HALOGEN FAMILY

Fluorine is a chemical element with symbol F and atomic number 9. It is the lightest halogen and exists as a highly toxic pale yellow diatomic gas at standard conditions. As the most electronegative element, it is extremely reactive, as it reacts with almost all other elements, except for helium and neon.
Among the elements, fluorine ranks 24th in universal abundance and 13th in terrestrial abundance. Fluorite, the primary mineral source of fluorine which gave the element its name, was first described in 1529; as it was added to metal ores to lower their melting points for smelting, the Latin verb fluo meaning "flow" gave the mineral its name.

• Allotropes:    alpha, beta
• Appearance: GAS(very pale yellow), LIQUID(bright yellow), SOLID(alpha is opaque & beta is transparent)
• Standard Atomic Weight: 18.998
• Atomic Number(Z): 9
• Group: group 17 (halogens)
• Period: period 2
• Block: p block
• Element Category: reactive non-metal
• PHASE AT STP : gas
• Melting Point : 53.48 K [-219.67 ° C]
• Boiling Point : 85.03 K [-188.11 ° C]
• Density(at STP) : 1.696 g/L
• Heat of Vaporization : 6.51 kj/mol
• Covalent Radius : 64pm
• Crystal Structure : cubic crystal structure for fluorine
• Discovery :  Andre-Marie Ampere(1810)

               
                            Reactions of Fluorine
Because of its reactivity, elemental fluorine is never found in nature and no other chemical element can displace fluorine from its compounds. Fluorine bonds with almost any element, both metals and nonmetals, because it is a very strong oxidizing agent. It is very unstable and reactive since it is so close to its ideal electron configuration. It forms covalent bonds with nonmetals, and since it is the most electronegative element, is always going to be the element that is reduced. It can also form a diatomic element with itself (F2F2), or covalent bonds where it oxidizes other halogens (ClFClF, ClF3ClF3, ClF5ClF5). It will react explosively with many elements and compounds such as Hydrogen and water. Elemental Fluorine is slightly basic, which means that when it reacts with water it forms OH−.

                             3F2 + 2H2O→O2 + 4HF
                             3F2 + 2H2O→O2 + 4HF
When combined with Hydrogen, Fluorine forms Hydrofluoric acid (HFHF), which is a weak acid. This acid is very dangerous and when dissociated can cause severe damage to the body because while it may not be painful initially, it passes through tissues quickly and can cause deep burns that interfere with nerve function.
                              HF + H2O→H3O+ + F−
                              HF + H2O→H3O+ + F−
There are also some organic compounds made of Fluorine, ranging from nontoxic to highly toxic. Fluorine forms covalent bonds with Carbon, which sometimes form into stable aromatic rings. When Carbon reacts with Fluorine the reaction is complex and forms a mixture of CF4CF4, C2F6C2F6, an C5F12C5F12.

                             C(s) + F2(g)→CF4(g) + C2F6 + C5F12
                             C(s) + F2(g)→CF4(g) + C2F6 + C5F12
Fluorine reacts with Oxygen to form OF2OF2 because Fluorine is more electronegative than Oxygen. The reaction goes:
                             2F2 + O2→2OF2
                             2F2 + O2→2OF2
Fluorine is so electronegative that sometimes it will even form molecules with noble gases like Xenon, such as the the molecule Xenon Difluoride, XeF2XeF2.

                              Xe + F2→XeF2
                              Xe + F2→XeF2
Fluorine also forms strong ionic compounds with metals. Some common ionic reactions of Fluorine are:
                              F2 + 2NaOH→O2 + 2NaF + H2
                              F2 + 2NaOH→O2 + 2NaF + H2
                              4F2 + HCl + H2O→3HF + OF2 + ClF3
                              4F2 + HCl + H2O→3HF + OF2 + ClF3
                              F2 + 2HNO3→2NO3F + H2


                     ISOTOPES
Only one isotope of fluorine occurs naturally in abundance, the stable isotope 19F. It has a high magnetogyric ratio and exceptional sensitivity to magnetic fields; because it is also the only stable isotope, it is used in magnetic resonance imaging.Seventeen radioisotopeswith mass numbers from 14 to 31 have been synthesized, of which 18F is the most stable with a half-life of 109.77 minutes. Other radioisotopes have half-lives less than 70 seconds; most decay in less than half a second. The isotopes 17F and 18F undergo β+ decayand electron capture, lighter isotopes decay by proton emission, and those heavier than 19F undergo β− decay (the heaviest ones with delayed neutron emission). Two metastable isomers of fluorine are known, 18mF, with a half-life of 162(7) nanoseconds, and 26mF, with a half-life of 2.2(1) milliseconds.
                      INDUSTRIAL APPLICATIONS
Fluorite mining, which supplies most global fluorine, peaked in 1989 when 5.6 million metric tons of ore were extracted. Chlorofluorocarbon restrictions lowered this to 3.6 million tons in 1994; production has since been increasing. Around 4.5 million tons of ore and revenue of US$550 million were generated in 2003; later reports estimated 2011 global fluorochemical sales at $15 billion and predicted 2016–18 production figures of 3.5 to 5.9 million tons, and revenue of at least $20 billion. Froth flotation separates mined fluorite into two main metallurgical grades of equal proportion: 60–85% pure metspar is almost all used in iron smelting whereas 97%+ pure acidspar is mainly converted to the key industrial intermediate hydrogen fluoride.

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.

• 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

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.

Bromine was discovered independently by two chemists, carl jacob lowig and antoine balard. Lowig isolated bromine from a mineral water spring which was situated in his hometown. Lowig used a solution of the mineral salt saturated with chlorine and and extracted the bromine with diethyl ether. After evaporation of the ether a brown liquid remained which was used as a sample of his work. Balard found bromine chemicals in the ash of seaweed from the salt marshes of Montpellier. The seaweed was used to produce iodine, but also contained bromine. Balard distilled the bromine from a solution of seaweed ash saturated with chlorine. The properties of the resulting substance were intermediate between those of chlorine and iodine; thus he tried to prove that the substance was iodine monochloride (ICl), but after failing to do so he was sure that he had found a new element, and named it muride, derived from the Latin word muria for brine.
After the French chemists Louis Nicolas Vauquelin, Louis Jacques Thénard, and Joseph-Louis Gay-Lussac approved the experiments of the young pharmacist Balard, the results were presented at a lecture of the Académie des Sciencesand published in Annales de Chimie et Physique. In his publication, Balard states that he changed the name from muride to brôme on the proposal of M. Anglada. Brôme (bromine) derives from the Greek βρωμος (stench). Other sources claim that the French chemist and physicist Joseph-Louis Gay-Lussac suggested the name brôme for the characteristic smell of the vapors. Bromine was not produced in large quantities until 1858, when the discovery of salt deposits in Stassfurt enabled its production as a by-product of potash.


Bromine is a chemical element with symbol Br and atomic number 35. It is the third-lightest halogen, and is a fuming red-brown liquid at room temperature that evaporates readily to form a similarly coloured gas. Its properties are thus intermediate between those of chlorine and iodine. Isolated independently by two chemists, Carl Jacob Löwig (in 1825) and Antoine Jérôme Balard (in 1826), its name was derived from the Ancient Greek βρῶμος ("stench"), referencing its sharp and disagreeable smell




The Hunsdiecker reaction (also called the Borodin reaction or the Hunsdiecker–Borodin reaction) is a name reaction in organic chemistry whereby silver salts of carboxylic acids react with a halogen to produce an organic halide.[1] It is an example of both a decarboxylation and a halogenation reaction as the product has one fewer carbon atoms than the starting material (lost as carbon dioxide) and a halogen atom is introduced its place. The reaction was first demonstrated by Alexander Borodin in his 1861 reports of the preparation of methyl bromide from silver acetate.[2][3] Shortly after, the approach was applied to the degradation of fatty acids in the laboratory of Adolf Lieben.[4][5] However, it is named for Cläre Hunsdiecker and her husband Heinz Hunsdiecker, whose work in the 1930s[6][7] developed it into a general method.[1] Several reviews have been published,[8][9] and a catalytic approach has been developed.[10]
 
 Occurrence in nature
 
 Bromine is too reactive to exist as a free element in nature. Instead, it occurs in compounds, the most common of which are sodium bromide (NaBr) and potassium bromide (KBr). These compounds are found in seawater and underground salt beds. These salt beds were formed in regions where oceans once covered the land. When the oceans evaporated (dried up), salts were left behind—primarily sodium chloride (NaCl), potassium chloride (KCl), and sodium and potassium bromide. Later, movements of the Earth's crust buried the salt deposits. Now they are buried miles underground. The salts are brought to the surface in much the same way that coal is mined.
Bromine is a moderately abundant element. Its abundance in the Earth's crust is estimated to be about 1.6 to 2.4 parts per million. It is far more abundant in seawater where it is estimated at about 65 parts per million.
In some regions, the abundance of bromine is even higher. For example, the Dead Sea (which borders Israel and Jordan), has a high level of dissolved salts. The abundance of bromine there is estimated to be 4,000 parts per million. The salinity, or salt content, is so high that nothing lives in the water. This is why it is called the Dead Sea.


Isotopes


Two naturally existing isotopes of bromine exist, bromine-79 and bromine-81. Isotopes are two or more forms of an element. Isotopes differ from each other according to their mass number. The number written to the right of the element's name is the mass number. The mass number represents the number of protons plus neutrons in the nucleus of an atom of the element. The number of protons determines the element, but the number of neutrons in the atom of any one element can vary. Each variation is an isotope.
At least 16 radioactive isotopes of bromine are known also. A radioactive isotope is one that breaks apart and gives off some form of radiation. Radioactive isotopes are produced when very small particles are fired at atoms. These particles stick in the atoms and make them radioactive.
No isotope of bromine has any important commercial use.
The salinity, or salt content, is so high that nothing lives in the water. This is why it is called the Dead Sea.

Iodine is a chemical element with symbol I and atomic number 53. The heaviest of the stable halogens, it exists as a lustrous, purple-black non-metallic solid at standard conditions that melts to form a deep violet liquid at 114 degrees Celsius, and boils to a violet gas at 184 degrees Celsius. The element was discovered by the French chemist Bernard Courtois in 1811. It was named two years later by Joseph Louis Gay-Lussac from this property.
Iodine occurs in many oxidation states, including iodide , iodate, and the various periodate anions. It is the least abundant of the stable halogens, being the sixty-first most abundant element. It is the heaviest essential mineral nutrient. Iodine is essential in the synthesis of thyroid hormones.[4] Iodine deficiency affects about two billion people and is the leading preventable cause of intellectual disabilities.
The dominant producers of iodine today are Chile and Japan. Iodine and its compounds are primarily used in nutrition. Due to its high atomic number and ease of attachment to organic compounds, it has also found favour as a non-toxic radiocontrast material. Because of the specificity of its uptake by the human body, radioactive isotopes of iodine can also be used to treat thyroid cancer. Iodine is also used as a catalyst in the industrial production of acetic acid and some polymers.

• Appearance: lustrous metallic gray, violet as a gas
• Standard Atomic Weight: 126.904
• Atomic Number(Z): 53
• Group: group 17 (halogens)
• Period: period 5
• Block: p block
• Element Category: reactive non-metal
• PHASE AT STP : solid
• Melting Point : 386.85 K [113.7 ° C]
• Boiling Point : 457.4 K [184.3 ° C]
• Density : 4.933 g/ml
• Heat of Vaporization : 41.57 kj/mol
• Covalent Radius : 139+-3pm
• Crystal Structure : orthorhombic
• Discovery :  Bernard Courtosis (1811)

                   HISTORY
In 1811, iodine was discovered by French chemist Bernard Courtois who was born to a manufacturer of saltpetre (an essential component of gunpowder). At the time of the Napoleonic Wars, saltpetre was in great demand in France. Saltpetre produced from French nitre beds required sodium carbonate, which could be isolated from seaweed collected on the coasts of Normandy and Brittany. To isolate the sodium carbonate, seaweed was burned and the ash washed with water. The remaining waste was destroyed by adding sulfuric acid. Courtois once added excessive sulfuric acid and a cloud of purple vapour rose. He noted that the vapour crystallised on cold surfaces, making dark crystals Courtois suspected that this material was a new element but lacked funding to pursue it further.
                     PROPERTIES
Iodine is the fourth halogen, being a member of group 17 in the periodic table, below fluorine, chlorine, and bromine; it is the heaviest stable member of its group. Like the other halogens, it is 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, although in keeping with periodic trends, it is the weakest oxidising agent among the stable halogens: it has the lowest electronegativity among them, just 2.66 on the Pauling scale (compare fluorine, chlorine, and bromine at 3.98, 3.16, and 2.96 respectively; astatine continues the trend with an electronegativity of 2.2). Elemental iodine hence forms diatomic molecules with chemical formula I2, where two iodine atoms share a pair of electrons in order to each achieve a stable octet for themselves; at high temperatures, these diatomic molecules reversibly dissociate a pair of iodine atoms. Similarly, the iodide anion, I−, is the strongest reducing agent among the stable halogens, being the most easily oxidised back to diatomic I2. The halogens darken in colour as the group is descended: fluorine is a very pale yellow gas, chlorine is greenish-yellow, and bromine is a reddish-brown volatile liquid. Iodine conforms to the prevailing trend, being a shiny black crystalline solid that melts at 114 °C and boils at 183 °C to form a violet gas. This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group (though astatine may not conform to it, depending on how metallic it turns out to be). Specifically, the violet colour of iodine gas results from the electron transition between the highest occupied antibonding πg molecular orbital and the lowest vacant antibonding σu molecular orbital.

                                 REACTIONS OF IODINE



                        CHEMISTRY AND COMPOUNDS
Though it is the least reactive of the halogens, iodine is still one of the more reactive elements. For example, while chlorine gas will halogenate carbon monoxide, nitric oxide, and sulfur dioxide (to phosgene, nitrosyl chloride, and sulfuryl chloride respectively), iodine will not do so. Furthermore, iodination of metals tends to result in lower oxidation states than chlorination or bromination; for example, rhenium metal reacts with chlorine to form rhenium hexachloride, but with bromine it forms only rhenium pentabromide and iodine can achieve only rhenium tetraiodide. By the same token, however, since iodine has the lowest ionisation energy among the halogens and is the most easily oxidised of them, it has a more significant cationic chemistry and its higher oxidation states are rather more stable than those of bromine and chlorine, for example in iodine heptafluoride.

                  APPLICATIONS
Unlike chlorine and bromine, which have one significant main use dwarfing all others, iodine is used in many applications of varying importance. About half of all produced iodine goes into various organoiodine compounds; another 15% remains as the pure element, another 15% is used to form potassium iodide, and another 15% for other inorganic iodine compounds. The remaining 5% is for minor uses. Among the major uses of iodine compounds are catalysts, animal feed supplements, stabilisers, dyes, colourants and pigments, pharmaceutical, sanitation (from tincture of iodine), and photography; minor uses include smog inhibition, cloud seeding, and various uses in analytical chemistry.

1) INTRODUCTION :

Astatine is a radioactive chemical element with symbol At and atomic number 85. It is the rarest naturally occurring element in the Earth's crust, occurring only as the decay product of various heavier elements. All of astatine's isotopes are short-lived; the most stable is astatine-210, with a half-life of 8.1 hours. A sample of the pure element has never been assembled, because any macroscopic specimen would be immediately vaporized by the heat of its own radioactivity.

SYMBOL: At

ELECTRONIC CONFIGURATION: [Xe] 4f14 5d10 6s2 6p5

Discovered: 1940




2) CHARECTERISTICS :

Astatine is an extremely radioactive element; all its isotopes have short half-lives of 8.1 hours or less, decaying into other astatine isotopes, bismuth, polonium or radon.

A) PHYSICAL PROPERTY :

Most of the physical properties of astatine have been estimated, using theoretically or empirically derived methods. For example, halogens get darker with increasing atomic weight – fluorine is nearly colorless, chlorine is yellow-green, bromine is red-brown, and iodine is dark gray/violet. Astatine is sometimes described as probably being a black solid .
The melting and boiling points of astatine are also expected to follow the trend seen in the halogen series, increasing with atomic number. On this basis they are estimated to be 575 and 610 K (302 and 337 °C; 575 and 638 °F), respectively. Some experimental evidence suggests astatine may have lower melting and boiling points than those implied by the halogen trend. Astatine sublimes less readily than does iodine, having a lower vapor pressure. Even so, half of a given quantity of astatine will vaporize in approximately an hour if put on a clean glass surface at room temperature. The absorption spectrum of astatine in the middle ultraviolet region has lines at 224.401 and 216.225 nm, suggestive of 6p to 7s transitions.
The structure of solid astatine is unknown. As an analogue of iodine it may have an orthorhombic crystalline structurecomposed of diatomic astatine molecules, and be a semiconductor (with a band gap of 0.7 eV).


B) CHEMICAL PROPERTY :


The chemistry of astatine is "clouded by the extremely low concentrations at which astatine experiments have been conducted, and the possibility of reactions with impurities, walls and filters, or radioactivity by-products, and other unwanted nano-scale interactions." Many of its apparent chemical properties have been observed using tracer studies on extremely dilute astatine solutions, typically less than 10−10 mol·L−1. Some properties – such as anion formation – align with other halogens. Astatine has some metallic characteristics as well, such as plating onto a cathode, coprecipitating with metal sulfides in hydrochloric acid and forming a stable monatomic cation in aqueous solution. It forms complexes with EDTA, a metal chelating agent, and is capable of acting as a metal in antibody radiolabeling; in some respects astatine in the +1 state is akin to silver in the same state. Most of the organic chemistry of astatine is, however, analogous to that of iodine.


3) ISOTOPES :


There are 39 known isotopes of astatine, with atomic masses (mass numbers) of 191–229. Theoretical modeling suggests that 37 more isotopes could exist. No stable or long-lived astatine isotope has been observed, nor is one expected to exist.


4) SYNTHESIS