Halogen
Template:Periodic table (halogens) The halogens or halogen elements are a series of nonmetal elements from group 17 of the periodic table (formerly: VII, VIIA), comprising fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The artificially created element 117 (ununseptium), may also be a halogen.
The group of halogens is the only periodic table group which contains elements in all three familiar states of matter at standard temperature and pressure.
History and etymology
In 1842 the Swedish chemist Baron Jöns Jakob Berzelius proposed the term "halogen" – ἅλς (háls), "salt" or "sea", and γεν- (gen-), from γίγνομαι (gígnomai), "come to be" – for the four elements (fluorine, chlorine, bromine, and iodine) that produce a sea-salt-like substance when they form a compound with a metal.[1] Earlier, in 1811, the word "halogen" had been proposed as a name for the newly discovered element chlorine, but Davy's proposed term for this element eventually won out.
Group trends
Like other groups, the candidates of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior:
| Z | Element | No. of electrons/shell |
|---|---|---|
| 9 | fluorine | 2, 7 |
| 17 | chlorine | 2, 8, 7 |
| 35 | bromine | 2, 8, 18, 7 |
| 53 | iodine | 2, 8, 18, 18, 7 |
| 85 | astatine | 2, 8, 18, 32, 18, 7 |
The halogens show a series of trends when moving down the group—for instance, decreasing electronegativity and reactivity, and increasing melting and boiling point.
| Halogen | Standard Atomic Weight (u) | Melting Point (K) | Boiling Point (K) | Electronegativity (Pauling) |
| Fluorine | 18.998 | 53.53 | 85.03 | 3.98 |
| Chlorine | 35.453 | 171.60 | 239.11 | 3.16 |
| Bromine | 79.904 | 265.80 | 332.00 | 2.96 |
| Iodine | 126.904 | 386.85 | 457.40 | 2.66 |
| Astatine | (210) | 575.00 | 610 (?) | 2.20 |
Diatomic halogen molecules
| halogen | molecule | structure | model | d(X−X) / pm (gas phase) |
d(X−X) / pm (solid phase) |
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The elements become less reactive and have higher melting points as the atomic number increases.
Chemistry
Reactivity
Halogens are highly reactive, and as such can be harmful or lethal to biological organisms in sufficient quantities. This high reactivity is due to the atoms being highly electronegative due to their high effective nuclear charge. They can gain an electron by reacting with atoms of other elements. Fluorine is one of the most reactive elements in existence, attacking otherwise inert materials such as glass, and forming compounds with the heavier noble gases. It is a corrosive and highly toxic gas. The reactivity of fluorine is such that if used or stored in laboratory glassware, it can react with glass in the presence of small amounts of water to form silicon tetrafluoride (SiF4). Thus fluorine must be handled with substances such as Teflon (which is itself an organofluorine compound), extremely dry glass, or metals such as copper or steel which form a protective layer of fluoride on their surface.
The high reactivity of fluorine means that once it does react with something, it bonds with it so strongly that the resulting molecule is very inert and non-reactive to anything else. For example, Teflon is fluorine bonded with carbon.
Both chlorine and bromine are used as disinfectants for drinking water, swimming pools, fresh wounds, spas, dishes, and surfaces. They kill bacteria and other potentially harmful microorganisms through a process known as sterilization. Their reactivity is also put to use in bleaching. Sodium hypochlorite, which is produced from chlorine, is the active ingredient of most fabric bleaches and chlorine-derived bleaches are used in the production of some paper products. Chlorine also reacts with sodium to create sodium chloride, which is another name for table salt.
Hydrogen halides
The halogens all form binary compounds with hydrogen known as the hydrogen halides: hydrogen fluoride (HF), hydrogen chloride (HCl), hydrogen bromide (HBr), hydrogen iodide (HI), and hydrogen astatide (HAt). All of these are strong chemical acids when dissolved in water, with the exception of HF. However, hydrofluroic acid does have quite destructive properties towards animal tissues, including those of human beings. When in aqueous solution, the hydrogen halides are known as hydrohalic acids. The names of these acids are as follows: hydrofluoric acid, [[hydrochloric acid], hydrobromic acid, and hydroiodic acid. All of these acids must be handled with great care because they are dangerous!
Hydrogen astatide should also be a strong acid, but it is seldom included in presentations about hydrohalic acids because of the extreme radioactivity of astatine (via alpha decay).
Metal halides
The halogens form many compounds with metals. These compounds range from highly ionic compounds like sodium chloride, monomeric covalent compounds like uranium hexafluoride, and polymeric covalent compounds like palladium chloride. Metal halides are generally obtained by direct combination, or more commonly, neutralization of basic metal salt with a hydrohalic acid. They serve as useful entry points into inorganic chemistry.
Interhalogen compounds
The halogens react with each other to form interhalogen compounds. Diatomic interhalogen compounds such as BrF, ICl, and ClF bear resemblance to the pure halogens in some respects. The properties and behaviour of a diatomic interhalogen compound tend to be intermediate between those of its parent halogens. Some properties, however, are found in neither parent halogen. For example, Cl2 and I2 are soluble in CCl4, but ICl is not since it is a polar molecule due to the relatively large electronegativity difference between I and Cl.
Organohalogen compounds
Many synthetic organic compounds such as plastic polymers, and a few natural ones, contain halogen atoms; these are known as halogenated compounds or organic halides. Chlorine is by far the most abundant of the halogens, and the only one needed in relatively large amounts (as chloride ions) by humans. For example, chloride ions play a key role in brain function by mediating the action of the inhibitory transmitter GABA and are also used by the body to produce stomach acid. Iodine is needed in trace amounts for the production of thyroid hormones such as thyroxine. On the other hand, neither fluorine nor bromine are believed to be essential for humans. Organohalogens are also synthesized through the nucleophilic abstraction reaction.
Polyhalogenated compounds
Polyhalogenated compounds are industrially created compounds substituted with multiple halogens. Many of them are very toxic and bioaccumulate in humans, and have a very wide application range. They include the much maligned PCBs, PBDEs, and PFCs as well as numerous other compounds.
Reactivity with water
Fluorine reacts vigorously with water to produce oxygen (O2) and hydrogen fluoride (HF):[2]
- 2 F2(g) + 2 H2O(l) → O2(g) + 4 HF(aq)
Chlorine has maximum solubility of ca. 7.1 g Cl2 per kg of water at ambient temperature (21 °C).[3] Dissolved chlorine reacts to form hydrochloric acid (HCl) and hypochlorous acid, a solution that can be used as a disinfectant or bleach:
- Cl2(g) + H2O(l) → HCl(aq) + HClO(aq)
Bromine has a solubility of 3.41 g per 100 g of water,[4] but it slowly reacts to form hydrogen bromide (HBr) and hypobromous acid (HBrO):
- Br2(g) + H2O(l) → HBr(aq) + HBrO(aq)
Iodine, however, is minimally soluble in water (0.03 g/100 g water at 20 °C) and does not react with it.[5] However, iodine will form an aqueous solution in the presence of iodide ion, such as by addition of potassium iodide (KI), because the triiodide ion is formed.
Production
The element astatine is produced artificially nearly all the time, and only minute amounts are found in nature. Ununseptium is always produced artificially. Note that it is possible that ununseptium may not be a halogen. Nearly all the facts about ununseptium are predicted.
Applications
Drug discovery
In drug discovery, the incorporation of halogen atoms into a lead drug candidate results in analogues that are usually more lipophilic and less water soluble.[6] Consequently, halogen atoms are used to improve penetration through lipid membranes and tissues. It follows that there is a tendency for some halogenated drugs to accumulate in adipose tissue.
The chemical reactivity of halogen atoms depends on both their point of attachment to the lead and the nature of the halogen. Aromatic halogen groups are far less reactive than aliphatic halogen groups, which can exhibit considerable chemical reactivity. For aliphatic carbon-halogen bonds the C-F bond is the strongest and usually less chemically reactive than aliphatic C-H bonds. The other aliphatic-halogen bonds are weaker, their reactivity increasing down the periodic table. They are usually more chemically reactive than aliphatic C-H bonds. Consequently, the most common halogen substitutions are the less reactive aromatic fluorine and chlorine groups.
Biological role
The halogens chlorine and iodine are needed by our bodies in trace amounts. We eat chlorine in salts, and we eat iodine in vegetables. The other three halogens have no known biological role.
Toxicity
Halogens are very reactive elements and their toxicity derives from their strong oxidant properties. Fluorine is the most toxic and dangerous of the halogens, capable of killing a person within seconds. Chlorine is also toxic, and can kill a person if inhaled in an enclosed area. Bromine is less reactive and toxic than Chlorine but, being a liquid, is denser and the contact with it can be fatal. Besides, is more reactive towards water than Chlorine, and can lead to the formation of a higher quantity of acids in the lungs. Iodine is the less dangerous of halogens and can hardly kill a person if inhaled, but the ingestion can be fatal even at a low dose (28 mg/kg).
See also
References
- ^ Online Etymology Dictionary halogen.
- ^ The Oxidising Ability of the Group 7 Elements. Chemguide.co.uk. Retrieved on 2011-12-29.
- ^ of chlorine in water. Resistoflex.com. Retrieved on 2011-12-29.
- ^ Properties of bromine. bromaid.org
- ^ MSDS. Hazard.com (1998-04-21). Retrieved on 2011-12-29.
- ^ G. Thomas, Medicinal Chemistry an Introduction, John Wiley & Sons, West Sussex, UK, 2000.
Further reading
- N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, 2nd ed., Butterworth-Heinemann, Oxford, UK, 1997.
