CHLORINE
Chlorine is a chemical element with symbol Cl and atomic number 17. The second-lightest of the halogens, it appears between fluorine and bromine in the periodic table and its properties are mostly intermediate between them. Chlorine is a yellow-green gas at room temperature. It is an extremely reactive element and a strong oxidising agent: among the elements, it has the highest electron affinity and the third-highest electronegativity, behind only oxygen and fluorine.
The most common compound of chlorine, sodium chloride (common salt), has been known since ancient times. Around 1630, chlorine gas was first synthesised in a chemical reaction, but not recognised as a fundamentally important substance. Carl Wilhelm Scheelewrote a description of chlorine gas in 1774, supposing it to be an oxide of a new element. In 1809, chemists suggested that the gas might be a pure element, and this was confirmed by Sir Humphry Davy in 1810, who named it from Ancient Greek: χλωρός, translit. khlôros, lit. 'pale green' based on its colour.
PROPERTIES
Chlorine is the second halogen, being a nonmetal in group 17 of the periodic table. 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]3s23p5, 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. Corresponding to periodic trends, it is intermediate in electronegativity between fluorine and bromine (F: 3.98, Cl: 3.16, Br: 2.96, I: 2.66), and is less reactive than fluorine and more reactive than bromine. 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. It is intermediate in atomic radius between fluorine and bromine, and this leads to many of its atomic properties similarly continuing the trend from iodine to bromine upward, such as first ionisation energy, electron affinity, enthalpy of dissociation of the X2 molecule (X = Cl, Br, I), ionic radius, and X–X bond length. (Fluorine is anomalous due to its small size.)
Reactions with Water
Usually, reactions of chlorine with water are for disinfection purposes. Chlorine is only slightly soluble in water, with its maximum solubility occurring at 49° F. After that, its solubility decreases until 212° F. At temperatures below that range, it forms crystalline hydrates (usually Cl2Cl2) and becomes insoluble. Between that range, it usually forms hypochlorous acid (HOClHOCl). This is the primary reaction used for water/wastewater disinfection and bleaching.
Cl2 + H2O→HOCl + HCl
Cl2 + H2O→HOCl + HCl
At the boiling temperature of water, chlorine decomposes water
2Cl2 + 2H2O→4HCl + O2
2Cl2 + 2H2O→4HCl + O2
Reactions with Oxygen
Although chlorine usually has -1 oxidation state, it can have oxidation states of +1, +3, +4, or +7 in certain compounds, such as when it forms Oxoacids with the alkali metals
Reactions with Hydrogen
When H2 and Cl2 are exposed to sunlight or high temperatures, they react quickly and violently in a spontaneous reaction. Otherwise, the reaction proceeds slowly.
H2 + Cl2→2HCl
H2 + Cl2→2HCl
HCl can also be produced by reacting Chlorine with compounds containing Hydrogen, such as Hydrogen sulfide
Reactions with Halogens
Chlorine, like many of the other halogens, can form interhalogen compounds (examples include BrCl, ICl, ICl2). The heavier elements in one of these compounds acts as the central atom. For Chlorine, this occurs when it is bounded to fluorine in ClF, ClF3, and ClF5
Reactions with Metals
Chlorine reacts with most metals and forms metal chlorides, with most of these compounds being soluble in water. Examples of insoluble compounds include AgClAgCl and PbCl2PbCl2. Gaseous or liquid chlorine usually does not have an effect on metals such as iron, copper, platinum, silver, and steel at temperatures below 230°F. At high temperatures, however, it reacts rapidly with many of the metals, especially if the metal is in a form that has a high surface area (such as when powdered or made into wires).
EXAMPLE: OXIDIZING IRON
Chlorine can oxidizing iron
Cl2 + Fe→FeCl2
Cl2 + Fe→FeCl2
Half Reactions:
Fe→Fe+2 + 2e−
Fe→Fe+2 + 2e−
Cl2 + 2e−→2Cl−
ISOTOPES
Chlorine has two stable isotopes, 35Cl and 37Cl. These are its only two natural isotopes occurring in quantity, with 35Cl making up 76% of natural chlorine and 37Cl making up the remaining 24%. Both are synthesised in stars in the oxygen-burning and silicon-burning processes. Both have nuclear spin 3/2+ and thus may be used for nuclear magnetic resonance, although the spin magnitude being greater than 1/2 results in non-spherical nuclear charge distribution and thus resonance broadening as a result of a nonzero nuclear quadrupole moment and resultant quadrupolar relaxation. The other chlorine isotopes are all radioactive, with half-lives too short to occur in nature primordially. Of these, the most commonly used in the laboratory are 36Cl (t1/2 = 3.0×105 y) and 38Cl (t1/2 = 37.2 min), which may be produced from the neutron activation of natural chlorine.
The most stable chlorine radioisotope is 36Cl. The primary decay mode of isotopes lighter than 35Cl is electron capture to isotopes of sulfur; that of isotopes heavier than 37Cl is beta decay to isotopes of argon; and 36Cl may decay by either mode to stable 36S or 36Ar.
ORGANOCHLORINE COMPOUNDS
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.
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. The haloform reaction, using chlorine and sodium hydroxide, is also able to generate alkyl halides from methyl ketones, and related compounds. 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 (PCl5) or thionyl chloride (SOCl2). The last is very convenient in the laboratory because all side products are gaseous and do not have to be distilled out.
The most common compound of chlorine, sodium chloride (common salt), has been known since ancient times. Around 1630, chlorine gas was first synthesised in a chemical reaction, but not recognised as a fundamentally important substance. Carl Wilhelm Scheelewrote a description of chlorine gas in 1774, supposing it to be an oxide of a new element. In 1809, chemists suggested that the gas might be a pure element, and this was confirmed by Sir Humphry Davy in 1810, who named it from Ancient Greek: χλωρός, translit. khlôros, lit. 'pale green' based on its colour.
- Appearance: pale yellow-green gas
- Standard Atomic Weight: 35.446
- Atomic Number(Z): 17
- Group: group 17 (halogens)
- Period: period 3
- Block: p block
- Element Category: reactive non-metal
- PHASE AT STP : gas
- Melting Point : 171.6 K [-101.5 ° C]
- Boiling Point : 239.11 K [-34.04 ° C]
- Density(at STP) : 3.2 g/L
- Heat of Vaporization : 20.41 kj/mol
- Covalent Radius : 102+-4pm
- Crystal Structure : orthorhombic
- Discovery : Carl Wilhelm Scheele(1774)
PROPERTIES
Chlorine is the second halogen, being a nonmetal in group 17 of the periodic table. 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]3s23p5, 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. Corresponding to periodic trends, it is intermediate in electronegativity between fluorine and bromine (F: 3.98, Cl: 3.16, Br: 2.96, I: 2.66), and is less reactive than fluorine and more reactive than bromine. 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. It is intermediate in atomic radius between fluorine and bromine, and this leads to many of its atomic properties similarly continuing the trend from iodine to bromine upward, such as first ionisation energy, electron affinity, enthalpy of dissociation of the X2 molecule (X = Cl, Br, I), ionic radius, and X–X bond length. (Fluorine is anomalous due to its small size.)
Reactions with Water
Usually, reactions of chlorine with water are for disinfection purposes. Chlorine is only slightly soluble in water, with its maximum solubility occurring at 49° F. After that, its solubility decreases until 212° F. At temperatures below that range, it forms crystalline hydrates (usually Cl2Cl2) and becomes insoluble. Between that range, it usually forms hypochlorous acid (HOClHOCl). This is the primary reaction used for water/wastewater disinfection and bleaching.
Cl2 + H2O→HOCl + HCl
Cl2 + H2O→HOCl + HCl
At the boiling temperature of water, chlorine decomposes water
2Cl2 + 2H2O→4HCl + O2
2Cl2 + 2H2O→4HCl + O2
Reactions with Oxygen
Although chlorine usually has -1 oxidation state, it can have oxidation states of +1, +3, +4, or +7 in certain compounds, such as when it forms Oxoacids with the alkali metals
| Oxidation State | Compound |
| +1 | NaClO |
| +3 | NaClO2 |
| +5 | NaClO3 |
| +7 | NaClO4 |
Reactions with Hydrogen
When H2 and Cl2 are exposed to sunlight or high temperatures, they react quickly and violently in a spontaneous reaction. Otherwise, the reaction proceeds slowly.
H2 + Cl2→2HCl
H2 + Cl2→2HCl
HCl can also be produced by reacting Chlorine with compounds containing Hydrogen, such as Hydrogen sulfide
Reactions with Halogens
Chlorine, like many of the other halogens, can form interhalogen compounds (examples include BrCl, ICl, ICl2). The heavier elements in one of these compounds acts as the central atom. For Chlorine, this occurs when it is bounded to fluorine in ClF, ClF3, and ClF5
Reactions with Metals
Chlorine reacts with most metals and forms metal chlorides, with most of these compounds being soluble in water. Examples of insoluble compounds include AgClAgCl and PbCl2PbCl2. Gaseous or liquid chlorine usually does not have an effect on metals such as iron, copper, platinum, silver, and steel at temperatures below 230°F. At high temperatures, however, it reacts rapidly with many of the metals, especially if the metal is in a form that has a high surface area (such as when powdered or made into wires).
EXAMPLE: OXIDIZING IRON
Chlorine can oxidizing iron
Cl2 + Fe→FeCl2
Cl2 + Fe→FeCl2
Half Reactions:
Fe→Fe+2 + 2e−
Fe→Fe+2 + 2e−
Cl2 + 2e−→2Cl−
ISOTOPES
Chlorine has two stable isotopes, 35Cl and 37Cl. These are its only two natural isotopes occurring in quantity, with 35Cl making up 76% of natural chlorine and 37Cl making up the remaining 24%. Both are synthesised in stars in the oxygen-burning and silicon-burning processes. Both have nuclear spin 3/2+ and thus may be used for nuclear magnetic resonance, although the spin magnitude being greater than 1/2 results in non-spherical nuclear charge distribution and thus resonance broadening as a result of a nonzero nuclear quadrupole moment and resultant quadrupolar relaxation. The other chlorine isotopes are all radioactive, with half-lives too short to occur in nature primordially. Of these, the most commonly used in the laboratory are 36Cl (t1/2 = 3.0×105 y) and 38Cl (t1/2 = 37.2 min), which may be produced from the neutron activation of natural chlorine.
The most stable chlorine radioisotope is 36Cl. The primary decay mode of isotopes lighter than 35Cl is electron capture to isotopes of sulfur; that of isotopes heavier than 37Cl is beta decay to isotopes of argon; and 36Cl may decay by either mode to stable 36S or 36Ar.
ORGANOCHLORINE COMPOUNDS
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.
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. The haloform reaction, using chlorine and sodium hydroxide, is also able to generate alkyl halides from methyl ketones, and related compounds. 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 (PCl5) or thionyl chloride (SOCl2). The last is very convenient in the laboratory because all side products are gaseous and do not have to be distilled out.
