Halogens Chemical properties of halogens Inter halogen compounds
Physical properties of halogens Occurence and extraction of the halogens Chlor–alkali industry
Uses of halogens Noble gases Compounds of noble gases
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Halogens and Noble Gases:
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Gases form very important components of many chemical reactions. In the periodic table of elements gases are mainly found in Group VII and Group VIII. The Group VII gases are known as halogen gases and the Group VIII gases are known as the noble gases. Let us look at their chemical properties a little more closely.
Halogens:
The halogen family in the Group VII of the modern periodic table has elements Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I) and Astatine (At).
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Halogens exist in the elemental form as diatomic molecules. At room temperature fluorine and chlorine exist as ‘gas‘; bromine is liquid and Iodine is a solid. The final member of the group, namely astatine, is intensely radioactive. Its most stable isotope has a half–life of only 8.3 hours. Fluorine is the most electronegative element found in nature – it loves to attract an electron and fill its shell.
The term ‘halogen’ is derived from Greek origin hals(sea) and halos(salt). It means ‘salt producing’ since these elements are found in the salts present in seawater.
The natural salt that we obtain from evaporation of seawater, is known as common salt or table salt. It is predominantly sodium chloride. It has small amounts of other chlorides and iodides (such as NaI, MgI2, MgCl2 etc.).
Table salt (NaCl) in its pure form does not absorb moisture. A substance is said to be deliquescent if it absorbs moisture from the air and dissolves in the water. NaCl is not a deliquescent compound. Salt made from seawater has a little bit of magnesium chloride (MgCl2) – which is deliquescent. It absorbs moisture from air and then dissolves in it. This makes the table salt moist, especially in the rainy season.
Iodized salt is a modern day commercially available salt – which is NaCl mixed with NaI. Iodine is necessary, in trace amounts, for our thyroid glands to function properly; it regulates our growth and body temperature.
Chlorine gas is used as a disinfectant and bleaching agent. In swimming pools mild chloric acid HOCl is mixed in water – the nascent oxygen released is useful for killing bacteria in the water. This is the reason why swimming pools always smell of chlorine!!.
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Chemical properties of halogens:
Halogens are a group of nonmetals that appear in the periodic table just before the noble gases.
The chemical behavior of the halogens is the result of the seven electrons in their outer shell. There are two electrons in the s sub–shell and five in the p sub–shell. Thus the addition of only one further electron by either ionic or covalent bonding will confer a noble gas configuration to the halogen atom.
F – 2 electrons in 1s, 2 in 2s, and 5 in 2p
Cl – 2 electrons in 1s, 2 in 2s, 6 in 2p, 2 in 3s, 5 in 3p
Br – 2 electrons in 1s, 2 in 2s, 6 in 2p, 2 in 3s, 6 in 3p, 2 in 4s, 10 in 3d, 5 in 4p
They are strong oxidizing agents as they borrow electrons. The most common oxidation state for the group VII elements is – 1, although other oxidation states do exist. They attain noble gas configuration either by gaining one electron to form compounds like NaCl, KI, LiF or by sharing to form a single covalent bond ( HCl, HBr, CCl4, BF3).
As the atomic number increases the group VII elements become less reactive. They also become less volatile and darker in color. Fluorine is a pale yellow gas at room temperature, whilst chlorine is a dark greenish yellow. Bromine is a dark red liquid giving off a dense red vapor, whilst iodine is a shiny, greyish black crystalline solid, which can be sublimed to a purple vapor.
Fluorine is a small atom, hence its nine electrons are firmly bound. It is difficult to remove electrons. Therefore fluorine has the highest ionization energy in the group. The ionization energy of iodine is lowest. Similarly, the electro–negativity of fluorine is the highest and for iodine it is the lowest.
Fluorine can replaces other halogens from halide compounds either in solution phase or even in dry state. Chlorine is the second best oxidizing agent of the group and oxidizes bromide and iodide ions but not the fluoride ion. Bromine is an oxidizing agent like chlorine, although not so powerful, it oxidizes iodine ion but not chloride ion. Iodine is a weaker oxidizing agent than either bromine or chlorine. It does not oxidize either bromide or chloride ion. In general any halogen of low atomic number will oxidize halide ions of higher atomic number.
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(Here A, B, D are any metal or non–metal ions)
Fluorine is the most electronegative and always has oxidation numberof –1. Other halogens show +ve oxidation states with more electronegative elements. Most of these compounds contain oxygen which has electronegativity between fluorine and chlorine.
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Four general points can be made on these oxidation states. They are
- All the halogens form compounds in which their oxidation state is –1.
- Fluorine exists only in the –1 state in its compounds.
- Both chlorine and iodine form compounds in which their oxidation state is +7.
- Bromoine‘s highest oxidation state is +5.
The halogens react strongly with all the more electropositive elements. When they react with metals, they remove some or all of the outer metal electrons and themselves become reduced to form negative halide ions. This is seen particularly dearly in the reactions with the alkali and alkaline–earth metals.
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The vigour of the reaction between a particular halogen and a metal is a function of the position of the metal in the activity series (in other words, how electropositive it is) and the particular halogen (fluorine always reacts more vigorously than bromine, for example).
Fluorine reacts readily with every metal, even those such as gold and platinum which are usually regarded as being very unreactive. Iodine will also react with these elements, but only very slowly even at very high temperatures.
Metals with low ionization energies form highly ionic halides with high melting and boiling point for example NaCl, KBr.
Metals with high ionization energies form covalent halides with low melting and boiling points for example SnCl4, SnCl2; PbCl4, PbCl2; SbCl2, SbCl2.
If a metal exhibits more than one oxidation state, the halides of the metal in the highest oxidation state will generally be the most covalent. Thus SnCl4 is more covalent than SnCl2.
Fluorine, being the most electro – negative, has the highest ionic character among metal halogen bonds. The ionic character of the M–X bond decreases as
M – F > M – Cl > M – Br > M – I
For example, AlF3 is ionic while AlI3 is essentially covalent.
Most metal halides are soluble in water. The solubility increases as we go from the fluoride to the iodide. For example, CaF2 is only sparingly soluble in water, but CaCl2 is soluble in water.
In reactions with the non–metals, the halogens usually achieve a noble gas configuration through covalent bonding. Fluorine forms only one covalent bond, but the other halogens can form three, five and seven bonds.
Nowhere is the trend in the reactivity of the group VII elements seen more clearly than in their reactions with hydrogen.
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The above Table shows that under all normal conditions, a mixture of fluorine and hydrogen reacts explosively to form hydrogen fluoride, whilst iodine will only react slowly and incompletely with hydrogen even at 3000C and with a platinum catalyst. Whilst the differences in the reactivities of the halogens with the non–metals are demonstrated particularly dramatically in their reactions with hydrogen, the pattern is similar for all the reactions with non–metals.
The hydrogen halides are acids. The boiling points of the halides increase on passing down the group, with the exception of hydrogen fluoride. The very high boiling point of this compound is due to the relatively strong hydrogen bonding, which results from the extremely electronegative character of the fluoride ion.
Fluorine reacts with alkalis in a different way to chlorine, bromine and iodine, always forming a mixture of fluoride ions and oxygen difluoride:
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When solid ionic halides are reacted with strong acids various products may be obtained, depending upon the oxidizing power of the acid. With concentrated phosphoric (V) acid, the hydrogen halides are evolved as gases, while with concentrated sulphuric acid the products depend upon the halide.
| Acid Added |
fluoride | chloride | bromide | Iodide |
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Concentrated
Phosphoric(V)
| HF(g) evolved | HCl(g) evolved | HBr(g) evolved | HI(g) evolved | Concentrated Sulphuric acid
| HF(g)
evolved | HCl(g) evolved | HBr(g) evolved
plus some Br 2(aq) | HI(g) evolved plus I2(aq)
produced |
The reaction of halides with acids
Fluorine is such a powerful oxidizing agent that it reacts with water, producing oxygen and hydrogen fluoride:
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HF is a very dangerous fuming acid.
Hydrogen halides can be prepared by the action of water on phosphorus trihalide.
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Hydrogen fluoride and Hydrogen chloride are prepared by treating soluble metallic chlorides and fluorides (NaCl or NaF), respectively, with concentrated sulphuric acid [H2SO4]. HBr and HI cannot be prepared by this method because sulphuric acid oxidizes the Br– or I– ions to free Halogens.
A convenient method to obtain hydrogen iodide is to treat a suspension of Iodine in water with H2S gas.
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In gas phase all the hydrogen halides HX are covalent diatomic molecules. HF exists with hydrogen bonding and shows anomalous behavior.
Except HF all other hydrogen halides ionize in aqueous solutions and act as strong acids. HF is only partially iodized and so is the weakest acid.
Halogens also form oxides and oxyacids. Some of the examples are listed below.
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Halic acids [HXO3–] are powerful oxidizing agents. e.g .Chlorous acid (HClO2), Periodic acid etc.
Inter halogen compounds:
Halogens react with each other, under suitable conditions of concentration, temperature and pressure, to form a series of binary compounds of the formula XX'n ,where X and X' are different halogen atoms and n is 1,3,5 or 7.
With n=1, all the interhalogen compounds (XX') except IF are known. Other compounds are ClF3, BrF3, ICl3, ClF5, BrF35, IF5 and IF7. All inter halogen compounds are strong oxidizing agents.
The bonds in the interhalogen compounds are essentially covalent because of the small electro–negativity difference; the atom having larger electro – negativity acquires a negative charge. The molecular structures of these compounds are very interesting. ClF3, BrF3 and ICl3 are all T shaped molecules.
Interhalogen compounds are generally more reactive than the halogens. Fluorine has recently been found to react with Xenon, one of the noble gases, under various conditions to form the crystalline solids XeF2, XeF4 and XeF6, which are reasonably stable.
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Physical properties of halogens:
Here is a table listing some of the physical properties of the group VII elements.
Patterns in the properties of the halogens
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The density is measured for the liquid at its boiling point. The atomic radius increases down the group as the number of electron shells increases. As we would expect from this, the electron affinities decrease from chlorine to bromine to iodine. However, at the top of the group, the electron affinity of fluorine is close to that of bromine. On the face of it this is surprising, since we should normally expect an extra electron added to a small atom to be attracted strongly by the positive nucleus. The explanation is that, the electrons in the fluorine atom are close together, and the repulsive forces experienced by an extra electron added to the atom make the formation of the F– ion less favourable than expected.
The decrease in volatility of the elements down the group from gaseous fluorine and chlorine through liquid bromine to the solid iodine is the result of increasingly strong Van der Waals forces between the molecules as the relative molecular mass increases. These increased intermolecular forces also account for the observed increases in melting points, boiling points and enthalpy changes of vaporisation.
Because the halogens form simple, non–polar molecules, they dissolve readily in non–polar organic solvents such as tetrachloromethane. However, chlorine, bromine and iodine are sparingly soluble in water too, although chlorine also undergoes a very slow reaction with the water forming initially chloric(I) acid, HClO(aq), which decomposes to give oxygen and hydrochloric acid:
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Sunlight and catalysts such as platinum speed up the second reaction. It is chloric acid that gives the familiar 'swimming pool' smell to water in which chlorine is dissolved.
The physical and chemical properties of halogens are summarized in the figure below.
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Occurence and extraction of the halogens
The halogens are so reactive that they are not found free in nature, but combined with other elements. Indeed fluorine, the most electronegative element, will combine with almost every other element, including some of the noble gases.
Chlorine too is very electronegative and reacts directly with all other elements except carbon, nitrogen, oxygen and the noble gases. Chlorine also forms compounds with both carbon and oxygen, but not by the direct combination of the elements.
Chlorine is by far the most common of the halogens in nature, being found in the form of sodium chloride in seawater and rock salt, as well as playing a part in all living organisms. Every kilogram of seawater contains around 30 g of sodium chloride. Fluorine is the next most abundant halogen, usually occurring as fluorspar (fluorite), CaF2 and cryolite, Na3AlF6. These deposits are generally quite thin and are rarely economically workable. Sometimes these fluoride containing ores occur as semiprecious minerals, which are mined, polished and used for their appearance rather than the fluoride they contain.
Bromine and iodine are much rarer elements. About 70 parts per milliliter of seawater are bromides which, surprisingly, can be extracted economically. Iodides are present as 0.05 p.p.m. in seawater, although certain seaweeds are capable of concentrating this greatly. The main source of iodine is sodium iodate (V), NaIO3. These occur as impurities in the nitrate deposits (Saltpeter).
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Silver components of bromine and iodine are light sensitive, and are used in the production of photographic filmand also for X–ray pictures.
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Iodine crystals are black and shiny and produce a purple vapour. This vapour kills bacteria and, can be used to prevent infection of wounds and also to make drinking water safe.
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Extracting the halogen from their sources leads to difficulties due to their powerful oxidising tendencies. They are usually obtained by oxidation of the halide ion, but in the case of fluoride there is no oxidising agent strong enough to be used in the extraction process and so the oxidation is performed by electrolysis. The electrolysis is carried out using potassium fluoride dissolved in liquid anhydrous hydrogen fluoride, as the fluorine produced would react with water. A graphite anode and steel cathode is used.
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Twenty–nine million tons or more of chlorine are used throughout the world each year, and for the large–scale extraction needed to provide this quantity of the element a very efficient manufacturing method is needed. Chlorine is used as a cheap industrial oxidant in the manufacture of bromine, as bleach and a germicide, but more importantly it is vital for the manufacture of many everyday materials.
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Chlorine compounds are also used as disinfectants. Bleach (sodium chlorate) is a well known example around the home, and swimming pools and water companies frequently rely on chlorine–based products to keep the water free of bacteria. The halogens are reactive non–metals rarely used as the elements However, their compounds are widely used in a variety of solutions.
About 70% of the industrial use of chlorine goes into the manufature of products such as the plastic PVC (polyvinyl chloride).
Chlor – alkali industry:
The production of chlorine by electrolysis of brine (a concentrated solution of rock salt in water), and the associated production of sodium hydroxide and hydrogen, is the basis of the massive chlor – alkali industry. A flowing mercury cathode is used, as shown in the figure below, and the sodium produced dissolves in the mercury to form an amalgam. After extracting the sodium, the mercury is recycled. The use of mercury in this way means that all the products of the process are useful. It also means that no energy input is required to produce the liquid metal for the electrode, as mercury is a liquid at room temperature.
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The flowing mercury cell is the source of chlorine sodium hydroxide all resulting from the electrolysis of brine.
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During the electrolysis, chlorine is produced and liberated at the graphite anodes, and sodium is produced and it gets dissolved in the mercury cathode.
At the anode (+):
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At the cathode ( – ):
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The chlorine is collected and pressurised for storage, whilst the sodium–mercury amalgam passes on into a second ‘soda cell’. Here the sodium reacts with water to form sodium hydroxide solution (caustic soda) and hydrogen:
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Thus three products of immense use to the chemical industries are made, and the mercury cathode is ready to be used again.
Similarly, fluorine is also obtained from KHF2 (fused) using copper tube and graphite electrodes.
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Bromine is obtained from seawater by oxidation of bromides.
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The bromine is absorbed in Na2CO3 solution giving a mixture of NaBr and NaBrO3, which when acidified and distilled gives bromine.
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A commercial source of free iodine is the reduction of sodium iodate by means of sodium hydrogen sulphite:–
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Uses of halogens:
The following highlight the uses of halogens.
- Fluoride ion, used widely in fluorination of water, under controlled conditions, inhibits tooth decay.
- Covalently bound fluorine–carbon compounds are known as fluorocarbons. These substances are very inert and strong oxidizing agents.
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- Teflon, the most resistant material important in industry is a polymer of fluorine ( CF2–CF2) n – polytetrafluoroethylene. It has very good thermo stability and is therefore used as a coating material.
- Some fluorinated derivatives of methane are used as refrigerants (commercially known as chlorofluorcarbons, freons).
- Chlorine is a very important industrial chemical. Its major use is as a bleaching agent, especially for wood pulp and paper, and also for bleaching cotton cloth. It is used in the preparation of bleaching powder and in the disinfecting water supplies. Solutions of NaOCl are used as household laundry bleaches.
- Potassium chlorate [KClO3] is used as the oxidizing agent in matchsticks and fire works.
- The most important commercial use of bromine is as silver bromide (AgBr), one of the components of photographic emulsion. Exposure to light causes silver bromide to breakdown into free silver resulting in the dark areas of a negative photograph.
- Bromine is also used in the production of ethylene di bromide C2H4Br2 that is added to petroleum containing the antiknock additive, tetraethyl lead. In the process of burning petroleum, lead is produced and deposited on the engine. Ethylene di bromide converts lead into lead bromide, which is eliminated in the exhaust.
- The most important use of iodine is as an antiseptic. A solution of iodine in ethyl alcohol, known as ‘Tincture Iodine‘ is marketed for home use. The iodine in the tincture oxidizes any germs in the open wound and prevents infection.
- Iodine and bromine are used as gas in fluorescent lamps. Iodine combines with tungsten so that the filament lasts longer. When bromine gas is used, the light from a fluorescent lamp looks a bit reddish.
- Halogen compounds are used to make weed killers, pesticides, etc.
Noble gases:
We are all familiar with colorful advertisements with bright lights. These lights are also called Neon lights because of the gas used in them. Neon emits light when electricity excites the atomic electrons. Neon belongs to a family of gases known as noble gases or inert gases, which means unreactive. They belong to the Group VIII of the periodic table starting with Helium He, Neon Ne, Argon Ar, Krypton Kr, Xenon Xe and Radon Rn.
If you see the electrons in the atomic orbit in noble gases, they are completely filled. This family of Noble gases has been put as group VIII [or zero] , at the end of the periodic table because they were not known when the periodic table was first thought about. Nobody knew they were there because they were so unreactive. The bond between the atoms is so weak that at ordinary temperatures they exist as gases.
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All the members of this family are listed below along with some of their characteristics.
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| Name | Symbol | Atomic number | Electronic configuration |
| Helium | He | 2 | 2. |
| Neon | Ne | 10 | 2,8. |
| Aragon | Ar | 18 | 2,8,8. |
| Krypton | Kr |
36 | 2,8,18,8 |
| Xenon | Xe | 54 | 2,8,18,18,8 |
| Radon | Rn | 86 | 2,8,18,32,18,8
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The last orbital in all is full or completed. There is no unpaired electron. So the valency is zero.
Helium has two and other gases have eight electrons in their outermost shell. So, these shells are full. This makes them very stable. They do not react with other atoms, so they exist as single atoms or they are monoatomic gases.
This group has very interesting 18th century history. Cavedish, while passing an electric discharge in a mixture of air and oxygen to prepare oxides of nitrogen and then removing them and oxygen completely, discovered that there always remained a gaseous residue [1/120 of the original volume of air]. This he thought might be a new element. After a lapse of a few years, during a total eclipse of the sun on 18th August 1868, spectroscopic examination of the sun‘s chromosphere indicated a new atomic spectral line, not observed from any of the known elements. This suggested the presence of a new element in the sun and the name Helium from Greek word ‘helios’ (meaning sun) was given to it. Ramsay in 1895 while making a spectroscopic examination of a gas obtained from the mineral Clevite, observed helium lines. Thus its presence on the earth was also established. During the same year another member named "Argon" (the lazy one) was discovered and named so because of its inert character. In the next four years other gases namely Krypton, Xenon and Neon were also isolated.
This family was also called rare gases because they occur in relatively small amounts in the atmosphere. Helium is also found as the occluded gas in some of the radioactive minerals and can be released on heating. Others are obtained by fractional distillation of liquid air.
Properties of noble gases:
All noble gases are very similar in their physical and chemical properties.
They are all colorless, tasteless and odorless mono atomic gases.
All have valence shells, which are closed octets. This accounts for their inert nature. The elements form low boiling liquids, the boiling points decreasing rapidly with the decrease in molecular mass.
Helium has the lowest boiling point and melting point. It crystallizes in a hexagonal closed packed (FCC) structure. Liquid Helium has some peculiar properties not possessed by any other substance known. Thus liquid helium He(I) {at 4.12 K] on further cooling [to 2.178K] changes to another allotrope of Helium, called Helium(II), with negligible viscosity. Such liquids with very low viscosity are known as Superfluids and there is effectively no resistance to their flow. Therefore, when kept in a container Helium even goes uphill and flows out of the vessel! Clever engineering tricks, such as a constriction, etc., have to be made to the dewar holding liquid helium, so that this loss is prevented.
The solubility of noble gases in water is much more than that of oxygen or nitrogen. It increases with the increase in atomic mass of the element that is noble gas with larger atoms is more soluble than a gas with smaller atoms.
Physical properties of group 18 elements
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Uses of noble gases:
The noble gases are used in following applications:
- For welding and cutting of metals where high temperatures are required
- For maintaining inert atmosphere in metallurgical and other processes
- As filler gases in electronic and lighting industry. For example Ar is used in normal incandescent bulbs so that the evapouration of the tungsten filament is reduced.
- In cryogenics (low temp science)
- Helium being light and noninflammable is used to fill weather balloons and airships.
- Radon is in the treatment of cancer.
Compounds of noble gases:
The preparation of stable noble gas compounds is a landmark development in Chemistry. Noble gases, which were called ‘inert‘, lost their inertness in 1962, when Neil Bartlett published the first valid account of a true chemical reaction between Xe and platinum hexafluoride (PtF6). This success was based on the argument that since oxygen react with PtF6 to form O2+ PtF6– complex, and since the ionization potentials of oxygen and xenon are the same, this reaction should occur with Xe also. This work was followed by synthesis of a variety of xenon compounds and several radon (Rn) and krypton (Kr) compounds. No compound of He and Ar have yet been prepared. Xe forms three types of stable compounds – fluorides, oxyfluorides and oxides called xenets and perxenets. Recently it has been shown that Xe also forms xenon dichloride (XeCl2).
Earlier ‘clathrates‘ were the only compounds in which noble gas molecules were found. But these are not true compounds because noble gases were only trapped in the cavities in crystalline lattices of organic or inorganic compounds. No true bonds, either ionic or covalent, were formed. These were Van Waal‘s bonds.
General formula for clathrates may be written as X(2N)12H2O, where N = noble gas such as Ar, Xe, Cr, and X = CH3COOH, CH2Cl2, CHCl3, or CCl4. In such compound the water molecule provides the host lattice, the other compounds, which form host. The organic or the water molecule, forms a cage that traps the noble gas atoms.
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Halogen compounds have also been used over the last few decades as weed killers and pesticides. A variety of complex chlorinated organic molecules have been developed as effective pesticides. Some of these were targeted at the insects and other invertebrates that attack and destroy crops grown for food around the world. Others were aimed at the vectors of tropical diseases, such as the anopheles mosquito that carries malaria. One of the first pesticides to be used in this way was DDT, dichlorodiphenyltrichloroethane. DDT removed the threat of malaria for thousands of millions of people and saved millions of lives throughout the world, and also helped stop the spread of typhus and yellow fever.
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The female Anopheles mosquito spreads the malaria parasite when she bites. Halogen containg compounds help control the mosquitos and save millions from debilitating disease and death.
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Whilst DDT and other early halogen–containing compounds were very effective at their job, it was subsequently realized that they do not break down naturally but remain in the environment. The level of these toxic chemicals used originally was not high – it was sufficient to kill the pests but not a threat to larger organisms.
However, because they are so stable, they have accumulated in the tissues of animals within food chains and have severely reduced the numbers of some carnivorous birds. This is illustrated in the diagram. These pesticides have also made their way into the human food chain. As many of the organisms they were used against have also developed a degree of resistance, the use of organohalogen pesticides is now very carefully regulated.
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DDT used on mosquito larvae passed from prey to predator, reaching higher concentrations at each level. Although it killed some herons, its most important effect was on their reproduction. The birds become much less fertile and chicks were less likely to survive. Halogenated compounds can be very effective in pest contol, but their persistence in biological systems may result in unforeseen damage to other organisms within food chain.
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Halogens Chemical properties of halogens Inter halogen compounds
Physical properties of halogens Occurence and extraction of the halogens Chlor–alkali industry
Uses of halogens Noble gases Compounds of noble gases |
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