What are the Halogens in the Periodic Table? - Quora Quora Daniel James Berger , PhD in organic chemistry The elements in Group 17 of the periodic table ( The Old fashioned name for Group 17 is Group VII ) are known collectively as The Halogens. They are the elements Fluorine, ( atomic number 9 ), Chlorine, ( atomic number 17 ), Bromine ( atomic number 35 ), Iodine ( atomic number 53 ), Astatine ( atomic number 85 ), and a newly discovered element not yet given a name, that has an atomic number of 117, ( Discovered in 2010). However both Astatine and Element No. 117 are too radioactive to be made in quantities that can be weighed. They are man made elements that have very short half lives; so decay into other elements before one has the chance to make any of their compounds. So in most laboratories Halogen chemistry is essentially about the study of Fluorine, Chlorine, Bromine & Iodine. The word Halogen was coined in 1842 by the Swedish Chemist Berzelius ( 1779 - 1848 ).

It is a made up Swedish word which literally means "Salt producer" in Swedish. A name in keeping with the reactive nature of these elements. Lots of Halogen salts are known. Eg. Sodium chloride. Potassium fluoride. Sodium Bromide. These elements are all members of the same group, because all of their atoms have the same number of electrons in their outermost shell or orbit. Group 17 elements all have 7 electrons in their outermost orbit. ​ ​The Periodic Table highlighting Group 17. Fluorine As mentioned previously fluorine has an atomic number of 9. This means that an atom of fluorine will have in addition to 9 protons in the nucleus of its atom, it will have a total of 9 electrons orbiting its nucleus. An individual proton always has a charge of +1 and an individual electron always has a charge of -1. We have 9 protons in an atom of fluorine each having a charge of +1, plus 9 electrons in that same atom orbiting the nucleus of that atom. ( +1 X 9 ) + (-1 X 9 ) = 0. So the atom is electrically neutral.

This is true for all neutral atoms of fluorine. The 9 electrons are arranged in 2 shells or orbit. 2 of the 9 electrons orbit in the 1st shell or orbit. The remaining 7 electrons orbit in a larger shell. The 7 electrons in the outermost orbit of all Halogens are the electrons that participate in chemical reactions. During chemical reactions the number of electrons in the outermost orbit can go up or down. Up to a max of 8, and down to 0 ( in theory ). Compare with Na+1 or with Li+1 ​ Written Feb 7, 2016 The halogens are the family of chemical elements that includes fluorine (atomic symbol F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The halogens make up Group VIIA of the Periodic Table of the elements. Elemental halogens are diatomic molecules. However, due to their high reactivity, the halogens are never found in nature in native form. The family name means "salt-forming," from the Greek for salt, "hals", and for generating "genes".

The salinity of the oceans on Earth is due in large part to such halogen salts (halides) as sodium chloride (NaCl) and potassium iodide (KI). Halogens display physical and chemical properties typical of nonmetals. They have relatively low melting and boiling points that increase steadily down the group. Near room temperature, the halogens span all of the physical states: Fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. All of the elements are colored, with the color becoming more intense moving down the group. Fluorine gas is pale yellow, and chlorine gas is a yellowish green. Liquid bromine and its vapors are brownish red. Solid iodine appears as shiny, dark gray crystals, and the vapors are a deep purple. The halogens are poor thermal and electrical conductors in all phases, and as solids they are brittle and crumbly. The halogens have distinctive, unpleasant odors, will burn exposed flesh, and are toxic. The neutral atoms of the halogens possess seven outer electrons. An additional electron can be added to halogen atoms to form singly charged negative ions. These ions have a closed outer-shell configuration.

Electronegativity is a measure of the ability of an atom of one element to remove an electron from an atom of another element. As a group, the halogens are among the most electronegative elements. Fluorine has the highest electronegativity of all the elements. Halogens are so reactive that all the elements except helium and neon have been found to react with at least one of the halogens. Fluorine is always assigned a formal oxidation number of –1, whereas the other halogens can exhibit a range of oxidation numbers 526 Views · View Upvotes · Answer requested by Properties Of Halogens, Atomic Radii, Ionization Energy - Transtutors Properties Of Halogens Atomic and ionic radii A halogen atom has the smallest radius as compared to any other element in its period. This is due to the increased effective nuclear charge which results in greater attraction of the electrons by the nucleus. The atomic radii. Increase from fluorine to iodine down the group due to increase in number of shells. 220 (2) Ionization energy : Ionization energy of these elements are higher than those of the corresponding elements of group 16 due to increased nuclear charge.

these values decrease systematically as we move down the group from F to 1 1008 Thus iodine which has a comparatively low value of I.E., has a tendency to lose an electon to form positive iodinium ion, I+ and thus shown electropositive or metallic character. (3) Electronegativity : Fluorine is the most electronegative element in the periodic table. With increase in atomic number down the group, the electronegativity decreases. The decreasing order of electronegativity is F > Cl > Br > I. Email Based Homework Assignment Help in Properties of Halogens Transtutors is the best place to get answers to all your doubts regarding properties of halogens, atomic radii, ionization energy and electronegativity with examples. You can submit your school, college or university level homework or assignment to us and we will make sure that you get the answers you need which are timely and also cost effective. Our tutors are available round the clock to help you out in any way with chemistry. Live Online Tutor Help for Properties of Halogens Transtutors has a vast panel of experienced chemistry tutors who specialize in properties of halogensand can explain the different concepts to you effectively.

You can also interact directly with our chemistry tutors for a one to one session and get answers to all your problems in your school, college or university level inorganic chemistry. Our tutors will make sure that you achieve the highest grades for your chemistry assignments. We will make sure that you get the best help possible for exams such as the AP, AS, A level, GCSE, IGCSE, IB, Round Square etc. Related Topics Halogens - Chemistry Encyclopedia - uses, elements, gas, number, name, symbol, salt, atom Halogens - Chemistry Encyclopedia Halogens Photo by: The halogens are the family of chemical elements that includes fluorine (atomic symbol F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The halogens make up Group VIIA of the Periodic Table of the elements. Elemental halogens are diatomic molecules. However, due to their high reactivity, the halogens are never found in nature in native form. The family name means "salt-forming," from the Greek for salt, halos, and for generating genes.

The salinity of the oceans on Earth is due in large part to such halogen salts (halides) as sodium chloride (NaCl) and potassium iodide (KI). Halogens display physical and chemical properties typical of nonmetals. They have relatively low melting and boiling points that increase steadily down the group. Near room temperature, the halogens span all of the physical states: Fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. All of the elements are colored, with the color becoming more intense moving down the group. Fluorine gas is pale yellow, and chlorine gas is a yellowish green. Liquid bromine and its vapors are brownish red. Solid iodine appears as shiny, dark gray crystals, and the vapors are a deep purple. The halogens are poor thermal and electrical conductors in all phases, and as solids they are brittle and crumbly. The halogens have distinctive, unpleasant odors, will burn exposed flesh, and are toxic. The neutral atoms of the halogens possess seven outer electrons. An additional electron can be added to halogen atoms to form singly charged negative ions. These ions have a closed outer-shell configuration.

Electronegativity is a measure of the ability of an atom of one element to remove an electron from an atom of another element. As a group, the halogens are among the most electronegative elements. Fluorine has the highest electronegativity of all the elements. Halogens are so reactive that all the elements except helium and neon have been found to react with at least one of the halogens. Fluorine is always assigned a formal oxidation number of –1, whereas the other halogens can exhibit a range of oxidation numbers. The ability of halogens to form chemical compounds with all metals and most nonmetals has led to a wide variety of uses for these elements. Chlorine is used as a bleach and a disinfectant. Iodine has been used as a topical microbicide. Iodine and bromine are added to halogen lamps to lengthen Chlorine gas belongs to the halogen chemical family. the life of the filament and prevent darkening of the bulb. Chloride and iodide are essential dietary minerals for humans.

Organic compounds containing halogens are used as fire-retardants (halons), as refrigerants (Freons), and in nonstick coatings (Teflon). Silver bromide and silver iodide have been used in photographic emulsions since the early days of photography. Many halogenated compounds are toxic. A well-known example is DDT (dichlorodiphenyltrichloroethane), once a widely applied pesticide that was banned in the United States after severe environmental effects were discovered. All known isotopes of astatine are radioactive, with the longest-lived isotope having a half-life of about eight hours. Relatively little is known of the physical and chemical properties of astatine. However, it is predicted to have properties similar to iodine. SEE ALSO John Michael Nicovich Bibliography Lide, David R., ed. (2003). The CRC Handbook of Chemistry and Physics , 84th edition. Boca Raton, FL: CRC Press. Internet Resources BBC Bitesize - GCSE Chemistry - Group 7 - the halogens - Revision 2 acidic Having a pH less than 7. aqueous Dissolved in water to form a solution.

Shown as (aq) in chemical equations. atom All elements are made of atoms. An atom consists of a nucleus containing protons and neutrons, surrounded by electrons. conduct To allow electricity, heat or other energy forms to pass through. covalent bond A bond between atoms formed when atoms share electrons to achieve a full outer shell of electrons. electron Sub-atomic particle, with a negative charge and a negligible mass relative to protons and neutrons. element A substance made of one type of atom only. group 1 The first vertical column of elements in the periodic table, starting with lithium and ending with francium. Also called the alkali metals. halogen An element placed in Group 7 of the period table, which starts with fluorine and ends with astatine. ion Electrically charged particle, formed when an atom or molecule gains or loses electrons. ionic equation A chemical equation that shows how positively charged ions join with negatively charged ions to make a compound. molecule A collection of two or more atoms held together by chemical bonds. periodic table A tabular representation of all known elements in order based on atomic number, eg all the noble elements are found on the right of the periodic table.

reactivity series A list of substances in order of their reactivity, usually from most reactive to least reactive. solvent The liquid in which the solute dissolves to form a solution. vapour Vapour is a cloud of liquid particles. Steam is water vapour. BBC - GCSE Bitesize: Electron arrangement and chemical properties of Group 7 - Higher Next Electron arrangement and chemical properties of Group 7 - Higher The atoms [atoms: An atom is the smallest part of an element that still has the properties of that element, comprising electrons surrounding a nucleus of protons and neutrons. ] of the elementselements: All atoms of an element have the same atomic number, the same number of protons and electrons, and so the same chemical properties. in Group 7 (also called the halogens) have seven electrons [electrons: Sub-atomic particles, with a negative charge and a negligible mass relative to protons and neutrons. ] in their highest occupied energy level (the outer shell). This is why their chemical properties are similar. Seven outer electrons When you write the electronic structure of a halogen, the last number must be a 7.

When you draw or complete a diagram showing the electronic structure of a halogen, there must be seven dots or crosses in the outer circle. Gain of electrons In a reaction with a metal, each halogen atom gains an outer electron and becomes an ion [ion: Positively- or negatively-charged particles - eg positively charged hydrogen, sodium and potassium atoms. Ion charge helps determine a substance's acidity or alkalinity ] with a single negative charge, -1. Larger atoms = less reactive As you go down Group 7 the atoms become larger, and the highest occupied energy level (the outer shell) becomes further from the nucleus. The force of attraction between the positively-charged nucleus and a negatively-charged electron from another atom becomes weaker. As a result, it becomes harder to attract and gain an electron the larger the atom becomes. The more difficult it is to gain these outer electrons, the less reactive a halogen is. For example chlorine is less reactive than fluorine because the outer electrons in a chlorine atom are further from the nucleus than the outer electrons in a fluorine atom. It is harder for a chlorine atom to gain an electron than it is for a fluorine atom.

Astatine atoms, with 6 shells, are the largest atoms in Group 7. They are very unreactive - now you know why. Table showing different electron arrangements in Group 7 Diagram Properties of the Halogens About Watch and Favorite Watch Watching this resources will notify you when proposed changes or new versions are created so you can keep track of improvements that have been made. Favorite Favoriting this resource allows you to save it in the “My Resources” tab of your account. There, you can easily access this resource later when you’re ready to customize it or assign it to your students. Properties of the Halogens Halogens (fluorine, chlorine, bromine, iodine, astatine) are nonmetal elements that are highly electronegative and reactive. Learning Objective Describe the physical and chemical properties of halogens. Key Points Halogens are nonmetals in group 17 (or VII) of the periodic table. Down the group, atom size increases. As a diatomic molecule , fluorine has the weakest bond due to repulsion between electrons of the small atoms. Due to increased strength of Van der Waals forces down the group, the boiling points of halogens increase.

Therefore, the physical state of the elements down the group changes from gaseous fluorine to solid iodine. Due to their high effective nuclear charge , halogens are highly electronegative . Therefore, they are highly reactive and can gain an electron through reaction with other elements. Halogens can be harmful or lethal to biological organisms in sufficient quantities. Terms Register for FREE to remove ads and unlock more features! Learn more Full Text The halogens are a series of non-metal elements from group 17 of the periodic table (formerly VII). The halogens include fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The artificially created element 117 (ununseptium) may also be considered a halogen. Physical Properties Atoms get bigger down the group as additional electron shells are filled. When fluorine exists as a diatomic molecule, the F–F bond is unexpectedly weak. This is because fluorine atoms are the smallest of the halogens—the atoms are bonded close together, which leads to repulsion between free electrons in the two fluorine atoms.

The boiling points of halogens increase down the group due to the increasing strength of Van der Waals forces as the size and relative atomic mass of the atoms increase. This change manifests itself in a change in the phase of the elements from gas (F2, Cl2) to liquid (Br2), to solid (I2). The halogens are the only periodic table group containing elements in all three familiar states of matter (solid, liquid, and gas) at standard temperature and pressure . Halogens represents all of the three familiar states of matter: (left to right) chlorine is a gas, bromine is a liquid, and iodine is a solid. Highly reactive fluorine is not included in the picture. Chemical Properties Electronegativity is the ability of an atom to attract electrons or electron density towards itself within a covalent bond . Electronegativity depends upon the attraction between the nucleus and bonding electrons in the outer shell. This, in turn, depends on the balance between the number of protons in the nucleus, the distance between the nucleus and bonding electrons, and the shielding effect of inner electrons.

In hydrogen halides (HX, where X is the halogen), the H-X bond gets longer as the halogen atoms get larger. This means the shared electrons are further from the halogen nucleus, which increases the shielding of inner electrons. This means electronegativity decreases down the group. Halogens are highly reactive, and they can be harmful or lethal to biological organisms in sufficient quantities. This reactivity is due to high electronegativity and high effective nuclear charge. Halogens can gain an electron by reacting with atoms of other elements. Fluorine is one of the most reactive elements. It reacts with otherwise inert materials such as glass , and it forms compounds with the heavier noble gases . It is a corrosive and highly toxic gas. Fluorine's reactivity means that once it does react with something, it bonds so strongly that the resulting molecule is inert and non-reactive. Fluorine can react with glass in the presence of small amounts of water to form silicon tetrafluoride (SiF4). Thus fluorine must be handled with substances like the inert organofluorine compound Teflon.

Fluorine reacts vigorously with water to produce oxygen (O2) and hydrogen fluoride: $2 F_2 (g) + 2 H_2O (l) \rightarrow O_2 (g) + 4 HF (aq)$ Chlorine has maximum solubility of 7.1 g per kg of water at ambient temperature (21 °C). Dissolved chlorine reacts to form hydrochloric acid (HCl) and hypochlorous acid (HClO), a solution that can be used as a disinfectant or bleach: $Cl_2 (g) + H_2O (l) \rightarrow HCl (aq) + HClO (aq)$ Bromine has a solubility of 3.41 g per 100 g of water. It slowly reacts to form hydrogen bromide (HBr) and hypobromous acid (HBrO): $Br_2 (g) + H_2O (l) \rightarrow HBr (aq) + HBrO (aq)$ Iodine is minimally soluble in water, with a solubility of 0.03 g per 100 g water.

However, iodine will form an aqueous solution in the presence of iodide ion . This occurs with the addition of potassium iodide (KI), forming a triiodide ion. The Halogens (Group 17) About Watch and Favorite Watch Watching this resources will notify you when proposed changes or new versions are created so you can keep track of improvements that have been made. Favorite Favoriting this resource allows you to save it in the “My Resources” tab of your account. There, you can easily access this resource later when you’re ready to customize it or assign it to your students. The Halogens (Group 17) The halogens are a series of highly reactive, nonmetal elements from Group 17 of the periodic table. Learning Objective Describe the general properties of the halogen elements. Key Points All halogens are electronegative: they gain electrons very quickly, making them the most reactive of all chemical elements . Halogens easily dissociate into atomic particles and can combine with surrounding elements to form compounds . Halogens are generally poisonous; each has a different level of toxicity. Halogens are poor conductors of heat and electricity.

Halogens can exist in all three classical states of matter: solid , liquid , and gas . Terms Full Text Overview The halogens are nonmetallic elements in Group 17 of the periodic table. They include fluorine, chlorine, bromine, iodine, and astatine. Ununseptium, which is not a naturally occurring element, is also often considered a halogen. As typical nonmetals, they have low melting and boiling points . A unique characteristic of halogens is that they are the only element group that includes elements that belong to all three classical states of matter: solid, liquid, and gas at room temperature . The halogens, as a group, are extremely reactive. In their elemental form, they are found as diatomic molecules . However, due to their extremely high reactivity (electronegativity), they are not found in their elemental form in any natural environment on Earth. Safety and Storage for Elemental Halogens Halogens are generally toxic , or even lethal, to biological organisms in their elemental forms. They tend to combine readily, and explosively, with metals , particularly with alkali metals .

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 it is used or stored in laboratory glassware, it can react with glass in the presence of small amounts of water to form silicon tetrafluoride (SiF4). Therefore, 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 that form a protective layer of fluoride on their surface. The high reactivity of fluorine also 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 a polymer of fluorine bonded with carbon. The halogens are highly reactive, diatomic molecules. Reactivity of Halogens 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 acids when dissolved in water, with the exception of HF. However, hydrofluoric acid is quite destructive to animal tissue. 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 extremely dangerous. Some of these acids are also widely used in industry. The halogens can also react with each other to form interhalogen compounds. Reactivity of Halogens and Water Fluorine reacts vigorously with water to produce oxygen (O2) and hydrofluoric acid (HF): $2F_2 (g) + 2H_2O (l) \rightarrow O_2 (g) + 4HF (aq)$ Chlorine has a maximum solubility of about 7.1 g Cl2/kg H2O at ambient temperature (21 °C).

Dissolved chlorine reacts to form hydrochloric acid (HCl) and hypochlorous acid (HClO), a solution that can be used as a disinfectant or bleach: $Cl_2(g) + H_2O(l) \rightarrow HCl(aq) + HClO(aq)$ Bromine has a solubility of 3.41 g/100 g H2O, but it slowly reacts to form hydrogen bromide (HBr) and hypobromous acid (HBrO): $Br_2(g) + H_2O(l) \rightarrow HBr(aq) + HBrO(aq)$ Iodine, however, is minimally soluble in water (0.03 g/100 g H2O at 20 °C) and does not react with it. However, iodine will form an aqueous solution in the presence of iodide ions , such as with the addition of potassium iodide (KI), because the triiodide ion (I3-) is formed.

Halogen elements facts, information, pictures | Encyclopedia.com articles about Halogen elements Iodine (I): 53 Astatine (At): 85 On the periodic table, as displayed in chemistry labs around the world, the number of columns and rows does not vary, since these configurations are the result of specific and interrelated properties among the elements. There are always 18 columns; however, the way in which these are labeled differs somewhat from place to place. Many chemists outside the United States refer to these as 18 different groups of elements; however, within the United States, a somewhat different system is used. In many American versions of the chart, there are only eight groups, sometimes designated with Roman numerals. The 40 transition metals in the center are not designated by group number, nor are the lanthanides and actinides, which are set apart at the bottom of the periodic table. The remaining eight columns are the only ones assigned group numbers. In many ways, this is less useful than the system of 18 group numbers; however, it does have one advantage. ELECTRON CONFIGURATIONS AND BONDING.

In the eight-group system, group number designates the number of valence electrons. The valence electrons, which occupy the highest energy levels of an atom, are the electrons that bond one element to another. These are often referred to as the "outer shell" of an atom, though the actual structure is much more complex. In any case, electron configuration is one of the ways halogens can be defined: all have seven valence electrons. Because the rows in the periodic table indicate increasing energy levels, energy levels rise as one moves up the list of halogens. Fluorine, on row 2, has a valence-shell configuration of 2s22p5; while that of chlorine is 3s23p5. Note that only the energy level changes, but not the electron configuration at the highest energy level. The same goes for bromine (4s24p5), iodine (4s24p5), and astatine (5s25p5). All members of the halogen family have the same valence-shell electron configurations, and thus tend to bond in much the same way.

As we shall see, they are inclined to form bonds more readily than most other substances, and indeed fluorine is the most reactive of all elements. Thus it is ironic that they are "next door" to the Group 8 noble gases, the least reactive among the elements. The reason for this, as discussed in the Chemical Bonding essay, is that most elements bond in such a way that they develop a valence shell of eight electrons; the noble gases are already there, so they do not bond, except in some cases—and then principally with fluorine. Characteristics of the Halogens In terms of the phase of matter in which they are normally found, the halogens are a varied group. Fluorine and chlorine are gases, iodine is a solid, and bromine is one of only two elements that exists at room temperature as a liquid. As for astatine, it is a solid too, but so highly radioactive that it is hard to know much about its properties. Despite these differences, the halogens have much in common, and not just with regard to their seven valence electrons. Indeed, they were identified as a group possessing similar characteristics long before chemists had any way of knowing about electrons, let alone electron configurations.

One of the first things scientists noticed about these five elements is the fact that they tend to form salts. In everyday terminology, "salt" refers only to a few variations on the same thing—table salt, sea salt, and the like. In chemistry, however, the meaning is much broader: a salt is defined as the result of bonding between an acid and a base. Many salts are formed by the bonding of a metal and a nonmetal. The halogens are all nonmetals, and tend to form salts with metals, as in the example of sodium chloride (NaCl), a bond between chlorine, a halogen, and the metal sodium. The result, of course, is what people commonly call "salt." Due to its tendency to form salts, the first of the halogens to be isolated—chlorine, in 1811—was originally named "halogen." This is a combination of the Greek words halos, or salt, and gennan, "to form or generate." In their pure form, halogens are diatomic, meaning that they exist as molecules with two atoms: F2, Cl2, and so on.

When bonding with metals, they form ionic bonds, which are the strongest form of chemical bond. In the process, halogens become negatively charged ions, or anions. These are represented by the symbols F−, Cl−, Br−, and I−, as well as the names fluoride, chloride, bromide, and iodide. All of the halogens are highly reactive, and will combine directly with almost all elements. Due to this high level of reactivity, the halogens are almost never found in pure form; rather, they have to be extracted. Extraction of halogens is doubly problematic, because they are dangerous. Exposure to large quantities can be harmful or fatal, and for this reason halogens have been used as poisons to deter unwanted plants and insects—and, in one of the most horrifying chapters of twentieth century military history, as a weapon in World War I. REAL-LIFE APPLICATIONS Chlorine Chlorine is a highly poisonous gas, greenish-yellow in color, with a sharp smell that induces choking in humans. Yet, it can combine with other elements to form compounds safe for human consumption.

Most notable among these compounds is salt, which has been used as a food preservative since at least 3000 b.c. Salt, of course, occurs in nature. By contrast, the first chlorine compound made by humans was probably hydrochloric acid, created by dissolving hydrogen chloride gas in water. The first scientist to work with hydrochloric acid was Persian physician and alchemist Rhazes (ar-Razi; c. 864-c. 935), one of the most outstanding scientific minds of the medieval period. Alchemists, who in some ways were the precursors of true chemists, believed that base metals such as iron could be turned into gold. Of course this is not possible, but alchemists in about 1200 did at least succeed in dissolving gold using a mixture of hydrochloric and nitric acids known as aqua regia. The first modern scientist to work with chlorine was Swedish chemist Carl W. Scheele (1742-1786), who also discovered a number of other elements and compounds, including barium, manganese, oxygen, ammonia, and glycerin.

However, Scheele, who isolated it in 1774, thought that chlorine was a compound; only in 1811 did English chemist Sir Humphry Davy (1778-1829) identify it as an element. Another chemist had suggested the name "halogen" for the alleged compound, but Davy suggested that it be called chlorine instead, after the Greek word chloros, which indicates a sickly yellow color. USES OF CHLORINE. The dangers involved with chlorine have made it an effective substance to use against stains, plants, animals—and even human beings. Chlorine gas is highly irritating to the mucous membranes of the nose, mouth, and lungs, and it can be detected in air at a concentration of only 3 parts per million (ppm). The concentrations of chlorine used against troops on both sides in World War I (beginning in 1915) was, of course, much higher. Thanks to the use of chlorine gas and other antipersonnel agents, one of the most chilling images to emerge from that conflict was of soldiers succumbing to poisonous gas.

Yet just as it is harmful to humans, chlorine can be harmful to microbes, thus preserving human life. As early as 1801, it had been used in solutions as a disinfectant; in 1831, its use in hospitals made it effective as a weapon against a cholera epidemic that swept across Europe . Another well-known use of chlorine is as a bleaching agent. Until 1785, when chlorine was first put to use as a bleach, the only way to get stains and unwanted colors out of textiles or paper was to expose them to sunlight, not always an effective method. By contrast, chlorine, still used as a bleach today, can be highly effective—a good reason not to use regular old-fashioned bleach on anything other than white clothing. (Since the 1980s, makers of bleaches have developed all-color versions to brighten and take out stains from clothing of other colors.) Calcium hydrocholoride (CaOCl), both a bleaching powder and a disinfectant used in swimming pools, combines both the disinfectant and bleaching properties of chlorine. This and the others discussed here are just some of many, many compounds formed with the highly reactive element chlorine.

Particularly notable—and controversial—are compounds involving chlorine and carbon. CHLORINE AND ORGANIC COMPOUNDS. Chlorine bonds well with organic substances, or those containing carbon. In a number of instances, chlorine becomes part of an organic polymer such as PVC (polyvinyl chloride), used for making synthetic pipe. Chlorine polymers are also applied in making synthetic rubber, or neoprene. Due to its resistance to heat, oxidation, and oils, neoprene is used in a number of automobile parts. The bonding of chlorine with substances containing carbon has become increasingly controversial because of concerns over health and the environment, and in some cases chlorine-carbon compounds have been outlawed. Such was the fate of DDT, a pesticide soluble in fats and oils rather than in water. When it was discovered that DDT was carcinogenic, or cancer-causing, in humans and animals, its use in the United States was outlawed. Other, less well-known, chlorine-related insecticides have likewise been banned due to their potential for harm to human life and the environment. Among these are chlorine-containing materials once used for dry cleaning.

Also notable is the role of chlorine in chlorofluorocarbons (CFCs), which have been used in refrigerants such as Freon, and in propellants for aerosol sprays. CFCs tend to evaporate easily, and concerns over their effect on Earth 's atmosphere have led to the phasing out of their use. Fluorine Fluorine has the distinction of being the most reactive of all the elements, with the highest electronegativity value on the periodic table. Because of this, it proved extremely difficult to isolate. Davy first identified it as an element, but was poisoned while trying unsuccessfully to decompose hydrogen fluoride. Two other chemists were also later poisoned in similar attempts, and one of them died as a result. French chemist Edmond Fremy (1814-1894) very nearly succeeded in isolating fluorine, and though he failed to do so, he inspired his student Henri Moissan (1852-1907) to continue the project.

One of the problems involved in isolating this highly reactive element was the fact that it tends to "attack" any container in which it is placed: most metals, for instance, will burst into flames in the presence of fluorine. Like the others before him, Moissan set about to isolate fluorine from hydrogen fluoride by means of electrolysis—the use of an electric current to cause a chemical reaction—but in doing so, he used a platinum-iridium alloy that resisted attacks by fluorine. In 1906, he received the Nobel Prize for his work, and his technique is still used today in modified form. PROPERTIES AND USES OF FLUORINE. A pale green gas of low density, fluorine can combine with all elements except some of the noble gases. Even water will burn in the presence of this highly reactive substance. Fluorine is also highly toxic, and can cause severe burns on contact, yet it also exists in harmless compounds, primarily in the mineral known as fluorspar, or calcium fluoride. The latter gives off a fluorescent light (fluorescence is the term for a type of light not accompanied by heat), and fluorine was named for the mineral that is one of its principal "hosts".

Beginning in the 1600s, hydrofluoric acid was used for etching glass, and is still used for that purpose today in the manufacture of products such as light bulbs. The oil industry uses it as a catalyst—a substance that speeds along a chemical reaction—to increase the octane number in gasoline. Fluorine is also used in a polymer commonly known as Teflon, which provides a non-stick surface for frying pans and other cooking-related products. Just as chlorine saw service in World War I, fluorine was enlisted in World War II to create a weapon far more terrifying than poison gas: the atomic bomb. Scientists working on the Manhattan Project, the United States' effort to develop the bombs dropped on Japan in 1945, needed large quantities of the uranium-235 isotope. This they obtained in large part by diffusion of the compound uranium hexafluoride, which consists of molecules containing one uranium atom and six fluorine anions. FLUORIDATION OF WATER. Long before World War II, health officials in the United States noticed that communities having high concentration of fluoride in their drinking water tended to suffer a much lower incidence of tooth decay.

In some areas the concentration of fluoride in the water supply was high enough that it stained people's teeth; still, at the turn of the century—an era when dental hygiene as we know it today was still in its infancy—the prevention of tooth decay was an attractive prospect. Perhaps, officials surmised, it would be possible to introduce smaller concentrations of fluoride into community drinking water, with a resulting improvement in overall dental health. After World War II, a number of municipalities around the United States undertook the fluoridation of their water supplies, using concentrations as low as 1 ppm. Within a few years, fluoridation became a hotly debated topic, with proponents pointing to the potential health benefits and opponents arguing from the standpoint of issues not directly involved in science. It was an invasion of personal liberty, they said, for governments to force citizens to drink water which had been supplemented with a foreign substance. During the 1950s, in fact, fluoridation became associated in some circles with Communism—just another manifestation of a government trying to control its citizens. In later years, ironically, antifluoridation efforts became associated with groups on the political left rather than the right.

By then, the argument no longer revolved around the issue of government power; instead the concern was for the health risks involved in introducing a substance lethal in large doses. Fluoride had meanwhile gained application in toothpastes. Colgate took the lead, introducing "stannous fluoride" in 1955. Three years later, the company launched a memorable advertising campaign with commercials in which a little girl showed her mother a "report card" from the dentist and announced "Look, Ma! No cavities!" Within a few years, virtually all brands of toothpaste used fluoride; however, the use of fluoride in drinking water remained controversial. As late as 1993, in fact, the issue of fluoridation remained heated enough to spawn a study by the U.S. National Research Council. The council found some improvement in dental health, but not as large as had been claimed by early proponents of fluoridation. Furthermore, this improvement could be explained by reference to a number of other factors, including fluoride in toothpastes and a generally heightened awareness of dental health among the U.S. populace. CHLOROFLUOROCARBONS.

Another controversial application of fluorine is its use, along with chlorine and carbon, in chlorofluorocarbons. As noted above, CFCs have been used in refrigerants and propellants; another application is as a blowing agent for polyurethane foam. This continued for several decades, but in the 1980s, environmentalists became concerned over depletion of the ozone layer high in Earth's atmosphere. Unlike ordinary oxygen (O2), ozone or O3 is capable of absorbing ultraviolet radiation from the Sun , which would otherwise be harmful to human life. It is believed that CFCs catalyze the conversion of ozone to oxygen, and that this may explain the "ozone hole," which is particularly noticeable over the Antarctic in September and October. As a result, a number of countries signed an agreement in 1996 to eliminate the manufacture of halocarbons, or substances containing halogens and carbon. Manufacturers in countries that signed this agreement, known as the Montreal Protocol, have developed CFC substitutes, most notably hydrochlorofluorocarbons (HCFCs), CFC-like compounds also containing hydrogen atoms. The ozone-layer question is far from settled, however.

Critics argue that in fact the depletion of the ozone layer over Antarctica is a natural occurrence, which may explain why it only occurs at certain times of year. This may also explain why it happens primarily in Antarctica, far from any place where humans have been using CFCs. (Ozone depletion is far less significant in the Arctic, which is much closer to the population centers of the industrialized world.) In any case, natural sources, such as volcano eruptions, continue to add halogen compounds to the atmosphere. Bromine Bromine is a foul-smelling reddish-brown liquid whose name is derived from a Greek word meaning "stink." With a boiling point much lower than that of water—137.84°F (58.8°C)—it readily transforms into a gas. Like other halogens, its vapors are highly irritating to the eyes and throat. It is found primarily in deposits of brine, a solution of salt and water. Among the most significant brine deposits are in Israel's Dead Sea, as well as in Arkansas and Michigan .

Credit for the isolation of bromine is usually given to French chemist Antoine-Jérôme Balard (1802-1876), though in fact German chemist Carl Löwig (1803-1890) actually isolated it first, in 1825. However, Balard, who published his results a year later, provided a much more detailed explanation of bromine's properties. The first use of bromine actually predated both men by several millennia. To make their famous purple dyes, the Phoenicians used murex mollusks, which contained bromine. (Like the names of the halogens, the word "Phoenicians" is derived from Greek—in this case, a word meaning "red" or "purple," which referred to their dyes.) Today bromine is also used in dyes, and other modern uses include applications in pesticides, disinfectants, medicines, and flame retardants. At one time, a compound containing bromine was widely used by the petroleum industry as an additive for gasoline containing lead.

Ethylene dibromide reacts with the lead released by gasoline to form lead bromide (PbBr2), referred to as a "scavenger," because it tends to clean the emissions of lead-containing gasoline. However, leaded gasoline was phased out during the late 1970s and early 1980s; as a result, demand for ethylene dibromide dropped considerably. HALOGEN LAMPS. The name "halogen" is probably familiar to most people because of the term "halogen lamp." Used for automobile headlights, spotlights, and floodlights, the halogen lamp is much more effective than ordinary incandescent light. Incandescent "heat-producing" light was first developed in the 1870s and improved during the early part of the twentieth century with the replacement of carbon by tungsten as the principal material in the filament, the area that is heated. Tungsten proved much more durable than carbon when heated, but it has a number of problems when combined with the gases in an incandescent bulb. As the light bulb continues to burn for a period of time, the tungsten filament begins to thin and will eventually break.

At the same time, tungsten begins to accumulate on the surface of the bulb, dimming its light. However, by adding bromine and other halogens to the bulb's gas filling—thus making a halogen lamp—these problems are alleviated. As tungsten evaporates from the filament, it combines with the halogen to form a gaseous compound that circulates within the bulb. Instead of depositing on the surface of the bulb, the compound remains a gas until it comes into contact with the filament and breaks down. It is then redeposited on the filament, and the halogen gas is free to combine with newly evaporated tungsten. Though a halogen bulb does eventually break down, it lasts much longer than an ordinary incandescent bulb and burns with a much brighter light. Also, because of the decreased tungsten deposits on the surface, it does not begin to dim as it nears the end of its life. Iodine First isolated in 1811 from ashes of seaweed, iodine has a name derived from the Greek word meaning "violet-colored"—a reference to the fact it forms dark purple crystals.

During the 1800s, iodine was obtained commercially from mines in Chile , but during the twentieth century wells of brine in Japan, Oklahoma , and Michigan have proven a better source. Among the best-known properties of iodine is its importance in the human diet. The thyroid gland produces a growth-regulating hormone that contains iodine, and lack of iodine can cause a goiter, a swelling around the neck. Table salt does not naturally contain iodine; however, sodium chloride sold in stores usually contains about 0.01% sodium iodide, added by the manufacturer. Iodine was once used in the development of photography: during the early days of photographic technology, the daguerreotype process used silver plates sensitized with iodine vapors. Iodine compounds are used today in chemical analysis and in synthesis of organic compounds. Astatine Just as fluorine has the distinction of being the most reactive, astatine is the rarest of all the elements. Long after its existence was predicted, chemists still had no luck finding it in nature, and it was only created in 1940 by bombarding bismuth with alpha particles (positively charged helium nuclei).

The newly isolated element was given a Greek name meaning "unstable." Indeed, none of astatine's 20 known isotopes is stable, and the longest-lived has a half-life of only 8.3 hours. This has only added to the difficulties involved in learning about this strange element, and therefore it is difficult to say what applications, if any, astatine may have. The most promising area involves the use of astatine to treat a condition known as hyperthyroidism, related to an overly active thyroid gland. WHERE TO LEARN MORE "The Chemistry of the Halogens." Purdue University Department of Chemistry (Web site). < (May 20, 2001). "Halogens." Chemical Elements.com (Web site). < (May 20, 2001). "Halogens." Corrosion Source (Web site). < (May 20, 2001). "Halogens" (Web site). < (May 20, 2001). "The Halogens" (Web site). <~jfromm/elements/halogen.htm> (May 20, 2001). Knapp, Brian J.

Chlorine, Fluorine, Bromine, and Iodine. Henley-on-Thames, England : Atlantic Europe, 1996. Oxlade, Chris. Elements and Compounds. Chicago : Heinemann Library, 2001. Stwertka, Albert. A Guide to the Elements. New York : Oxford University Press, 1998. "Visual Elements: Group VII—The Halogens" (Web site). < (May 20, 2001). KEY TERMS ANION: The negative ion that results when an atom gains one or more electrons. An anion (pronounced "AN-ie-un") of an element is never called, for instance, the chlorine anion. Rather, an anion involving a single element is named by adding the suffix -ide to the name of the original element—in this case, "chloride." Other rules apply for more complex anions. ATOMIC NUMBER: The number of protons in the nucleus of an atom. Since this number is different for each element, elements are listed on the periodic table of elements in order of atomic number. CHEMICAL SYMBOL: A one-or two-letter abbreviation for the name of an element.

DIATOMIC: A term describing an element that exists as molecules composed of two atoms. All of the halogens arediatomic. ELECTROLYSIS: The use of an electrical current to cause a chemical reaction. HALF-LIFE: The length of time it takes a substance to diminish to one-half its initial amount. HALOGENS: Group 7 of the periodic table of elements, including fluorine, chlorine, bromine, iodine, and astatine. The halogens are diatomic, and tend to form salts; hence their name, which comes from two Greek terms meaning "salt-forming." ION: An atom that has lost or gained one or more electrons, and thus has a net electric charge. IONIC BONDING: A form of chemical bonding that results from attractions between ions with opposite electricalcharges. ISOTOPES: Atoms that have an equal number of protons, and hence are of the same element, but differ in their number of neutrons. This results in a difference ofmass. Isotopes may be either stable or unstable—that is, radioactive. PERIODIC TABLE OF ELEMENTS: A chart that shows the elements arranged in order of atomic number.

Vertical columns within the periodic table indicate groups or "families" of elements with similar chemical characteristics. POLYMER: A large molecule containing many small units that hook together. RADIOACTIVITY: A term describing a phenomenon whereby certain materials are subject to a form of decay brought about by the emission of high-energy particles. "Decay" in this sense does not mean "rot"; instead, radioactive isotopes continue to emit particles, changing into isotopes of other elements, until they becomestable. SALT: A compound formed by the reaction of an acid with a base. Salts are usually formed by the joining of a metal and a nonmetal. VALENCE ELECTRONS: Electrons that occupy the highest energy levels in anatom, and are involved in chemical bonding. The halogens all have seven valence electrons. Cite this article Pick a style below, and copy the text for your bibliography. MLA Chemical activity: The tendency to form chemical compounds. Compound: The combination of two or more elements in a definite mass ratio. Radioactive: The tendency of an element to break down spontaneously into one or more other elements. Synthesized: Prepared by scientists in a laboratory; not a naturally occurring process.

Because it is so reactive, fluorine itself has few uses. One exception is its role as an oxidizing (burning) agent in rocket fuels. The vast majority of fluorine, however, is used to make compounds. One of the most interesting of those compounds is hydrofluoric acid. This compound has been used since the 1600s to etch glass. Perhaps the most familiar application of fluorine compounds is in toothpaste additives. Scientists have discovered that the addition of tiny amounts of fluoride in a person's diet can decrease the number of dental caries (cavities) that develop. Today, many kinds of toothpastes include stannous fluoride to improve a person's dental health. For many years, the most important group of fluorine compounds used commercially were the chlorofluorocarbons (CFCs). The CFCs were developed and used as refrigerants, blowing agents for polyurethane foam, and propellants in spray cans. At one time, more than 700 million kilograms (1.5 billion pounds) of CFCs were produced in a single year. In the 1980s, however, scientists found that CFCs break down in the atmosphere.

The chlorine formed as a result of this breakdown attacks the ozone layer in Earth 's stratosphere (the part of Earth's atmosphere that extends 7 to 31 miles [11 to 50 kilometers] above the surface). The loss of the ozone layer is a serious problem for humans since ozone screens out radiation that causes skin cancer and other damage to plants and animals on Earth. Today, scientists are exploring the use of another class of fluorine compounds—the hydrochlorofluorocarbons, or HCFCs—as replacements for CFCs. Chlorine Chlorine was first prepared in the 1770s by Swedish chemist Carl Wilhelm Scheele (1742–1786), who thought it was a compound. It was later identified as an element by English chemist Humphry Davy (1778–1829). Davy suggested the name of chlorine for the element because of its greenish-yellow color. (The Greek word for "greenish-yellow" is chloros. ) Chlorine occurs most abundantly in sodium chloride, which is obtained from seawater and from underground deposits of rock salt formed from seas that have dried up.

To obtain chlorine, an electrical current is passed through brine, a water solution of sodium chloride. Properties and uses. Chlorine gas is toxic. It attacks the respiratory tract, causing coughing, congestion, and flu-like symptoms. In high doses, it can be fatal. For this reason, chlorine was used as a chemical weapon during World War I (1914–18). Chlorine is also very reactive, although less so than fluorine. It forms compounds with almost every other element. Among the most important of those compounds are sodium chloride (table salt), potassium chloride, hydrochloric acid, and calcium chloride. Chlorine consistently ranks among the top ten chemicals produced in the United States . Some of chlorine's uses depend on its toxic effects. For example, chlorine is now widely used as a disinfectant in municipal water systems, swimming pools, and sewage treatment plants. Many organic (carbon-containing) compounds of chlorine are used as pesticides, herbicides, and fungicides. These compounds kill unwanted insects, weeds, fungi, and other plants and animals. The use of these compounds is often associated with undesirable environmental effects, however.

DDT, for example, is a chlorine-containing compound that was once one of the most popular pesticides ever produced. But its harmful effects on fish, birds, and other animals in the environment eventually led to bans on its use in many industrialized nations. Chlorination The mention of chlorine brings summertime and swimming pools to mind for most people. Chlorine is added to pools and spas to kill bacteria in water that might otherwise cause disease. The process of adding chlorine to a swimming pool is called chlorination. Chlorination is used for other purification purposes also, as in the purification of public water supplies. Chlorination can be done in various ways. In some cases, gaseous chlorine is pumped directly into water. In other cases, a compound containing chlorine, such as sodium or calcium hypochlorite, is added to water. When that compound breaks down, chlorine and other purifying substances are released to the water. The term chlorination applies more generally to any chemical reaction in which chlorine is added to some other substance. For example, chlorine and methane gas can be reacted with each other to form a series of chlorine-containing compounds.

The best known of that series are trichloromethane (also known as chloroform; used as an anesthetic) and tetrachloromethane (also known as carbon tetrachloride; used as a solvent and a refrigerant). Chlorine is also used in the bleaching of paper, pulp, and textiles. The largest single application of the element is in the preparation of a large variety of compounds, including organic chlorides that are the starting point in the manufacture of plastics and other kinds of polymers (chemical compounds that consist of repeating structural units). One of the most important of these polymers is polyvinyl chloride (PVC), from which plastic pipe and many other plastic products are made. Another is neoprene, a synthetic form of rubber that is resistant to the effects of heat, oxidation, and oils. Neoprene is widely used in automobile parts. Bromine Bromine was discovered in 1826 by French chemist Antoine-Jérôme Balard (1802–1876). Balard chose the name bromine from the Greek word for "stink," because of its strong and disagreeable odor. Like chlorine, bromine is obtained from brine.

Chlorine gas is used to convert bromide compounds in brine to elemental bromine. Properties and uses. Bromine is a beautiful reddish-brown liquid that vaporizes (changes to a gas) easily. The vapors are irritating to the eyes and throat. The liquid is highly corrosive and can cause serious burns if spilled on the skin. Bromine is chemically less active than fluorine and chlorine but more active than iodine. Like chlorine, bromine can be used as a disinfectant. In fact, some water treatment systems have converted from chlorination to bromination as a way of purifying water. For many years, one of the most important compounds of bromine was ethylene dibromide, an additive in leaded gasolines. Since leaded gasoline has been removed from the market, this use has declined. The product in which most people are likely to encounter compounds of bromine is in photographic film. Tiny crystals of silver bromide undergo a chemical change when exposed to light. This change is responsible for the image produced when photographic film is used to take a picture. Bromine is also used to make a number of organic products that function as pesticides.

The most popular of these currently is methyl bromide, a fumigant (another word for a substance used to destroy pests). Methyl bromide is used as a spray for potatoes, tomatoes, and other agricultural crops. The halons are a group of organic compounds that contain bromine along with at least one other halogen. The halons are popular as flame retardants. However, scientists have found that, like the CFCs, they appear to cause damage to Earth's ozone layer. For that reason, their use has been largely reduced throughout the world. Iodine Iodine was discovered accidentally in 1811 by French chemist Bernard Courtois (1777–1838). Courtois was burning seaweed to collect potassium nitrate when he noticed that a beautiful violet vapor was produced. When the vapor cooled, it changed to dark, shiny, metallic-like crystals. Humphry Davy later suggested the name iodine for the element from the Greek word iodos, for "violet." As with chlorine and bromine, iodine is obtained from seawater.

It can also be produced from Chile saltpeter (sodium nitrate), in which it occurs as an impurity in the form of sodium iodate (NaIO3). Properties and uses. Iodine vapor is irritating to the eyes and respiratory system. It is highly toxic if ingested. Iodine is the least active of the common halogens (not counting astatine). The human body uses iodine to make thyroxine, an important hormone (chemical messenger) produced by the thyroid gland. (The thyroid is a gland located in the neck that plays an important role in metabolism—a term used to describe processes of energy production and use by the body.) If insufficient amounts of iodine are present in the diet, a person may develop a condition known as goiter, a sometimes noticeable enlargement of the thyroid gland. Once the relationship between iodine and goiter were first discovered, manufacturers of table salt began to add iodine (in the form of sodium iodide) to their product (iodized salt). This practice has largely eliminated the problem of goiter in modern developed nations. Iodine is also used commercially in a variety of products including dyes, specialized soaps, lubricants, photographic film, medicines, and pharmaceuticals.

Astatine Astatine is generally regarded as one of the rarest naturally occurring elements. According to some estimates, no more than 44 milligrams of the element are to be found in Earth's crust. It is hardly surprising, then, that the element was first produced synthetically. In 1940, three physicists at the University of California at Berkeley—D. R. Corson, K. R. Mackenzie, and Emilio Segrè (1905–1989)—made astatine by bombarding the element bismuth with alpha particles in a cyclotron (a particle accelerator or atom-smasher). About 24 isotopes (forms) of astatine exist, all of them radioactive. The most long-lived has a half-life of 8.3 hours, meaning that half of a sample of the element disappears in 8.3 hours. Because it is so rare and has such a short half-life, astatine is one of the most poorly understood of all chemical elements. It has no practical applications at this time.

[See also Element, chemical; Organic chemistry; Periodic table ] Cite this article halogen element | Facts, Definition, Properties, & Uses | Britannica.com Halogen element Alternative Titles: Group 17 element, Group VIIa element Related Topics iodine (I) Halogen element, any of the six nonmetallic elements that constitute Group 17 (Group VIIa) of the periodic table. The halogen elements are fluorine (F), chlorine (Cl), bromine (Br), iodine (I), astatine (At), and tennessine (Ts). They were given the name halogen, from the Greek roots hal- (“salt”) and -gen (“to produce”), because they all produce sodium salts of similar properties, of which sodium chloride—table salt , or halite —is best known. Modern version of the periodic table of the elements. Encyclopædia Britannica, Inc. Because of their great reactivity, the free halogen elements are not found in nature. In combined form, fluorine is the most abundant of the halogens in Earth’s crust.

The percentages of the halogens in the igneous rocks of Earth’s crust are 0.06 fluorine, 0.031 chlorine, 0.00016 bromine, and 0.00003 iodine. Astatine and tennessine do not occur in nature, because they consist of only short-lived radioactive isotopes . The halogen elements show great resemblances to one another in their general chemical behaviour and in the properties of their compounds with other elements. There is, however, a progressive change in properties from fluorine through chlorine, bromine, and iodine to astatine—the difference between two successive elements being most pronounced with fluorine and chlorine. Fluorine is the most reactive of the halogens and, in fact, of all elements, and it has certain other properties that set it apart from the other halogens. Chlorine is the best known of the halogen elements. The free element is widely used as a water-purification agent, and it is employed in a number of chemical processes. Sodium chloride, of course, is one of the most familiar chemical compounds.

Fluorides are known chiefly for their addition to public water supplies to prevent tooth decay , but organic fluorides are also used as refrigerants and lubricants. Iodine is most familiar as an antiseptic , and bromine is used chiefly to prepare bromine compounds that are used in flame retardants and as general pesticides . In the past ethylene dibromide was extensively used as an additive in leaded gasoline . Oxidation cobalt (Co) Probably the most important generalization that can be made about the halogen elements is that they are all oxidizing agents ; i.e., they raise the oxidation state, or oxidation number , of other elements—a property that used to be equated with combination with oxygen but that is now interpreted in terms of transfer of electrons from one atom to another. In oxidizing another element, a halogen is itself reduced; i.e., the oxidation number 0 of the free element is reduced to −1. The halogens can combine with other elements to form compounds known as halides —namely, fluorides , chlorides , bromides , iodides , and astatides.

Many of the halides may be considered to be salts of the respective hydrogen halides , which are colourless gases at room temperature and atmospheric pressure and (except for hydrogen fluoride) form strong acids in aqueous solution. Indeed, the general term salt is derived from rock salt , or table salt (sodium chloride). The tendency of the halogen elements to form saltlike (i.e., highly ionic) compounds increases in the following order: astatine < iodine < bromine < chlorine < fluorine . Fluorides are usually more stable than the corresponding chlorides, bromides, or iodides. (Often astatine is omitted from general discussions of the halogens because less is known about it than about the other elements.) Ionic bond Encyclopædia Britannica, Inc. The oxidizing strength of the halogens increases in the same order—i.e., from astatine to fluorine. Therefore, of the halogen elements, elemental fluorine is prepared with the greatest difficulty and iodine with the least. As a class, the halogen elements are nonmetals, but astatine shows certain properties resembling those of the metals.

Electronic structure Periodic Table of the Elements The chemical behaviour of the halogen elements can be discussed most conveniently in terms of their position in the periodic table of the elements. In the periodic table the halogens make up Group 17 (according to the numbering system adopted by the International Union of Pure and Applied Chemistry), the group immediately preceding the noble gases. The halogen atoms carry seven valence electrons in their outermost electron shell . These seven outermost electrons are in two different kinds of orbitals , designated s (with two electrons) and p (with five). Potentially, a halogen atom could hold one more electron (in a p orbital), which would give the resulting halide ion the same arrangement (configuration) as that of the noble gas next to it in the periodic table. These electron configurations are exceptionally stable. This pronounced tendency of the halogens to acquire an additional electron renders them strong oxidizers. Britannica Stories EU Considers Rules For Robots At room temperature and atmospheric pressure the halogen elements in their free states exist as diatomic molecules . In molecular fluorine (F2) the atoms are held together by a bond made from the union of a p orbital from each atom, with such a bond being classed as a sigma bond .

It should be mentioned that the dissociation energy for fluorine (the energy necessary to break the F−F bond) is over 30 percent smaller than that of chlorine but is similar to that of iodine (I2). The weakness of the F−F single bond compared with chlorine can be ascribed to the small size of fluorine resulting in a decreased overlap of bonding orbitals and an increased repulsion of the nonbonding orbitals. In iodine, however, the p orbitals are more diffuse, which means the bond becomes weaker than in chlorine or bromine. Relative reactivity The great reactivity of fluorine largely stems from the relatively low dissociation energy, a standard measure for bond energies, of the F−F bond (37.7 kilocalories per mole) and its ability to form stable strong bonds with essentially all the other elements. Fluorine (F2) and chlorine (Cl2) are gases at room temperature. Bromine (Br2) is a reddish-brown liquid at room temperature and is—apart from mercury —the only element that is liquid at 20 °C (68 °F) and atmospheric pressure.

Iodine (I2) forms dark violet crystals under these conditions. In the solid state the halogen elements form molecular lattices, and the sublimation energies rise with increasing size of the molecules. Connect with Britannica Facebook Twitter YouTube Instagram Pinterest The energy released in the formation of an ion from a free atom and an electron (brought up from an infinite distance) is called the electron affinity . The electron affinities for the halogen atoms all are high and show only slight differences from one another. It is known, however, that the oxidizing properties (ability to take up an electron by formation of a bond with another atom) increase from astatine to fluorine. This increase can be attributed to the low dissociation energy and the high electron affinity of fluorine combined with the strength of the resulting fluorine-hetero atom bond, resulting in a large heat of reaction. While the fluoride ion exhibits no reducing properties, the iodide ion is a mild reducing agent. Within a molecule in which atoms are held together by a shared electron pair (i.e., by a covalent or nonionic bond), the tendency of an atom to attract the shared electrons may be expressed by an electronegativity value.

According to American chemist Linus Pauling , “Electronegativity is the power of an atom in a molecule to attract electrons to itself.” Fluorine possesses the highest electronegativity of all elements, and there is a decrease in electronegativity within the family of the halogen elements from fluorine through chlorine, bromine, and iodine to astatine. Fluorine replaces any other halide ion from its compounds, as shown in the following equations. Chlorine, however, replaces only bromide, iodide, and astatide ions, and bromine only iodide and astatide ions. Free fluorine, chlorine, bromine, and iodine are expected to replace astatide ions. The halogen elements all form compounds with hydrogen , the hydrogen halides . The energy of the hydrogen-halogen bond increases strongly from iodide to fluoride. Hydrogen fluoride in the crystalline state consists of infinite zigzag chains, as shown in the diagram, in which H represents the hydrogen atoms and (as before) F the fluorine atoms; the solid lines represent covalent bonds between the hydrogen and fluorine atoms within the molecules, and the dotted lines represent secondary bonds, called hydrogen bonds .

The hydrogen bonds between hydrogen fluoride molecules are considerably weaker (7 kilocalories per mole) than those within the molecules (135 kilocalories per mole), yet they are retained to a great extent in the liquid state. Similar hydrogen bonding exists in the other hydrogen halides, but it is considerably weaker. The large difference in hydrogen bonding between hydrogen fluoride and the other hydrogen halides accounts for the relatively high melting and boiling points of hydrogen fluoride as compared with those of hydrogen chloride and the other hydrogen halides. The hydrogen-halogen bond energies also decrease considerably in going from hydrogen fluoride to hydrogen iodide. Britannica Lists & Quizzes Editor Picks: Exploring 10 Types of Basketball Movies The ionization energies of the halogens are generally high, but they fall markedly with increasing atomic number . Fluorine is the only halogen that does not form compounds with positive oxidation states—i.e., states in which it has lost, rather than gained, electrons. This property is related to fluorine’s having the highest electronegativity of all elements; i.e., it does not give up its electrons to other elements. All halogens possess the oxidation state 0 in their diatomic forms.

Fluorine exhibits the oxidation states of −1 (F− ion) and +1 (hypofluorous acid). The principal oxidation states of chlorine, bromine, and iodine are −1, +1, +3, +5, and +7. The oxyacids are compounds in which halogen atoms are joined to oxygen atoms. The oxyacids are all powerful oxidizing agents, which can be reduced to the corresponding hydrogen halides—the oxidation numbers changing from positive to −1 in the process. The oxidizing strength of the oxyanions increases with increasing oxidation number of the halogen atom. All the molecules and ions in which halogen atoms possess four valence electron pairs are tetrahedral, as, for example, in the perchlorate ion (ClO4)−. Those employing five valence electron pairs, such as chlorine trifluoride (ClF3), have structures derived from a trigonal bipyramidal arrangement of electron pairs. However, since electron lone pairs (i.e., electron pairs that do not bond atoms together) are not located by techniques that analyze structure, only the positions of the fluorine atoms (attached to bonding pairs) are seen.

Thus, ClF3 has a T shape resulting from the placement of fluorine atoms at both axial and at one equatorial position of the trigonal bipyramid, with lone electron pairs in the remaining two equatorial positions. Molecules with six valence electron pairs have structures derived from octahedral geometry for the electron pairs; e.g., iodine pentafluoride (IF5) has a square pyramidal structure resulting from the bonding of fluorine atoms by five of the six octahedral electron pairs. The unique binary compound iodine heptafluoride (IF7) has a pentagonal bipyramidal arrangement of fluorine atoms. Trending Topics

Question: Chemistry - which of the halogens has the lowest atomic number?
Answer: FLOURINE