent.
Chlorination can be accomplished in a number of different forms. These processes are important to understand at a basic level if we are to fully understand the risks in each process and stick to the facts when assessing safety:.
Gas chlorine, while a hazardous material, is actually responsible for very few incidents in water and wastewater treatment facilities. In over 95 percent of all treatment facilities the amount of gas chlorine on hand is relatively small, by comparison to industrial use in non-water and wastewater treatment applications.
Pharmaceuticals and chemical manufacturing account for the vast majority of chlorine usage, and railroad tank cars are not an uncommon sight at those facilities. Another use that requires those same railroad tank cars of liquefied chlorine gas is sodium hypochlorite bleach manufacturing. Yes, there are incidents where operations personnel improperly handle the gas chlorine containers or accidentally remove a regulator or fitting without turning off the cylinder valve, as there are incidents with the handling of virtually every hazardous chemical known to mankind.
But, with chlorine gas, these are very isolated incidents, and, most importantly, no one dies. It can be very unpleasant, and preventive hospitalization may be necessary for observation, but in the end, personnel return to their jobs in virtually all cases. And, even more importantly, there has been no loss of life, mass evacuations, or mass hospitalizations as a result of chlorine gas leaks at water and wastewater treatment facilities in the USA in more than 20 years, where lb.
This is the vast majority of treatment facilities. Where one ton containers were in use, a very few incidents called for erring on the side of caution and people were evacuated from an area around the facility until the leakage was brought under control. But again, no loss of life. Nothing could be further from the truth! The other forms of chlorination, detailed above, have all had their significant accidents, and on a very regular basis.
Large fires and explosions have rocked calcium hypochlorite storage facilities, with the resultant toxic clouds carried over large areas. The media reports state that chlorine gas was released, but rarely, if ever, do they mention that the cause was anything other than chlorine gas containers.
A large water treatment facility in Michigan had used gas chlorination for many decades and had an enviable safety record. Concerns by city officials about the storage of gas chlorine containers led them to change the facility over to sodium hypochlorite. Very large storage tanks were installed in the facility for the bleach.
A tank truck delivering aluminum sulfate for part of the treatment process, accidentally pumped the chemical into one of the bleach tanks, and 12 workers went to the hospital with chemical inhalation distress. Countless cases of personnel accidentally spilling or adding the wrong chemical around bleach storage areas results in hundreds of injuries every year. One actually lifted the roof off of a storage building. And people have been injured.
A large sodium hypochlorite generator at a treatment plant in West Palm Beach, FL, exploded due to a failure of the hydrogen venting system, and peppered the walls of the chlorination room with plastic shrapnel. Luckily, this happened in the very early morning hours when no one was present in the room. Every treatment technology has some risks. Chlorine, in all of its forms, continues to offer us the most affordable and most well understood means of eliminating water borne diseases.
Both government as well as private researchers, in addition to health agencies and water treatment professionals, acknowledge that chlorination is a vital part of our water infrastructure and should remain so into the foreseeable future. Even the newest, limited application, treatment processes still require some chlorination to protect a public water distribution system from contamination after the water leaves the treatment plant.
Knowing all of this, there is still a movement to eliminate or reduce the use of gas chlorination. 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 than 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.
Beyond a mixture of HCl and H 2 O, the system separates completely into two separate liquid phases. 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: extreme inertness and reluctance to participate in chemical reactions the noble gases , with the exception of xenon in the highly unstable XeCl 2 and XeCl 4 ; extreme nuclear instability hampering chemical investigation before decay and transmutation many of the heaviest elements beyond bismuth ; and having an electronegativity higher than chlorine's oxygen and fluorine so that the resultant binary compounds are formally not chlorides but rather oxides or fluorides of chlorine.
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: for example, it will attack carbon monoxide to form carbonyl chlorofluoride, COFCl.
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. 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 , of which it is the anhydride. 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 stable towards dimerisation due to the delocalisation of the unpaired electron.
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: the analogous reaction with anhydrous hydrogen fluoride does not proceed to completion.
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.
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: it is known from the decomposition of aqueous chlorine dioxide.
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: chlorocarbons are typically denser than water due to the higher atomic weight of chlorine versus hydrogen, and aliphatic organochlorides are alkylating agents because chloride is a leaving group. Alkanes and aryl alkanes may be chlorinated under free-radical conditions, with UV light. However, the extent of chlorination is difficult to control: the reaction is not regioselective and often results in a mixture of various isomers with different degrees of chlorination, though this may be permissible if the products are easily separated.
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.Watch Tattooed Teen Fucking Car Gear Shift on Cam video on xHamster - the ultimate collection of free Home Made & Masturbation hardcore porn tube movies! Watch Tattooed Teen Fucking Car Gear Shift on Cam video on xHamster - the ultimate collection of free Home Made & .