Liquid Chlorine Pool Shock - Commercial Grade 12.5% Concentrated Strength - 1 Gallon

Product Information

Dichlorine monoxide (Cl 2O) 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. [53] It’s also a good product to use for a daily treatment during times of heavy use. How to use liquid chlorine for your pool

Gédéon, Andras (2006). Science and technology in medicine. Springer. pp.181–82. ISBN 978-0-387-27874-2. Archived from the original on 2015-12-31. 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. [68] The chlorine produced is available in cylinders from sizes ranging from 450g to 70kg, as well as drums (865kg), tank wagons (15tonnes on roads; 27–90tonnes by rail), and barges (600–1200tonnes). [69] Applications Snelders, H. A. M. (1971). "J. S. C. Schweigger: His Romanticism and His Crystal Electrical Theory of Matter". Isis. 62 (3): 328–38. doi: 10.1086/350763. JSTOR 229946. S2CID 170337569.

How Does an Automatic Pool Chlorinator Work?

The chlorine to ammonia nitrogen ratio characterizes what kind of residual is produced. Are there Other Uses for Chlorine? Chlorination can be done at any time/point throughout the water treatment process - there is not one specific time when chlorine must be added. Each point of chlorine application will subsequently control a different water contaminant concern, thus offering a complete spectrum of treatment from the time the water enters the treatment facility to the time it leaves. a b "Chlorine: Product Datasheet" (PDF). Bayer MaterialScience AG. 2008-04-21. Archived from the original (PDF) on September 15, 2012 . Retrieved 2013-12-17. NOAA Office of Response and Restoration, US GOV. "Chlorine". noaa.gov. Archived from the original on 15 October 2015 . Retrieved 25 August 2015.

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 (3.16) and carbon (2.55), the carbon in a C–Cl bond is electron-deficient and thus electrophilic. 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. [56] 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: chlorine melts at −101.0°C and boils at −34.0°C. 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. [35] The halogens darken in colour as the group is descended: thus, while fluorine is a pale yellow gas, chlorine is distinctly yellow-green. This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group. [35] Specifically, the colour of a halogen, such as chlorine, results from the electron transition between the highest occupied antibonding π g molecular orbital and the lowest vacant antibonding σ u molecular orbital. [36] The colour fades at low temperatures, so that solid chlorine at −195°C is almost colourless. [35] Chlorine gas was first used as a weapon on April 22, 1915 at the Second Battle of Ypres by the German Army. [31] [32] The effect on the allies was devastating because the existing gas masks were difficult to deploy and had not been broadly distributed. [33] [34] Properties Chlorine, liquefied under a pressure of 7.4 bar at room temperature, displayed in a quartz ampule embedded in acrylic glass Solid chlorine at −150°CChlorine is detectable with measuring devices in concentrations as low as 0.2 parts per million (ppm), and by smell at 3ppm. Coughing and vomiting may occur at 30ppm and lung damage at 60ppm. About 1000ppm can be fatal after a few deep breaths of the gas. [14] The IDLH (immediately dangerous to life and health) concentration is 10ppm. [113] Breathing lower concentrations can aggravate the respiratory system and exposure to the gas can irritate the eyes. [114] When chlorine is inhaled at concentrations greater than 30ppm, it reacts with water within the lungs, producing hydrochloric acid (HCl) and hypochlorous acid (HOCl). The membrane cell process". Euro Chlor. Archived from the original on 2011-11-11 . Retrieved 2007-08-15. Audi, Georges; Bersillon, Olivier; Blachot, Jean; Wapstra, Aaldert Hendrik (2003), "The N UBASE evaluation of nuclear and decay properties", Nuclear Physics A, 729: 3–128, Bibcode: 2003NuPhA.729....3A, doi: 10.1016/j.nuclphysa.2003.11.001 Scheele observed several of the properties of chlorine: the bleaching effect on litmus, the deadly effect on insects, the yellow-green color, and the smell similar to aqua regia. [14] He called it " dephlogisticated muriatic acid air" since it is a gas (then called "airs") and it came from hydrochloric acid (then known as "muriatic acid"). [13] He failed to establish chlorine as an element. [13] Chlorine gas was first used by French chemist Claude Berthollet to bleach textiles in 1785. [25] [26] Modern bleaches resulted from further work by Berthollet, who first produced sodium hypochlorite in 1789 in his laboratory in the town of Javel (now part of Paris, France), by passing chlorine gas through a solution of sodium carbonate. The resulting liquid, known as " Eau de Javel" (" Javel water"), was a weak solution of sodium hypochlorite. This process was not very efficient, and alternative production methods were sought. Scottish chemist and industrialist Charles Tennant first produced a solution of calcium hypochlorite ("chlorinated lime"), then solid calcium hypochlorite (bleaching powder). [25] These compounds produced low levels of elemental chlorine and could be more efficiently transported than sodium hypochlorite, which remained as dilute solutions because when purified to eliminate water, it became a dangerously powerful and unstable oxidizer. Near the end of the nineteenth century, E. S. Smith patented a method of sodium hypochlorite production involving electrolysis of brine to produce sodium hydroxide and chlorine gas, which then mixed to form sodium hypochlorite. [27] This is known as the chloralkali process, first introduced on an industrial scale in 1892, and now the source of most elemental chlorine and sodium hydroxide. [28] In 1884 Chemischen Fabrik Griesheim of Germany developed another chloralkali process which entered commercial production in 1888. [29]

Now, just follow the manufacturer’s instructions for how much to add to bring it up to your desired level. If you’re still a little confused, you can use an online chlorine calculator to help you with the math. Liquid chlorine vs powder chlorine Do not enter the pool until the chlorine levels are below 3 ppm! How much liquid chlorine should you add? During the Paris cholera outbreak of 1832, 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. [77] These "putrid miasmas" were thought by many to cause the spread of "contagion" and "infection" – both words used before the germ theory of infection. 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 1854 London, [78] though three other reputable sources that describe that famous cholera epidemic do not mention the incident. [79] [80] [81] One reference makes it clear that chloride of lime was used to disinfect the offal and filth in the streets surrounding the Broad Street pump – a common practice in mid-nineteenth century England. [79] :296 Semmelweis and experiments with antisepsis Ignaz SemmelweisHelmont, Johannes (Joan) Baptista Van, Encyclopedia.Com: "Others were chlorine gas from the reaction of nitric acid and sal ammoniac; … "

Joseph Cotruvo, Victor Kimm, Arden Calvert. "Drinking Water: A Half Century of Progress." EPA Alumni Association. March 1, 2016. UV rays will eat up your chlorine like yesterday’s lunch if you don’t have anything in the pool to protect it—this means you’re constantly having to add more chlorine, costing time and money.

Weaponry: Use of Chlorine Gas Cylinders in World War I". historynet.com. 2006-06-12. Archived from the original on 2008-07-02 . Retrieved 2008-07-10. The hypochlorite ions also disproportionate further to produce chloride and chlorate (3 ClO − ⇌ 2 Cl − + ClO − 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. Chloric acid (HOClO 2) is a strong acid that is quite stable in cold water up to 30% concentration, but on warming gives chlorine and chlorine dioxide. Evaporation under reduced pressure allows it to be concentrated further to about 40%, but then it decomposes to perchloric acid, chlorine, oxygen, water, and 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: [55] ClO − Lewis, Kenneth A. (2010). "Ch. 9 Hypochlorination – Sodium Hypochlorite" (PDF). White's Handbook of Chlorination and Alternative Disinfectants. Hoboken, NJ: Wiley. p.452. doi: 10.1002/9780470561331.ch9. ISBN 978-0-470-56133-1. Archived (PDF) from the original on 2022-10-09. [ permanent dead link] Elemental chlorine solutions dissolved in chemically basic water (sodium and calcium hypochlorite) were first used as anti- putrefaction agents and disinfectants in the 1820s, in France, long before the establishment of the germ theory of disease. This practice was pioneered by Antoine-Germain Labarraque, who adapted Berthollet's "Javel water" bleach and other chlorine preparations. [30] Elemental chlorine has since served a continuous function in topical antisepsis (wound irrigation solutions and the like) and public sanitation, particularly in swimming and drinking water. [14]