Chlorine
 Chlorine gas in an ampoule. | |||||
| General properties | |||||
|---|---|---|---|---|---|
| Name, symbol | Chlorine, Cl | ||||
| Appearance | Greenish gas | ||||
| Chlorine in the periodic table | |||||
| |||||
| Atomic number | 17 | ||||
| Standard atomic weight (Ar) | 35.45 | ||||
| Group, block | , p-block | ||||
| Period | period 17 | ||||
| Electron configuration | [Ne] 3s2 3p5 | ||||
per shell | 2, 8, 7 | ||||
| Physical properties | |||||
| Pale green | |||||
| Phase | Gas | ||||
| Melting point | 171.6 K (−101.5 °C, −150.7 °F) | ||||
| Boiling point | 239.11 K (−34.04 °C, −29.27 °F) | ||||
| Density at (0 °C and 101.325 kPa) | 3.2 g/L | ||||
| when liquid, at | 1.5625 g/cm3 | ||||
| Critical point | 416.9 K, 7.991 MPa | ||||
| Heat of fusion | 6.406 kJ/mol | ||||
| Heat of | 20.41 kJ/mol | ||||
| Molar heat capacity | 33.949 J/(mol·K) | ||||
| pressure | |||||
| Atomic properties | |||||
| Oxidation states | 7, 6, 5, 4, 3, 2, 1, −1 (a strongly acidic oxide) | ||||
| Electronegativity | Pauling scale: 3.16 | ||||
| energies |
1st: 1251.2 kJ/mol 2nd: 2298 kJ/mol 3rd: 3822 kJ/mol | ||||
| Covalent radius | 102±4 pm | ||||
| Van der Waals radius | 175 pm | ||||
| Miscellanea | |||||
| Crystal structure | Orthorhombic | ||||
| Thermal conductivity | 8.9×10−3 W/(m·K) | ||||
| Electrical resistivity | >10 Ω·m (at 20 °C) | ||||
| Magnetic ordering | Diamagnetic | ||||
| CAS Registry Number | 7782-50-5 | ||||
| Discovery and first isolation | Carl Wilhelm Scheele (1774) | ||||
Chlorine is the second-lightest halogen, with the symbol Cl and atomic number 17. It has a sickly green colur and a distinctive smell, recognizable to many at low concentrations as 'the smell of pool centers' due to its use as a water disinfecting agent.
Contents
Properties
Chemical
Chlorine is a strong oxidizer with 7 valence electrons. Its unstable electron configuration results in high reactivity. Because of this, chlorine usually exists on earth in the form of a halide salt, and free chlorine is rare. Like fluorine, elemental chlorine forms a highly reactive diatomic gas.[1] Chlorine, like other halogens, forms many oxoanions, negatively charged ions containing oxygen. Most notably, these are hypochlorite(ClO-), chlorite(ClO2-), chlorate(ClO3-), and perchlorate(ClO4-).
Unlike hydrochloric acid, elemental chlorine easily corrodes copper, especially in moist air.[2]
Physical
Chlorine is a yellow-greenish gas, with a powerful odor similar to that of boiling hypochlorite solutions. It is heavier than air, and slightly soluble in water, 3.26 g/L.
Availability
While liquid (as in liquefied, and not aqueous solution) chlorine is sold by gas companies, it is hard to get hold of as it's very toxic and corrosive.
Chlorine is better produced from OTC products.
Production
There are many methods to generating chlorine gas, due to it being such a commonly used ion.
Potassium permanganate will oxidize hydrochloric acid to chlorine gas, as will manganese dioxide. In both cases, manganese(II) chloride will be produced as well.
- 2 KMnO4 + 16 HCl → 2 KCl + 2 MnCl2 + 5 Cl2 + 8 H2O
A hypochlorite and hydrochloric acid will produce chlorine; either a solution of sodium hypochlorite or calcium hypochlorite. A violent reaction with a lot of foam may take place in the case of the latter, and starting small scale is a must to get a sense for the reaction before any large scale chlorine production is attempted.
- MOCl + HCl → MCl + Cl2 + H2O
A popular way of making chlorine on Sciencemadness is using hydrochloric acid and trichloroisocyanuric acid (TCCA). TCCA can be found as slow release chlorine tablets for swimming pools. This reaction is favorable because it not too expensive, produces a large amount of chlorine over an extended period of time (while hypochlorites tend to violently produce all the chlorine right on mixing with the acid), leaves no awful byproducts (such as MnO2) and the reaction speed at standard concentrations and temperatures is not too fast nor too slow for most applications.
Projects/Experiments
Chlorine can be used to produce anhydrous metal chlorides, such as aluminum(III) chloride, iron(II) or iron(III) chloride, and many others, which cannot be made in solution, due to formation of hydrates which are irreversible, and decompose to the metal oxide and HCl gas upon heating.
- 2 Al + 3 Cl2 → 2 AlCl3
- 2 Fe + 3 Cl2 → 2 FeCl3
Other projects:
- Make hypochlorites
- Make sulfur dichloride and disulfur dichloride
- Make interhalogen compounds
- Alkane halogenation
- Element collecting
Handling
Safety
Elemental chlorine is extremely toxic and corrosive to mot materials, be it organic or inorganic. Because it is heavier than air, it tends to accumulate at the bottom of poorly ventilated spaces. It is highly irritant to eyes, mucous membrane, throat and lungs, even short term exposure may cause injury.
Chlorine gas is a strong oxidizer, which may react with flammable materials. It is notorious for reacting with iron at high temperatures, in a strong exothermic reaction, known as chlorine-iron fire.
Storage
Liquefied chlorine must be stored in cold places, away from any source of heat. Chlorine can be liquefied at room temperature, at a pressure of 7.4 bar.
Chlorine releasing chemicals, such as bleach and TCCA should be stored in closed bottles, usually covered with a bag or in a box, that must be opened form time to time to release the pressure.
The storage area for both chemicals should not contain any metal parts susceptible to chlorine attack.
Disposal
As it is toxic and has an irritating smell, it is recommended to neutralize chlorine before disposing of it. In gaseous form, elemental chlorine can be neutralized with ammonia at low concentrations, reaction that produces nitrogen gas and ammonium chloride, though chloramines may also be produced as side products. This should not be used at high concentrations of chlorine in air, where chloramines may be produced in higher amount.
Aqueous chlorine however, should never be neutralized with ammonia, as it will generate toxic chloramines. Acids should also be avoided. Hydrogen peroxide will neutralize bleach and release oxygen. Ascorbic acid and its salts are also good at neutralizing chlorine. Other good neutralizing agents are certain sulfur compounds, such as sulfites, bisulfites, metabisulfites, thiosulfites.
Gallery
A sample of liquid chlorine
References
- ↑ "Handbook of Toxicology of Chemical Warfare Agents", Academic Press, 2009, p. 313.
- ↑ Handbook of Corrosion Data, by Bruce D. Craig, David S. Anderson, p. 271
Relevant Sciencemadness threads
- Chlorine - Illustrated Practical Guide
- Best overall chlorine generator?
- Making Chlorine
- Preparation of Chlorine
- Small scale chlorine generation.
- Passing chlorine gas
- Found an easy way to make chlorine gas from muriatic acid and bleach
- Ways to get pure chlorine gas in a container/separate from the generator
- What to do with chlorine gas ?
- Storing Chlorine as a liquid?
- Chlorine gas neutralization
