The materials to prepare students for practical lessons of inorganic chemistry

The materials to prepare students for practical lessons of inorganic chemistry

LESSON № 21.

Theme: 1. р-elements of the VІІ group. Halogens. 2. р-elements of the VІІІ group (noble gases).

Halogens

The halogens or halogen elements (/ˈhælɵdʒɨn/) are a group in the periodic table consisting of five chemically related elements, fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The artificially created element 117 (ununseptium) may also be a halogen. In the modern IUPAC nomenclature, this group is known as group 17.

The group of halogens is the only periodic table group that contains elements in all three familiar states of matter at standard temperature and pressure. All of the halogens form acids when bonded to hydrogen. Most halogens are typically produced from minerals of salts. The middle halogens, that is, chlorine, bromine and iodine, are often used as disinfectants. The halogens are also all toxic.

History

The fluorine mineral fluorospar was known as early as 1529. Early chemists realized that fluorine compounds contain an undiscovered element, but were unable to isolate it. In 1869, George Gore, an English chemist, ran a current of electricity through hydrofluoric acid and discovered fluorine, but he was unable to prove his results at the time. In 1886, Henri Moissan, a chemist in Paris, performed electrolysis on potassium bifluoride dissolved in waterless hydrofluoric acid, and successfully produced fluorine.

Hydrochloric acid was known to alchemists and early chemists. However, elemental chlorine was not produced until 1774, when Carl Wilhelm Scheele heated hydrochloric acid with manganese dioxide. Scheele called the element “dephlogisticated muriatic acid”, which is how chlorine was known for 33 years. In 1807, Humphrey Davy investigated chlorine and discovered that it is an actual element. Chlorine was used as a poison gas during World War I.

Bromine was discovered in the 1820s by Antoine-Jérôme Balard. Balard discovered bromine by passing chlorine gas through a sample of brine. He originally proposed the name muride for the new element, but the French Academy changed the element’s name to bromine.

Iodine was discovered by Bernard Courtois, who was using seaweed ash as part of a process for saltpeter manufacture. Courtois typically boiled the seaweed ash with water to generate potassium chloride. However, in 1811, Courtois added sulfuric acid to his process, and found that his process produced purple fumes that condensed into black crystals of iodine. Suspecting that he had discovered a new element, Courtois sent his sample to other chemists for investigation. Iodine was proven to be a new element by Joseph Gay-Lussac.

In 1931, Fred Allison claimed to have discovered element 85 with a magneto-optical machine, and named the element Alabamine, but was mistaken. In 1937, Jajendralal De claimed to have discovered element 85 in minerals, and called the element dakine, but he was also mistaken. An attempt at discovering element 85 in 1939 by Horia Hulublei and Yvette Cauchois via spectroscopy was also unsuccessful, as was an attempt in the same year by Walter Minder, who discovered an iodine-like element resulting from beta decay of radium. Element 85 was produced successfully in 1940 by Dale R. Corson, K.R. Mackenzie, and Emilio G. Segrè, who bombarded bismuth with alpha particles.

Etymology

In 1842, the Swedish chemist Baron Jöns Jakob Berzelius proposed the term “halogen” – ἅλς (háls), “salt” or “sea”, and γεν- (gen-), from γίγνομαι (gígnomai), “come to be” – for the four elements (fluorine, chlorine, bromine, and iodine) that produce a sea-salt-like substance when they form a compound with a metal. The word “halogen” had actually first been proposed in 1811 by Johann Salomo Christoph Schweigger as a name for the newly discovered element chlorine, but Davy’s proposed term for this element eventually won out, and Schweigger’s term was kept at Berzelius’ suggestion as the term for the element group that contains chlorine.

Fluorine’s name comes from the Latin word fluere, meaning “to flow”. Chlorine’s name comes from the Greek word chloros, meaning “greenish-yellow”. Bromine’s name comes from the Greek word bromos, meaning “stench”. Iodine’s name comes from the Greek word iodes, meaning “violet”. Astatine’s name comes from the Greek word astatos, meaning “unstable”.

 

The halogens include fluorine, chlorine, bromine, and iodine. Astatine is also in the group, but is radioactive and will not be considered here. A summary of atomic properties of the halogens is given in the following table. The free elemental halogens all consist of diatomic molecules X₂, where X may be fluorine, chlorine, bromine, or iodine (recall the microscopic picture of bromine). They are strong oxidizing agents and are readily reduced to the X^– ions, and so the halogens form numerous ionic compounds. Fluorine, the most electronegative element, has no positive oxidation states, but the other halogens commonly exhibit +1, +3, +5, and +7 states. Most compounds containing halogens in positive oxidation states are good oxidizing agents, however, reflecting the strong tendency of these elements to gain electrons.

Properties of the Group VIIA Elements

 

There is some variation among their physical properties and appearance. Fluorine and chlorine are both gases at room temperature, the former very pale yellow, and the latter yellow-green in color. Bromine is a red-brown liquid which vaporizes rather easily. Iodine forms shiny dark crystals and, when heated, sublimes (changes directly from solid to gas) to a beautiful violet vapor. All the gases produce a choking sensation when inhaled. Chlorine was used to poison soldiers on European battlefields in 1915 to 1918. Halogens are put to more humane uses such as to disinfect public water supplies by means of chlorination and to treat minor cuts by using an alcohol solution (tincture) of iodine. These applications depend on the ability of the halogens to destroy microorganisms which are harmful to humans.

HALOGEN BONDING (XB) is the non-covalent interaction that occurs between a halogen atom (Lewis acid) and a Lewis base. Although halogens are involved in other types of bonding (e.g. covalent), halogen bonding specifically refers to when the halogen acts as an electrophilic species.

Comparison between hydrogen and halogen bonding:

Hydrogen bonding:

Halogen bonding:

In both cases, D (donor) is the atom, group, or molecule that is electron rich and donates them to the electron poor species (H or X). H is the hydrogen atom involved in HB, and X is the halogen atom involved in XB. A (acceptor) is the electron poor species withdrawing the electron density from H or X, accordingly. H-A and X-A, when both atoms are considered together, are called hydrogen/halogen bond donors, accordingly, and D is HB/XB acceptor. An interesting difference between HB and XB is since halogen atoms are Lewis bases, a halogen atom can both donate and accept in a halogen bond.

A parallel relationship can easily be drawn between halogen bonding and hydrogen bonding (HB). In both types of bonding, an electron donor/electron acceptor relationship exists. The difference between the two is what species can act as the electron donor/electron acceptor. In hydrogen bonding, a hydrogen atom acts as the electron acceptor and forms a non-covalent interaction by accepting electron density from an electron rich site (electron donor). In halogen bonding, a halogen atom is the electron acceptor. Simultaneously, the normal covalent bond between H or X and A weakens, so the electron density on H or X appears to be reduced. Electron density transfers results in a penetration of the van der Waals volumes.

Halogens participating in halogen bonding include: iodine (I), bromine (Br), chlorine (Cl), and sometimes fluorine (F). All four halogens are capable of acting as XB donors (as proven through theoretical and experimental data) and follow the general trend: F < Cl < Br < I, with iodine normally forming the strongest interactions.

Dihalogens (I2, Br2, etc.) tend to form strong halogen bonds. The strength and effectiveness of chlorine and fluorine in XB formation depend on the nature of the XB donor. If the halogen is bonded to an electronegative (electron withdrawing) moiety, it is more likely to form stronger halogen bonds.

For example, iodoperfluoroalkanes are well-designed for XB crystal engineering. In addition, this is also why F2 can act as a strong XB donor, but fluorocarbons are weak XB donors because the alkyl group connected to the fluorine is not electronegative. In addition, the Lewis base (XB acceptor) tends to be electronegative as well and anions are better XB acceptors thaeutral molecules.

Halogen bonds are strong, specific, and directional interactions that give rise to well-defined structures. Halogen bond strengths range from 5–180 kJ/mol. The strength of XB allows it to compete with HB, which are a little bit weaker in strength. Halogen bonds tend to form at 180° angles, which was shown in Odd Hassel’s studies with bromine and 1,4-dioxane in 1954. Another contributing factor to halogen bond strength comes from the short distance between the halogen (Lewis acid, XB donor) and Lewis base (XB acceptor). The attractive nature of halogen bonds result in the distance between the donor and acceptor to be shorter than the sum of van der Waals radii. The XB interaction becomes stronger as the distance decreases between the halogen and Lewis base.

History

In 1863, Frederick Guthrie gave the first report on the ability of halogen atoms to form well-defined adducts with electron donor species. In his experiment, he added I2 to a saturated solution of ammonium nitrate to form NH₃I₂.. When the compound was exposed to air, it spontaneously decomposed into ammonia and iodine which allowed Guthrie to conclude that he had formed NH₃I₂.

In the 1950s, Robert S. Mulliken developed a detailed theory of electron donor-acceptor complexes, classifying them as being outer or inner complexes. Outer complexes were those in which the intermolecular interaction between the electron donor and acceptor were weak and had very little charge transfer. Inner complexes have extensive charge redistribution. Mulliken’s theory has been used to describe the mechanism by which XB formation occurs.

Around the same time period that Mulliken developed his theory, crystallographic studies performed by Hassel began to emerge and became a turning point in the comprehension of XB formation and its characteristics.

The first X-ray crystallography study from Hassel’s group came in 1954. In the experiment, his group was able to show the structure of bromine 1,4-dioxanate using x-ray diffraction techniques. The experiment revealed that a short intermolecular interaction was present between the oxygen atoms of dioxane and bromine atoms. The O−Br distance in the crystal was measured at 2.71 Å, which indicates a strong interaction between the bromine and oxygen atoms. In addition, the distance is smaller than the sum of the van der Waals radii of oxygen and bromine (3.35 Å). The angle between the O−Br and Br−Br bond is about 180°. This was the first evidence of the typical characteristics found in halogen bond formation and led Hassel to conclude that halogen atoms are directly linked to electron pair donor with a bond direction that coincides with the axes of the orbitals of the lone pairs in the electron pair donor molecule.

In the 1980s continued work was carried out using analytical methods such as infrared spectroscopy and Fourier transform spectroscopy. These methods allowed the isolation of complexes formed between Lewis bases and halogen molecules for further studies.

Applications

Crystal Engineering

Crystal engineering is a growing research area that bridges solid-state and supramolecular chemistry. This unique field is interdisciplinary and merges traditional disciplines such as crystallography, organic chemistry, and inorganic chemistry. In 1971, Schmidt first established the field with a publication on photodimerization in the solid-state. The more recent definition identifies crystal engineering as the utilization of the intermolecular interactions for crystallization and for the development of new substances with different desired physicochemical properties. Before the discovery of halogen bonding, the approach for crystal engineering involved using hydrogen bonding, coordination chemistry and inter-ion interactions for the development of liquid-crystalline and solid-crystalline materials. Furthermore, halogen bonding is employed for the organization of radical cationic salts, fabrication of molecular conductors, and creation of liquid crystal constructs. Since the discovery of halogen bonding, new molecular assemblies exist. Due to the unique chemical nature of halogen bonding, this intermolecular interaction serves as an additional tool for the development of crystal engineering.

The first reported use of halogen bonding in liquid crystal formation was by H. Loc Nguyen. In an effort to form liquid crystals, alkoxystilbazoles and pentafluoroiodobenzene were used. Previous studies by Metrangolo and Resnati demonstrated the utility of pentafluoroiodobenzene for solid-state structures. Various alkoxystilbazoles have been utilized for nonlinear optics and metallomesogens. Using another finding of Resnati (e.g. N−I complexes form strongly), the group engineered halogen-bonded complexes with iodopentafluorobenzene and 4-alkoxystilbazoles. X-ray crystallography revealed a N−I distance of 2.811(4) Å and the bonding angle to be 168.4°. Similar N−I distances were measured in solid powders. The N−I distance discovered is shorter than the sum of the Van Der Waals radii for nitrogen and iodine (3.53 Å). The single crystal structure of the molecules indicated that no quadrupolar interactions were present. Interestingly, the complexes in Figure 4 were found to be liquid-crystalline.

To test the notion of polarizability involvement in the strength of halogen bonding, bromopentafluorbenzene was used as a Lewis base. Consequently, verification of halogen bond complex formation wasn’t obtained. This finding provides more support for the dependence of halogen bonding on atomic polarizability. Utilizing similar donor-acceptor frameworks, the authors demonstrated that halogen bonding strength in the liquid crystalline state is comparable to the hydrogen-bonded mesogens.

Preparation of poly(diiododiacetylene)

Applications utilizing properties of conjugated polymers emerged from work done by Heeger, McDiaramid, and Shirakawa with the discovery that polyacetylene is a conducting, albeit difficult to process material. Since then, work has been done to mimic this conjugated polymer’s backbone (e.g., poly(p-phenylenevinylene)). Conjugated polymers have many practical applications, and are used in devices such as photovoltaic cells, organic light-emitting diodes, field-effect transistors, and chemical sensors. Goroff et al. prepared ordered poly(diiododiacetylene) (PIDA) via prearrangement of monomer (2) with a halogen bond scaffolding. PIDA is an excellent precursor to other conjugated polymers, as Iodine can be easily transformed. For instance, C−I cleavage is possible electrochemical reduction.

Crystal structures of monomer (2) are disordered materials of varying composition and connectivity. Hosts (3–7) were investigated for their molecular packing, primarily by studying co-crystals of monomer (2) and respective host. Both (3) and (4) pre-organized monomer (2), but steric crowding around the iodines prevented successful topological polymerization of the monomer. Hosts (5–7) utilize hydrogen bonds and halogen bonds to hold monomer (2) at an optimal distance from each other to facilitate polymerization.

In fact, when host 7 was used, polymerization occurred spontaneously upon isolation of the co-crystals. Crystal structures show the polymer strands are all parallel to the hydrogen-bonding network, and the host nitriles are each halogen-bonded to iodine atoms. Interestingly, half of the iodine atoms in (1) in the crystal are in close contact to the oxalamide oxygen atoms. Oxygen atoms of host 7 are acting as both hydrogen and halogen bond acceptors.

Above is a representation of a crystal structure of 1 and 7. As shown, the oxalamide oxygen (in purple) forms a hydrogen bond with the amide below (blue dashed line) and forms a weak halogen bond with the iodine on 1 (purple dashed line). This weak halogen bond further stabilizes this co-crystal. Halogen bond between the nitrile and iodine is represented with a red dashed line.

Porous structures

Porous structures have a variety of uses. Many chemists and material scientists are working to improve metal-organic frameworks (MOFs) to store hydrogen to use in cars. These highly organized crystalline inclusion complexes have potential uses in catalysis and molecular separation devices. Molecular organization is often controlled via intermolecular forces such as hydrogen bonding. However, utilizing hydrogen bonding often limits the range of pore sizes available due to close packing.

Pigge, et al., utilized halogen bonding interactions between amines, nitrogen heterocycles, carbonyl groups, and other organic halides, to construct their porous structures. This is significant because organic crystalline networks mediated by halogen bonds, an interaction significantly weaker than hydrogen bond, are rare.

Crystal structures of 1 and 2  were obtained in a variety of solvents, such as dichloromethane, pyridine, and benzene. The authors note that the porous inclusion complexes appear to be mediated in part by unprecedented I-π interactions and by halogen bond between iodine and carbonyl groups. The crystal structure come together in a triangular array and molecules of 2 are approximately symmetric. Additionally, all of the sets of halogen bonding interactions are not identical, and all of the intermolecular interactions between halogen and halogen bond acceptor slightly exceed the sum of the Van der Waals radius, signifying a slightly weaker halogen bond, which leads to more flexibility in the structure. The 2D layers stack parallel to each other to produce channels filled with solvent.

Solvent interactions are also noted in the formation of the hexagonal structures, especially in pyridine and chloroform. Initially, crystals that form these solutions form channeled structures. Over time, new needle-like solvate-free structures form are packed tighter together, and these needles are actually the thermodynamically favored crystal. The authors hope to use this information to better understand the complementary nature of hydrogen bonds and halogen bonds in order to design small molecules predict structures.

Halogen Bonding in Biological Macromolecules

For some time, the significance of halogen bonding to biological macromolecular structure was overlooked. Based on single-crystal structures in the protein data bank (PDB) (July 2004 version), a study by Auffinger and others on single crystals structures with 3 Å resolution or better entered into the PDB revealed that over 100 halogen bonds were found in six halogenated-based nucleic acid structures and sixty-six protein-substrate complexes for halogen-oxygen interactions. Although not as frequent as halogen-oxygen interactions, halogen-nitrogen and halogen-sulfur contacts were identified as well. These scientific findings provide a unique basis for elucidating the role of halogen bonding in biological systems.

On the bio-molecular level, halogen bonding is important for substrate specificity, binding and molecular folding. In the case of protein-ligand interactions, the most common charge-transfer bonds with polarizable halogens involve backbone carbonyls and/or hydroxyl and carboxylate groups of amino acid residues. Typically in DNA and protein-ligand complexes, the bond distance between Lewis base donor atoms (e.g. O, S, N) and Lewis acid (halogen) is shorter than the sum of their Van der Waals radius. Depending on the structural and chemical environment, halogen bonding interactions can be weak or strong. In the case of some protein-ligand complexes, halogen bonds are energetically and geometrically comparable to that of hydrogen bonding if the donor-acceptor directionality remains consistent. This intermolecular interaction has been shown to be stabilizing and a conformational determinant in protein-ligand and DNA structures.

For molecular recognition and binding, halogen bonding can be significant. An example of this assertion in drug design is the substrate specificity for the binding of IDD 594 to human aldose reductase. E.I. Howard reported the best resolution for this monomeric enzyme. This biological macromolecule consists of 316 residues, and it reduces aldoses, corticosteroids, and aldehydes. D-sorbitol, a product of the enzymatic conversion of D-glucose, is thought to contribute to the downstream effects of the pathology of diabetes. Hence, inhibiting this enzyme has therapeutic merit.

Aldehyde-based and carboxylate inhibitors are effective but toxic because the functional activity of aldehyde reductase is impaired. Carboxylate and aldehyde inhibitors were shown to hydrogen bond with Trp 111, Tyr 48, and His 110. The “specificity pocket,” created as a result of inhibitor binding, consists of Leu 300, Ala 299, Phe 122, Thr 113, and Trp 111. For inhibitors to be effective, the key residues of interaction were identified to be Thr 113 and Trp 111. IDD 594 was designed such that the halogen would provide selectivity and be potent. Upon binding, this compound induces a conformational change that causes halogen bonding to occur between the oxygen of the Thr and the bromine of the inhibitor. The bond distance was measured to be 2.973(4) Å. It is this O−Br halogen bond that contributes to the large potency of this inhibitor for human aldose reductase rather than aldehyde reductase.

Chemical Reactions and Compounds

All halogens are quite reactive, and in the natural world they always occur combined with other elements. Fluorine reacts so readily with almost any substance it contacts that chemists were not successful in isolating pure fluorine until 1886, although its existence in compounds had been known for many years. Chlorine, bromine, and iodine are progressively less reactive but still form compounds with most other elements, especially metals. A good example is mercury, whose reaction with bromine was discussed in the section covering macroscopic and microscopic views of a chemical reaction. Mercury reacts with other halogens in the same way:

 

Hg(l) + X₂(g, l, or s) → HgX₂ (s)      X = F, Cl, Br, or I

Already covered in the section on alkali metals, halogens react readily with alkali metals with the general form of:

2M + X₂ → 2MX       M = Li, Na, K, Rb, or Cs and X = F, Cl, Br, I

Iodine combines less vigorously with alkali metals than other halogens, but its reactions are analogous to the reactions of alkali metals with florine, chlorine and bromine. Compounds of an alkali metal and a halogen, such as sodium chloride, potassium fluoride, lithium bromide, or cesium iodide, have closely related properties. (All taste salty, for example.) They belong to a general category called salts, all of whose members are similar to ordinary table salt, sodium chloride. The term halogen is derived from Greek words meaning “salt former.”

Halogens also react with alkaline-earth metals in the general reaction:

M + X₂ → MCl₂      M = Be, Mg, Ca, Sr, Ba, or Ra and X = F, Cl, Br, I

Another vigorous reaction occurs when certain compounds containing carbon and hydrogen contact the halogens. Turpentine, C₁₀H₁₆, reacts quite violently. In the case of fluorine and chlorine the equation is but the products are different when bromine and iodine react.

C₁₀H₁₆(l) + 8X₂(g) → 10C(s) + 16HX(g)      X = F, Cl

Before the advent of the automobile, veterinarians used solid iodine and turpentine to disinfect wounds in horses’ hooves. This may have been because of the superior antiseptic qualities of the mixture. However, a more likely reason is the profound impression made on the owner of the horse by the great clouds of violet iodine vapor which sublimed as a result of the increase in temperature when the reaction occurred!

The violent reaction is due to α-pinene in turpentine. The relief of ring strain is highly exothermic. This temperature increase causes the sublimation leading to the impressive violet iodine vapor.

The halogens also react directly with hydrogen, yielding the hydrogen halides:

H₂ + X₂ → 2HX      X = F, Cl, Br, I

These compounds are all gases, are water soluble, and, except for HF, are strong acids in aqueous solution. They are conveniently prepared in the laboratory by acidifying the appropriate sodium or other halide:

NaCl(s) + H₃O⁺(aq) Na⁺(aq) + H₂O(l) + HCl(g)      (1)

The acid must be nonvolatile so that heating will drive off only the gaseous hydrogen halide. In the case of fluorides and chlorides, H₂SO₄ will do, but bromides and iodides are oxidized to Br₂ or I₂ by hot H₂SO₄ and so H₃PO₄ is used instead.

A reaction similar to Eq. (1) occurs when phosphate rock containing fluorapatite is treated with H₂SO₄ to make fertilizer:

Ca₁₀(PO₄)₆F₂ + 7H₂SO₄ + 3H₂O → 3Ca(H₂PO₄) ₂•H₂O + 7CaSO₄ + 2HF

The HF produced in this reaction can cause significant air- pollution problems. Fluorides are also emitted to the atmosphere in steelmaking and aluminum production. There is some evidence that fluorides, rather than sulfur dioxide, may have been responsible for human deaths in air-pollution episodes at Donora, Pennsylvania, and the Meuse Valley in Belgium.

The relative oxidizing strengths of the halogens can be illustrated nicely in the laboratory. If, for example, a solution of Cl₂ in H₂O is combined with a solution of NaI, the dark color of I₂ can be observed, showing that the Cl₂ has oxidized the I^–:

Cl₂(aq) + 2I^–(aq) → 2Cl^–(aq) + I₂(aq)

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The experimental solution is on the far left, and contain Cl₂ in water, which is covered by a layer of hexane, a nonpolar solvent which is immiscible with H₂O. The three other solutions, from left to right are a Cl₂ solution, a Br₂ solution, and an I₂ solution. When a solution with iodide ions is added to the experimental solution, nonpolar I₂ molecules are formed. They concentrate in the hexane layer, and a beautiful violet color can be observed, the same as I₂ solution. From such experiments it can be shown that the strongest oxidizing agent is F₂ (at the top of the group). F₂ will react with Cl^–, Br^–, and I^–. The weakest oxidizing agent, I₂, does not react with any of the halide ions.

The extremely high oxidizing power of F₂ makes it the only element which can combine directly with a noble gas.

Xe(g) + F₂(g) → XeF₂(s)

XeF₂(s) + F₂(g) → XeF₄(s)

XeF₄(s) + F₂(g) → XeF₆(s)

The reactions may be used to synthesize the three xenon fluorides, all of which are strong oxidizing agents. When an electrical discharge is passed through a mixture of Kr and F₂ at a low temperature, KrF₂ can be formed. This is the only compound of Kr, and it decomposes slowly at room temperature.

Fluorine is also set apart from the other halogens because of its ability to oxidize water:

3F₂ + 6H₂O → 4H₃O⁺ + 4F^– + O₂

Chlorine is also capable of oxidizing water, but it does so very slowly. Instead the reaction goes partway to completion.

Cl₂ + 2H₂O H₃O⁺ + Cl^– + HOCl

Hypochlorous acid, HOCl, is a weak acid. Small concentrations of hypobromous and hypoiodous acids can also be obtained in this way. In basic solution the halogen is completely consumed, producing the hypohalite anion:

Cl₂ + 2OH^– → Cl^– + H₂O + OCl^–

Since hypochlorite, OCl^–, could also be supplied from an ionic compound such as NaOCl, the latter is often used to chlorinate swimming pools.

Hypohalite ions disproportionate in aqueous solution:

3OCl^– → 2Cl^– + ClO₃^–

This reaction is rather slow for hypochlorite unless the temperature is above 75°C, but OBr^– and OI^– are consumed immediately at room temperature. Chlorate, ClO₃^–, bromate, BrO₃^–, and iodate, IO₃^–, salts can be precipitated from such solutions. All are good oxidizing agents. Potassium chlorate, KClO₃, decomposes, giving O₂ when heated in the presence of a catalyst:

2KClO₃ 2KCl + 3O₂

This is a standard laboratory reaction for making O₂.

If KClO₂ is heated without a catalyst, potassium perchlorate, KClO₄, may be formed. Perchlorates oxidize organic matter rapidly and often uncontrollably. They are notorious for exploding unexpectedly and should be handled with great care.

One other interesting group of compounds is the interhalogens, in which one halogen bonds to another. Some interhalogens, such as BrCl, are diatomic, but the larger halogen atoms have room for several smaller ones around them. Thus compounds such as ClF₃, BrF₃ and BrF₅, and IF₃, ICl₃, IF₅, and IF₇ can be synthesized. Note that the largest halogen atom I can accommodate three chlorines and up to seven fluorines around it. The following video showcases a reaction which involves some of these interhalogens:

The process begins with a test tube containing a layer of KI aqueous solution on top of CCl₄ below it. Chlorine is bubbled through the KI layer. Due to oxidizing strength of the halogens, Cl₂ reacts with I^– to form iodine, according to the reaction:

2I^–(aq ) + Cl₂(aq ) → I₂(aq ) + 2Cl^–(aq)

A brown triiodide ion is also formed in the aqueous layer, according to the reaction:

I^–(aq ) + I₂(aq ) → I3^–(aq )

A purple solution begins to form in the CCl₄ layer, as iodine dissolves in it. The iodine in the aqueous layer also reacts with the excess Cl₂ to form the red ICl, according to the following reaction:

I₂(aq ) + Cl₂(aq ) –> 2ICl(aq )

The final reaction takes place as more Cl₂ is added, which reacts with ICl, to form the yellow ICl₃. This reaction causes the aqueous solution to decolorize. This goes according to the reaction:

ICl(aq ) + Cl₂(aq ) → ICl₃(aq )

Among the most important compounds of halogens in positive oxidation states are the oxoacids of Cl, Br, and I and the corresponding oxoacid salts. In these compounds, the halogen shares its valence electrons with oxygen, a more electronegative element. (Electronegativities are O, 3.5; Cl, 3.0; Br, 2.8; I, 2.5.) The general formula for a halogen oxoacid, is HXO_n and the oxidation state of the halogen is +1, +3, +5, or +7, depending on the value of n.

Fluorine does not form any oxoacid since it is the strongest oxidizing agent. Chlorine, bromine and iodine mainly form four series of oxoacids namely hypohalous acid (HXO), halous acid (HXO₂) halic acid (HXO₃) and perhalic acid (HXO₄) as given below :

 

 

(i) Hybridized ion : In all these oxoacids, the halogen atom is sp³-hybridized.

(ii) Acidic character : All these acids are monobasic containing an—OH group. The acidic character of the oxoacids increases with increase in oxidatioumber, i.e., HClO < HClO₂ < HClO₃ < HClO₄ and the strength of the conjugate bases of these acids follows the order,

ClO > ClO> ClO > ClO₄^–

(iii) Oxidising power and thermal stability : The oxidizing power of these acids decreases as the oxidatioumber increases, i.e., HClO < HClO₂ < HClO₃ < HClO₄. Stability of oxoacids of chlorine in the increasing order is, HClO < HClO₂ < HClO₃ < HClO₄ and the increasing stability order of anions of oxoacids of chlorine is, CIO^– < CIO₂^– < CIO₃^– < CIO₄^–.

As the number of oxygen atoms in an ion increases there will be a greater dispersal of negative charge and thus greater will be the stability of ion formed. For different halogen having the name oxidatioumber, the thermal stability decreases with increase in atomic number i.e., it is in the order HClO > HBrO > HIO and ClO^– > BrO^– > IO^– However, in HXO₃ is most stable. The stability order being HClO₃ < HBrO₃ < HIO₃.

(iv) Perhalates are strong oxidizing agents, the oxidizing power is in the order, BrO₄^–, IO₄^– > CIO₄^–.

Thus BrO₄ is the strongest oxidizing agent (though its reaction is quite slow) and CIO₄^– is the weakest.

(v) The acidity of oxoacids of different halogens having the same oxidatioumber decreases with increase in the atomic size of the halogen i.e. HCIO₄ > HBrO₄ > HIO₄.

Only four acids have been isolated in pure form: perchloric acid (HClO₄), iodic acid (HIO₃) and the two periodic acids, metaperiodic acid (HIO₄) and paraperiodic acid (H₅IO₆) The others are stable only in aqueous solution or in the form of their salts. Chlorous acid (HClO₂) is the only known halous acid.

The acid strength of the halogen oxoacids increases with the increasing oxidation state of the halogen. For example, acid strength increases from HClO, a weak acid (Ka =3.5*10^-8) to a very strong acid (Ka>>1). The acidic proton is bonded to oxygen, not to the halogen, even though we usually write the molecular formula of these acids as HXO_n. All the halogen oxoacids and their salts are strong oxidizing agents.

A hypohalous acid is formed when Cl₂,Br₂, or I₂ dissolves in cold water. In this reaction, the halogen disproportionates, going to the +1 oxidation state in HOX and -1 the state in HX. The equilibrium lies to the left but is shifted to the right in basic solution:

Large amounts of aqueous sodium hypochlorite (NaOCl) are produced in the chlor-alkali industry when the Cl₂ gas and aqueous NaOH from the electrolysis of aqueous NaCl are allowed to mix. Aqueous NaOCl is a strong oxidizing agent and is sold in a 5% solution as chlorine bleach.

Further disproportionation of OCl^– to ClO₃^– and Cl^– is slow at room temperature but becomes fast at higher temperatures. Thus, when Cl₂ gas reacts with hot aqueous NaOH, it gives a solution that contains sodium chlorate rather NaClO₃ than NaOCl:

Chlorate salts are used as weed-killers and as strong oxidizing agents. Potassium chlorate, for example, is an oxidizer in matches, fireworks, and explosives. It also reacts vigorously with organic matter.

Sodium perchlorate (NaClO₄) is produced commercially by the electrolytic oxidation of aqueous sodium chlorate and is converted to perchloric acid by reaction with concentrated HCl:

The HClO₄ is then concentrated by distillation at reduced pressure.

Pure, anhydrous perchloric acid is a colorless, shock-sensitive liquid that decomposes explosively on heating. It is a powerful and dangerous oxidizing agent, violently oxidizing organic matter and rapidly oxidizing even silver and gold. Perchlorate salts are also strong oxidants, and they too must be handled with caution. Ammonium perchlorate (NH₅ClO₄) in fact, is the oxidizer in the solid booster rockets used to propel the space shuttle.

Iodine differs from the other halogens because it forms more than one perhalic acid. Paraperiodic acid (H₅IO₆) is obtained as white crystals (mp 128°C) when periodic acid solutions are evaporated. When heated to 100°C at reduced pressure, these crystals lose water and are converted to metaperiodic acid (HIO₄):

Metaperiodic acid is a strong monoprotic acid, whereas paraperiodic acid is a weak polyprotic acid (K_a1=5.1 *10^–4; K_a2 = 4.9 *10^–9). It has an octahedral structure in which a central iodine atom is bonded to one O atom and five OH groups:

Chlorine and bromine do not form perhalic acids of the type H₅XO₆) because their smaller sizes favor a tetrahedral structure over an octahedral one.

Halides.  

A halide is a binary compound, of which one part is a halogen atom and the other part is an element or radical that is less electronegative than the halogen, to make a fluoride, chloride, bromide, iodide, or astatide compound. Many salts are halides. All Group 1 metals form halides with the halogens and they are white solids.

NaF + HOH Û HF + NaOH.

CaF₂ + H₂SO₄ = CaSO₄ + 2HF

A halide ion is a halogen atom bearing a negative charge. The halide anions are fluoride (F⁻), chloride (Cl⁻), bromide (Br⁻), iodide (I⁻) and astatide (At⁻). Such ions are present in all ionic halide salts.

2Fe + 3Cl₂ = 2FeCl₃

H₂ + Cl₂ = 2HCl

NaCl + H₂SO₄ = NaHSO₄ + HCl­

2NaCl + H₂SO₄ = Na₂SO₄ = 2HCl­

 

Halides in organic chemistry

In organic chemistry halides represent a functional group. Any organic compound that contains a halogen atom can be considered a halide. Alkyl halides are organic compounds of the type R-X, containing an alkyl group R covalently bonded to a halogen X.

Fluorine

Fluorine (pronounced /flʊəriːn/, Latin: fluere, meaning “to flow”), is the chemical element with the symbol F and atomic number 9. Atomic fluorine is univalent and is the most chemically reactive and electronegative of all the elements. In its elementally isolated (pure) form, fluorine is a poisonous, pale, yellowish brown gas, with chemical formula F2. Like other halogens, molecular fluorine is highly dangerous; it causes severe chemical burns on contact with skin.

Fluorine’s large electronegativity and small atomic radius gives it interesting bonding characteristics, particularly in conjunction with carbon.

Notable characteristics

Pure fluorine (F₂) is a corrosive pale yellow or brown gas that is a powerful oxidizing agent. It is the most reactive and most electronegative of all the elements (4.0), and readily forms compounds with most other elements. It has an oxidatioumber -1, except when bonded to another fluorine in F₂ which gives it an oxidatioumber of 0. Fluorine even combines with argon, krypton, xenon, and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. The reaction with hydrogen occurs even at extremely low temperatures, using liquid hydrogen and solid fluorine. It is so reactive that metals, and even water, as well as other substances, burn with a bright flame in a jet of fluorine gas. It is far too reactive to be found in elemental form. In moist air it reacts with water to form also-dangerous hydrofluoric acid.

Fluorides are compounds that combine fluorine with some positively charged counterpart. They often consist of crystalline ionic salts. Fluorine compounds with metals are among the most stable of salts.

Hydrogen fluoride is a weak acid when dissolved in water. Consequently, fluorides of alkali metals produce basic solutions.

Video  Physical properties of halogens http://www.youtube.com/watch?v=cwcVjh559KY&feature=BF&list=ULwx2gSk2-sz4&index=5

 

Chemical uses of fluorine:

·                        Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMS (microelectromechanical systems) fabrication. Xenon difluoride is also used for this last purpose.

•  Hydrofluoric acid (chemical formula HF) is used to etch glass in light bulbs and other products.

• Fluorine is indirectly used in the production of low friction plastics such as Teflon (or polytetrafluoroethylene), and in halons such as freon.

• Along with some of its compounds, fluorine is used in the production of pure uranium from uranium hexafluoride and in the synthesis of numerous commercial fluorochemicals, including vitally important pharmaceuticals, agrochemical compounds, lubricants, and textiles.

• Sodium hexafluoroaluminate (cryolite), is used in the electrolysis of aluminium.

• In much higher concentrations, sodium fluoride has been used as an insecticide, especially against cockroaches.

• Fluorides have been used in the past to help molten metal flow, hence the name.

• Some researchers including US space scientists in the early 1960s have studied elemental fluorine gas as a possible rocket propellant due to its exceptionally high specific impulse. The experiments failed because fluorine proved difficult to handle, and its combustion products proved extremely toxic and corrosive.

• Compounds of fluorine such as fluoropolymers, potassium fluoride and cryolite are utilized in applications such as anti-reflective coatings and dichroic mirrors on account of their unusually low refractive index.

Dental and medical uses:

• Compounds of fluorine, including sodium fluoride (NaF), stannous fluoride (SnF₂) and sodium MFP, are used in toothpaste to prevent dental cavities. These compounds are also added to municipal water supplies, a process called water fluoridation, though a number of health concerns has sometimes led to controversy.

• Many important agents for general anesthesia such as sevoflurane, desflurane, and isoflurane are hydrofluorocarbon derivatives.

• The fluorinated antiinflammatories dexamethasone and triamcinolone are among the most potent of the synthetic corticosteroids class of drugs.

• Fludrocortisone (“Florinef”) is one of the most common mineralocorticoids, a class of drugs which mimics the actions of aldosterone.

• Fluconazole is a triazole antifungal drug used in the treatment and prevention of superficial and systemic fungal infections.

Compounds

Fluorine forms a variety of very different compounds, owing to its small atomic size and covalent behavior, and on the other hand, its oxidizing ability and extreme electronegativity. For example, hydrofluoric acid is extremely dangerous, while in synthetic drugs incorporating an aromatic ring (e.g. flumazenil), fluorine is used to prevent toxication or to delay metabolism.

The fluoride ion is basic, therefore hydrofluoric acid is a weak acid in water solution. However, water is not an inert solvent in this case: when less basic solvents such as anhydrous acetic acid are used, hydrofluoric acid is the strongest of the hydrohalogenic acids. Also, owing to the basicity of the fluoride ion, soluble fluorides give basic water solutions. The fluoride ion is a Lewis base, and has a high affinity to certain elements such as calcium and silicon. For example, deprotection of silicon protecting groups is achieved with a fluoride. The fluoride ion is poisonous.

Fluorine as a freely reacting oxidant gives the strongest oxidants known. Chlorine trifluoride, for example, can burn water and sand, both compounds of a weaker oxidant, oxygen.

Fluorine compounds involving noble gases were first synthesised by Neil Bartlett in 1962 –xenon hexafluoroplatinate, XePtF6, being the first. Fluorides of krypton and radon have also been prepared. Also argon fluorohydride has been prepared, although it is only stable at cryogenic temperatures.

The carbon-fluoride bond is covalent and very stable. The use of a fluorocarbon polymer, poly(tetrafluoroethene) or Teflon, is an example: it is thermostable and waterproof enough to be used in frying pans. Organofluorines may be safely used in applications such as drugs, without the risk of release of toxic fluoride. In synthetic drugs, toxication can be prevented. For example, an aromatic ring is useful but presents a safety problem: enzymes in the body metabolize some of them into poisonous epoxides. When the para position is substituted with fluorine, the aromatic ring is protected and epoxide is no longer produced.

The substitution of hydrogen for fluorine in organic compounds offers a very large number of compounds. An estimated fifth of pharmaceutical compounds and 30% of agrochemical compounds contain fluorine. The -CF₃ and -OCF₃ moieties provide further variation, and more recently the -SF₅ group.

 

Fluorite (CaF₂) crystals

This element is recovered from fluorite, cryolite, and fluorapatite.

Preparation

Industrial fluorine production starts with fluorspar (CaF₂), which is heated with sulfuric acid (H₂SO₄) to produce anhydrous hydrogen fluoride (HF). The hydrogen fluoride is added to potassium fluoride (KF) to make potassium bifluoride (KHF₂). Electrolysis of potassium bifluoride produces fluorine gas at the anode, and hydrogen gas at the cathode. This is essentially the same method employed by Moissan in 1886; the use of potassium bifluoride rather than hydrogen fluoride itself aids electrolysis by greatly increasing the conductivity.

2 CaF₂ + H₂SO₄ → 2 HF + CaSO₄

HF + KF → KHF₂

2 KHF₂ → 2 HF + H₂ + F₂

In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl Christe discovered a purely chemical preparation involving the reaction of solutions in anhydrous HF, K₂MnF₆, and SbF₅ at 150 °C:

K₂MnF₆ + 2SbF₅ → 2KSbF₆ + MnF₃ + ½F₂

Though not a practical synthesis, it demonstrates that electrolysis is not essential.

Safety

Elemental fluorine

Elemental fluorine (fluorine gas) is a highly toxic, corrosive oxidant, which can cause organic material, combustibles, or other flammable materials to ignite. It must be handled with great care and any contact with skin and eyes should be strictly avoided. Fluorine gas has a characteristic pungent odor that is detectable in concentrations as low as 20 nL/L. As it is so reactive, all materials of construction must be carefully selected. All metal surfaces must be passivated before exposure to fluorine.

Fluoride ion

Fluoride ions are also highly toxic and must also be handled with great care and any contact with skin and eyes should be strictly avoided.

Hydrogen fluoride and hydrofluoric acid

Contact of exposed skin with hydrofluoric acid solutions poses one of the most extreme and insidious industrial threats—one which is exacerbated by the fact that hydrofluoric acid damages nerves in such a way as to make such burns initially painless.

CaF₂ + H₂SO₄ = CaSO₄ + 2HF

The HF molecule is a weaker acid which is significantly non-dissociated in water, and the intact molecule is capable of rapidly migrating through lipid layers of cells which would ordinarily stop an ion or partly ionized acid, and the burns it produces are typically deep.

H₂O + HF = H₃O⁺ + F^–,   Ka = 7,2 × 10^-4.

HF + F^– = HF₂^–,   K = 5,1

2HF + H₂O Û H₃O⁺ + HF₂^–.

HF may react with calcium, permanently damaging the bone . More seriously, HF reaction with the body’s calcium inside cells can cause cardiac arrhythmias, followed by cardiac arrest brought on by sudden chemical changes within the body (hypocalcaemia). These cannot always be prevented with local or intravenous injection of calcium salts. Hydrofluoric acid spills over just 2.5% of the body’s surface area (about 75 in or 5 dm²), despite copious immediate washing, have been fatal. If the patient survives, hydrofluoric acid burns typically produce open wounds of an especially slow-healing nature.

SiO₂ + 4HF = SiF₄ + 2H₂O.

Anhydrous hydrogen fluoride will rapidly form hydrofluoric acid on contact with moisture; its physiological effects are then the same.

CHLORINE

Chlorine (IPA: /klɔərin/, Greek: χλωρóς chloros, meaning “pale green”), is the chemical element with atomic number 17 and symbol Cl. It is a halogen, found in the periodic table in group 17 (formerly VIIa or VIIb). As the chloride ion, which is part of common salt and other compounds, it is abundant iature and necessary to most forms of life, including humans. In its common elemental form (Cl2 or “dichlorine”) under standard conditions, it is a pale green gas about 2.5 times as dense as air. It has a disagreeable, suffocating odor that is detectable in concentrations as low as 3.5 ppm and is poisonous. Chlorine is a powerful oxidant and is used in bleaching and disinfectants. As a common disinfectant, chlorine compounds are used in swimming pools to keep them clean and sanitary. In the upper atmosphere, chlorine based molecules have been implicated in the destruction of the ozone layer.

Cl₂ + F₂ = 2ClF

H₂ + Cl₂ = 2HCl

Cl₂ + NaOH ® NaClO + NaCl + H₂O

Characteristics.

Chlorine gas in a plastic container. It is not advisable to store chlorine in this manner.

Chlorine gas is diatomic, with the formula Cl₂. It combines readily with all elements except O₂ and N₂ and the noble gases. Compounds with oxygen, nitrogen, and xenon are known but do not form by direct reaction of the elements. Chlorine is not as extremely reactive as fluorine. Pure chlorine gas does, however, support combustion of organic compounds such as hydrocarbons, although the carbon component tends to burn incompletely, with much of it remaining as soot. At 10 °C and atmospheric pressure, one liter of water dissolves 3.10 L of gaseous chlorine, and at 30°C, 1 L of water dissolves only 1.77 liters of chlorine.

Interaction with metals:

2Fe + 3Cl₂ = 2FeCl₃;

Cu + Cl₂ = CuCl₂.

Interaction with non-metals:

H₂ + Cl₂ = 2HCl

2S + Cl₂ = S2Cl₂

Si + 2Cl₂ = SiCl₄

2P + 5Cl = 2PCl₅.

Cl₂ + H₂O Û HCl + HClO

This element is a member of the salt-forming halogen series and is extracted from chlorides through oxidation often by electrolysis. As the chloride ion, Cl−, it is also the most abundant dissolved ion in ocean water.

Isotopes: Chlorine has isotopes with mass numbers ranging from 32 to 40. There are two principal stable isotopes, ³⁵Cl (75.77%) and ³⁷Cl (24.23%), giving chlorine atoms in bulk an apparent atomic weight of 35.5 g/mol.

³⁶Cl Trace amounts of radioactive ³⁶Cl exist in the environment, in a ratio of about 7×10⁻¹³ to 1 with stable isotopes. ³⁶Cl is produced in the atmosphere by spallation of ³⁶Ar by interactions with cosmic ray protons. In the subsurface environment, ³⁶Cl is generated primarily as a result of neutron capture by ³⁵Cl or muon capture by ⁴⁰Ca. ³⁶Cl decays to ³⁶S and to ³⁶Ar, with a combined half-life of 308,000 years. The half-life of this hydrophilic nonreactive isotope makes it suitable for geologic dating in the range of 60,000 to 1 million years. Additionally, large amounts of ³⁶Cl were produced by irradiation of seawater during atmospheric detonations of nuclear weapons between 1952 and 1958. The residence time of ³⁶Cl in the atmosphere is about 1 week. Thus, as an event marker of 1950s water in soil and ground water, ³⁶Cl is also useful for dating waters less than 50 years before the present. ³⁶Cl has seen use in other areas of the geological sciences, including dating ice and sediments.

Occurrence. Iature, chlorine is found primarily as the chloride ion, a component of the salt that is deposited in the earth or dissolved in the oceans — about 1.9% of the mass of seawater is chloride ions. Even higher concentrations of chloride are found in the Dead Sea and in underground brine deposits. Most chloride salts are soluble in water, thus, chloride-containing minerals are usually only found in abundance in dry climates or deep underground. Common chloride minerals include halite (sodium chloride), sylvite (potassium chloride), and carnallite (potassium magnesium chloride hexahydrate). Over 2000 naturally-occurring organic chlorine compounds are known.

Industrially, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water. Along with chlorine, this chloralkali process yields hydrogen gas and sodium hydroxide, according to the following chemical equation:

2 NaCl + 2 H₂O → Cl₂ + H₂ + 2 NaOH

Production

Gas extraction. Chlorine can be manufactured by electrolysis of a sodium chloride solution (brine). The production of chlorine results in the co-products caustic soda (sodium hydroxide, NaOH) and hydrogen gas (H₂). These two products, as well as chlorine itself, are highly reactive. Chlorine can also be produced by the electrolysis of a solution of potassium chloride, in which case the co-products are hydrogen and caustic potash (potassium hydroxide). There are three industrial methods for the extraction of chlorine by electrolysis of chloride solutions, all proceeding according to the following equations:

Cathode: 2 H⁺ (aq) + 2 e⁻ → H₂ (g)

Anode: 2 Cl⁻ (aq) → Cl₂ (g) + 2 e⁻

Overall process: 2 NaCl (or KCl) + 2 H₂O → Cl₂ + H₂ + 2 NaOH (or KOH)

In diaphragm cell electrolysis, an asbestos (or polymer-fiber) diaphragm separates a cathode and an anode, preventing the chlorine forming at the anode from re-mixing with the sodium hydroxide and the hydrogen formed at the cathode. The salt solution (brine) is continuously fed to the anode compartment and flows through the diaphragm to the cathode compartment, where the caustic alkali is produced and the brine is partially depleted. Diaphragm methods produce dilute and slightly impure alkali but they are not burdened with the problem of preventing mercury discharge into the environment and they are more energy efficient. Membrane cell electrolysis employ permeable membrane as an ion exchanger. Saturated sodium (or potassium) chloride solution is passed through the anode compartment, leaving at a lower concentration. This method is more efficient than the diaphragm cell and produces very pure sodium (or potassium) hydroxide at about 32% concentration, but requires very pure brine.

Other methods. Before electrolytic methods were used for chlorine production, the direct oxidation of hydrogen chloride with oxygen or air was exercised in the Deacon process:

4 HCl + O₂ → 2 Cl₂ + 2 H₂O

This reaction is accomplished with the use of Copper(II) Chloride as a catalyst and is performed at high temperarature (about 400°C). The amount of extracted chlorine is approximately 80%. Due to the extremely corrosive reaction mixture, industrial use of this method is difficult and several pilot trials failed in the past. Nevertheless, recent developments are promising.

Another earlier process to produce chlorine was to heat brine with acid and manganese dioxide.

2 NaCl + 2H₂SO₄ + MnO₂ → Na₂SO₄ + MnSO₄ + 2 H₂O + Cl₂

Using this process, chemist Carl Wilhelm Scheele was the first to isolate chlorine in a laboratory. The manganese can be recovered by the Weldon process.

Small amounts of chlorine gas can be made in the laboratory by putting concentrated hydrochloric acid in a flask with a side arm and rubber tubing attached. Manganese dioxide is then added and the flask stoppered. The reaction is not greatly exothermic. As chlorine is denser than air, it can be easily collected by placing the tube inside a flask where it will displace the air. Once full, the collecting flask can be stoppered.

MnO₂ + 4HCl = MnCl₂ = Cl₂ + 2H₂O

2KMnO₄ + 16HCl = 2KCl + 2MnCl₂ + Cl₂ + 8H₂O.

Another method for producing small amounts of chlorine gas in a lab is by adding concentrated hydrochloric acid (typically about 5M) to sodium hypochlorite or sodium chlorate solution.

KClO₃ + 6HCl = KCl + 3Cl₂ + 3H₂O

 

Compounds

For general references to the chloride ion (Cl⁻), including references to specific chlorides, see chloride. For other chlorine compounds see chlorate (ClO₃⁻), chlorite (ClO₂⁻), hypochlorite(ClO⁻), and perchlorate(ClO₄⁻), and chloramine (NH₂Cl).

Other chlorine-containing compounds include:

·               Fluorides: chlorine monofluoride (ClF), chlorine trifluoride (ClF₃), chlorine pentafluoride (ClF₅)

·               Oxides: chlorine dioxide (ClO₂), dichlorine monoxide (Cl₂O), dichlorine heptoxide (Cl₂O₇);

·               Acids: hydrochloric acid (HCl), chloric acid (HClO₃), and perchloric acid (HClO₄)

 

Oxidation states.

Chlorine exists in all odd numbered oxidation states from −1 to +7, as well as the elemental state of zero. Progressing through the states, hydrochloric acid can be oxidized using manganese dioxide, or hydrogen chloride gas oxidized catalytically by air to form elemental chlorine gas. The solubility of chlorine in water is increased if the water contains dissolved alkali hydroxide. This is due to disproportionation:

Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O

Cl₂ + NaOH ® NaClO + NaCl + H₂O

In hot concentrated alkali solution disproportionation continues:

2ClO⁻ → Cl⁻ + ClO₂⁻

ClO⁻ + ClO₂⁻ → Cl⁻ + ClO₃⁻

Sodium chlorate and potassium chlorate can be crystallized from solutions formed by the above reactions. If their crystals are heated, they undergo the final disproportionation step.

4ClO₃⁻ → Cl⁻ + 3ClO₄⁻

4KClO₃ = 3KClO₄ + KCl

2KClO₃  =  2KCl + 3O₂­

This same progression from chloride to perchlorate can be accomplished by electrolysis. The anode reaction progression is:

Each step is accompanied at the cathode by 2H₂O + 2e⁻ → 2OH⁻ + H₂ −0.83 volts

 

Applications and uses

Production of industrial and consumer products. Chlorine’s principal applications are in the production of a wide range of industrial and consumer products. For example, it is used in making plastics, solvents for dry cleaning and metal degreasing, textiles, agrochemicals and pharmaceuticals, insecticides, dyestuffs, etc.

Purification and disinfection. Chlorine is an important chemical for water purification, in disinfectants, and in bleach. It is used (in the form of hypochlorous acid) to kill bacteria and other microbes in drinking water supplies and public swimming pools. However, in most private swimming pools chlorine itself is not used, but rather sodium hypochlorite (household bleach), formed from chlorine and sodium hydroxide, or solid tablets of chlorinated isocyanurates. Even small water supplies are now routinely chlorinated.

2Ca(OH)₂ + 2Cl₂ = CaCl₂ + Ca(OCl)₂ + 2H₂O

CaOCl₂ + CO₂ = CaCO₃ + Cl₂­

 

Chemistry. Elemental chlorine is an oxidizer. It undergoes halogen substitution reactions with lower halide salts. For example, chlorine gas bubbled through a solution of bromide or iodide anions oxidizes them to bromine and iodine respectively.

1/2Cl₂(g) + 1e ® Cl^–(g),   DG= -240 kJ/mol

1/2Cl₂(g) + 1e ® Cl^–(s),   DG= -131 kJ/mol

2FeCl₂ = Cl₂ = 2FeCl₃

H₂SO₃ + Cl₂ + H₂O = H₂SO₄ + 2HCl

Like the other halogens, chlorine participates in free-radical substitution reactions with hydrogen-containing organic compounds. This reaction is often – but not invariably – non-regioselective, and hence may result in a mixture of isomeric products. It is often difficult to control the degree of substitution as well, so multiple substitutions are common. If the different reaction products are easily separated, e.g. by distillation, substitutive free-radical chlorination (in some cases accompanied by concurrent thermal dehydrochlorination) may be a useful synthetic route. Industrial examples of this are the production of methyl chloride, methylene chloride, chloroform and carbon tetrachloride from methane, allyl chloride from propylene, and trichloroethylene and tetrachloroethylene from 1,2-dichloroethane.

Like the other halides, chlorine undergoes electrophilic additions reactions, most notably, the chlorination of alkenes and aromatic compounds with a Lewis acid catalyst. Organic chlorine compounds tend to be less reactive iucleophilic substitution reactions than the corresponding bromine or iodine derivatives, but they tend to be cheaper. They may be activated for reaction by substituting with a tosylate group, or by the use of a catalytic amount of sodium iodide.

Chlorine is used extensively in organic and inorganic chemistry as an oxidizing agent and in substitution reactions because chlorine often imparts many desired properties to an organic compound, due to its electronegativity.

Chlorine compounds are used as intermediates in the production of a number of important commercial products that do not contain chlorine. Examples are: polycarbonates, polyurethanes, silicones, polytetrafluoroethylene, carboxymethyl cellulose and propylene oxide.

Other uses. Chlorine is used in the manufacture of numerous organic chlorine compounds: methyl chloride, methylene chloride, chloroform, vinylidene chloride, trichloroethylene, perchloroethylene, allyl chloride, epichlorohydrin, chlorobenzene, dichlorobenzenes and trichlorobenzenes.

Chlorine is also used in the production of chlorates and in bromine extraction.

Vinyl chloride. The largest application of organochlorine chemistry is the production of vinyl chloride. The annual production in 1985 was around 13 billion kilograms, almost all of which was converted into polyvinylchloride (PVC).

Chloromethanes. Most low molecular weight chlorinated hydrocarbons such as chloroform, dichloromethane, dichloroethene, and trichloroethane are useful solvents. These solvents tend to be relatively non-polar; they are therefore immiscible with water and effective in cleaning applications such as degreasing and dry cleaning. Several billion kilograms of chlorinated methanes are produced annually, mainly by chlorination of methane:

CH₄ + x Cl₂ → CH₄−xClx + x HCl

The most important is dichloromethane, which is mainly used as a solvent. Chloromethane is a precursor to chlorosilanes and silicones. Historically significant, but smaller in scale is chloroform, mainly a precursor to chlorodifluoromethane (CHClF2) and tetrafluoroethene which is used in the manufacture of Teflon.

Pesticides. Many pesticides contain chlorine. Notable examples include DDT, dicofol, heptachlor, endosulfan, chlordane, aldrin, dieldrin, endrin, mirex, kepone and pentachlorophenol. These can be either hydrophilic or hydrophobic depending on their molecular structure. Many of these agents have been banned in various countries, e.g. mirex, aldrin.

Insulators. Polychlorinated biphenyls (PCBs) were once commonly used electrical insulators and heat transfer agents. Their use has generally been phased out due to health concerns. PCBs were replaced by polybrominated diphenyl ethers (PBDEs), which bring similar toxicity and bioaccumulation concerns.

History. Chlorine was discovered in 1774 by Swedish chemist Carl Wilhelm Scheele, who called it dephlogisticated marine acid (see phlogiston theory) and mistakenly thought it contained oxygen. Scheele isolated chlorine by reacting MnO₂ with HCl.

4 HCl + MnO₂ → MnCl₂ + 2 H₂O + Cl₂

Scheele observed several of the properties of chlorine. The bleeching effect on litmus and the deadly effect on insects additional to the yellow green colour and the smell similar to aqua regina. Chlorine was given its current name in 1810 by Sir Humphry Davy, who insisted that it was in fact an element.

 

CHLORINATION OF WATER

Chlorine is added directly as a gas or as a concentrated solution in water (fig. a). It produces hypochlorous acid which acts as a powerful germicide. The apparatus used is called as a chlorinator. It has a large tower with baffle plates. From the top raw water and chlorine are introduced and mixed. Water once filtered generally has a chlorine content of 0·3?0·5 ppm.

Cl₂ + H₂O → HOCl + HCl (pH > 5)

HOCl ↔ H⁺ + OCl^– (pH > 8)

The form in which chlorine exists is dependent on the pH of the solution. At pH below 5 it exists as Cl₂, between pH 5·7 it exists as HOCl and above pH 8 it exists as OCl^–

Fig. Method for Chlorination

Factors affecting the efficiency of chlorine:

• Temperature of water: The death rate of micro-organisms is proportional to temperature. The efficiency of chlorination increases with increase in temperature.

• Time of contact: The death rate of micro-organisms is directly proportional to the number of microorganisms alive. Initially it is maximum and with time it decreases.

• pH of water: A small contact period of low pH value (5?6·5) is appropriate for destroying micro-organisms.

 

Safety. Chlorine is a toxic gas that irritates the respiratory system. Because it is heavier than air, it tends to accumulate at the bottom of poorly ventilated spaces. Chlorine gas is a strong oxidizer, which may react with flammable materials.

The number of people allergic to chlorine is very small. People who are allergic to chlorine cannot drink tap water, bathe in tap water or swim in pools. Dechlorinating bath salts are used to neutralize the chlorine in bath water. Otherwise, fresh water is boiled and cooled. Breathing lower concentrations can aggravate the respiratory system, and exposure to the gas can irritate the eyes. The toxicity of chlorine comes from its oxidizing power. When chlorine is inhaled at concentrations above 30 ppm, it begins to react with water and cells, which change it into hydrochloric acid (HCl) and hypochlorous acid (HClO).

When used at specified levels for water disinfection, the reaction of chlorine with water is not a major concern for human health. Other materials present in the water may generate disinfection by-products that are associated with negative effects on human health, however, the health risk is far lower than drinking undisinfected water.

Some types of organochlorides have significant toxicity to plants or animals, including humans. Dioxins, produced when organic matter is burned in the presence of chlorine, and some insecticides, such as DDT, are persistent organic pollutants which pose dangers when they are released into the environment. For example, DDT, which was widely used to control insects in the mid 20th century, also accumulates in food chains, and causes reproductive problems (i.e., eggshell thinning) in certain bird species.

When chlorinated solvents, such as carbon tetrachloride, are not disposed of properly, they accumulate in groundwater. Some highly reactive organochlorides such as phosgene have even been used as chemical warfare agents.

However, the presence of chlorine in an organic compound does not ensure toxicity. Some organochlorides are considered safe enough for consumption in foods and medicines. For example, peas and broad beans contain the natural chlorinated plant hormone 4-chloroindole-3-acetic acid (4-Cl-IAA); and the sweetener sucralose (Splenda) is widely used in diet products. As of 2004, there were at least 165 organochlorides approved worldwide for use as pharmaceutical drugs, including the natural antibiotic vancomycin, the antihistamine loratadine (Claritin), the antidepressant sertraline (Zoloft), the anti-epileptic lamotrigine (Lamictal), and the inhalation anesthetic isoflurane.

Rachel Carson brought the issue of DDT pesticide toxicity to public awareness with her 1962 book Silent Spring. While many countries have phased out the use of some types of organochlorides such as the US ban on DDT, persistent DDT, PCBs, and other organochloride residues continue to be found in humans and mammals across the planet many years after production and use have been limited. In Arctic areas, particularly high levels are found in marine mammals. These chemicals concentrate in mammals, and are even found in human breast milk. Males typically have far higher levels, as females reduce their concentration by transfer to their offspring through breast feeding.

PUBLIC SANITATION, DISINFECTION, AND ANTISEPSIS

·        Combating putrefaction

In France (as elsewhere) there was a need to process animal guts in order to make musical instrument strings, Goldbeater’s skin and other products. This was carried out in “gut factories” (boyauderies) as an odiferous and unhealthy business. In or about 1820, the Société d’encouragement pour l’industrie nationale offered a prize for the discovery of a method, chemical or mechanical, that could be used to separate the peritoneal membrane of animal intestines without causing putrefaction. It was won by Antoine-Germain Labarraque, a 44 year-old French chemist and pharmacist who had discovered that Berthollet’s chlorinated bleaching solutions (“Eau de Javel”) not only destroyed the smell of putrefaction of animal tissue decomposition, but also retarded the decomposition process itself.

Labarraque’s research resulted in chlorides and hypochlorites of lime (calcium hypochlorite) and of sodium (sodium hypochlorite) being employed not only in the boyauderies but also for the routine disinfection and deodorisation of latrines, sewers, markets, abattoirs, anatomical theatres and morgues. They were also used, with success, in hospitals, lazarets, prisons, infirmaries (both on land and at sea), magnaneries, stables, cattle-sheds, etc.; and for exhumations, embalming, during outbreaks of epidemic illness, fever, blackleg in cattle, etc.

·        Against infection and contagion

Labarraque’s chlorinated lime and soda solutions have been advocated since 1828 to prevent infection (called “contagious infection”, and presumed to be transmitted by “miasmas”) and also to treat putrefaction of existing wounds, including septic wounds. In this 1828 work, Labarraque recommended for the doctor to breathe chlorine, wash his hands with chlorinated lime, and even sprinkle chlorinated lime about the patient’s bed, in cases of “contagious infection.” In 1828, it was well known that some infections were contagious, even though the agency of the microbe was not to be realized or discovered for more than half a century.

During the Paris cholera outbreak of 1832, large quantities of so-called chloride of lime were used to disinfect the capital. This was not simply modern calcium chloride, but contained chlorine gas dissolved in lime-water (dilute calcium hydroxide) to form calcium hypochlorite (chlorinated lime). Labarraque’s discovery helped to remove the terrible stench of decay from hospitals and dissecting rooms, and, by doing so, effectively deodorised the Latin Quarter of Paris. These “putrid miasmas” were thought by many to be responsible for the spread of “contagion” and “infection” – both words used before the germ theory of infection. The use of chloride of lime was based on destruction of odors and “putrid matter.” One source has claimed that chloride of lime was used by Dr. John Snow to disinfect water from the cholera-contaminated well feeding the Broad Street pump in 1854 London. Three reputable sources that described the famous Broad Street pump cholera epidemic do not mention Snow performing any disinfection of water from that well. Instead, one reference makes it clear that chloride of lime was used to disinfect the offal and filth in the streets surrounding the Broad Street pump—a common practice in mid-nineteenth century England.

Chloride.

The chloride ion is formed when the element chlorine picks up one electron to form an an ion (negatively-charged ion) Cl⁻. The salts of hydrochloric acid HCl contain chloride ions and can also be called chlorides. An example is table salt, which is sodium chloride with the chemical formula NaCl. In water, it dissolves into Na⁺ and Cl⁻ ions.

The word chloride can also refer to a chemical compound in which one or more chlorine atoms are covalently bonded in the molecule. This means that chlorides can be either inorganic or organic compounds. The simplest example of an inorganic covalently-bonded chloride is hydrogen chloride, HCl.

Chemistry

Hydrochloric acid fumes turning pH paper red showing that the fumes are acidic

Hydrogen chloride is composed of diatomic molecules, each consisting of a hydrogen atom H and a chlorine atom Cl connected by a covalent single bond. Since the chlorine atom is much more electronegative than the hydrogen atom, the covalent bond between the two atoms is quite polar. Consequently, the molecule has a large dipole moment with a negative partial charge δ^– at the chlorine atom and a positive partial charge δ⁺ at the hydrogen atom. In part due to its high polarity, HCl is very soluble in water (and in other polar solvents).

Upon contact, H₂O and HCl combine to form hydronium cations H₃O⁺ and chloride anions Cl^– through a reversible chemical reaction:

HCl + H₂O → H₃O⁺ + Cl^–

The resulting solution is called hydrochloric acid and is a strong acid. The acid dissociation or ionization constant, K_a, is large, which means HCl dissociates or ionizes practically completely in water. Even in the absence of water, hydrogen chloride can still act as an acid. For example, hydrogen chloride can dissolve in certain other solvents such as methanol, protonate molecules or ions, and serve as an acid-catalyst for chemical reactions where anhydrous (water-free) conditions are desired.

HCl + CH₃OH → CH₃O⁺H₂ + Cl^–

Because of its acidic nature, hydrogen chloride is corrosive, particularly in the presence of moisture.

Production

Direct synthesis. In the chlor-alkali industry, brine (mixture of sodium chloride and water) solution is electrolyzed producing chlorine (Cl₂), sodium hydroxide, and hydrogen (H₂). The pure chlorine gas can be combined with hydrogen to produce hydrogen chloride.

Cl₂ + H₂ → 2HCl

As the reaction is exothermic, the installation is called an HCl oven or HCl burner. The resulting hydrogen chloride gas is absorbed in deionized water, resulting in chemically pure hydrochloric acid. This reaction can give a very pure product, e.g. for use in the food industry.

Laboratory methods

Small amounts of HCl gas for laboratory use can be generated in a HCl generator by dehydrating hydrochloric acid with either sulfuric acid or anhydrous calcium chloride. Alternatively, HCl can be generated by the reaction of sulfuric acid with sodium chloride:

NaCl + H₂SO₄ → NaHSO₄ + HCl

This reaction occurs at room temperature. Provided there is salt remaining in the generator and it is heated above 200 degrees Celsius, the reaction proceeds to;

NaCl + NaHSO₄ → HCl + Na₂SO₄

For such generators to function, the reagents should be dry.

HCl can also be prepared by the hydrolysis of certain reactive chloride compounds such as phosphorus chlorides, thionyl chloride (SOCl₂), and acyl chlorides. For example, cold water can be gradually dripped onto phosphorus pentachloride (PCl₅) to give HCl in this reaction:

PCl₅ + H₂O → POCl₃ + 2HCl

High purity streams of the gas require lecture bottles or cylinders, both of which can be expensive. In comparison, the use of a generator requires only apparatus and materials commonly available in a laboratory.

Safety of hydrochloric acid

Hydrogen chloride forms corrosive hydrochloric acid on contact with water found in body tissue. Inhalation of the fumes can cause coughing, choking, inflammation of the nose, throat, and upper respiratory tract, and in severe cases, pulmonary edema, circulatory system failure, and death. Skin contact can cause redness, pain, and severe skin burns. Hydrogen chloride may cause severe burns to the eye and permanent eye damage.

The gas, being strongly hydrophilic, can be easily scrubbed from the exhaust gases of a reaction by bubbling it through water, producing useful hydrochloric acid as a byproduct.

Any equipment handling hydrogen chloride gas must be checked on a routine basis; particularly valve stems and regulators. The gas requires the use of specialized materials on all whetted parts of the flow path, as it will interact with or corrode numerous materials hydrochloric acid alone will not; such as stainless and regular polymers.

A simple example of an organic covalently-bonded (an organochloride) chloride is chloromethane (CH₃Cl), often called methyl chloride.

Other examples of inorganic covalently-bonded chlorides that are used as reactants are:

·        phosphorus trichloride, phosphorus pentachloride, and thionyl chloride, all three of which reactive chlorinating reagents that have been used in a laboratory

·        disulfur dichloride (S₂Cl₂), used for vulcanization of rubber.

Bromine

Bromine (pronounced /broʊmin/, /broʊmaɪn/, /broʊmɪn/, Greek: βρῶμος, brómos, meaning “stench (of he-goats)”), is a chemical element with the symbol Br and atomic number 35. A halogen element, bromine is a red volatile liquid at standard room temperature that is intermediate in reactivity between chlorine and iodine. Bromine vapours are corrosive and toxic. Approximately 730,000,000 kg was produced in 1993. The main applications for bromine are in fire retardants and fine chemicals.

History. Bromine was discovered by Antoine Balard at the salt marshes of Montpellier in 1826, but was not produced in quantity until 1860. The French chemist and physicist Joseph-Louis Gay-Lussac suggested the name bromine due to the characteristic smell of the vapors. Some also suggest that it may have been discovered by Bernard Courtois, the man who discovered iodine.

Isotopes. Bromine has 2 stable isotopes: Br-79 (50.69%) and Br-81 (49.31%). At least another 23 isotopes are known to exist. Many of the bromine isotopes are fission products. Several of the heavier bromine isotopes from fission are delayed neutron emitters. All of the radioactive bromine isotopes are relatively short lived. The longest half life is the neutron deficient Br-77 at 2.376 days. The longest half life on the neutron rich side is Br-82 at 1.471 days. A number of the bromine isotopes exhibit metastable isomers. Stable Br-79 exhibits a radioactive isomer, with a half life of 4.86 seconds. It decays by isomeric transition to the stable ground state.

Notable characteristics. Bromine is the only liquid nonmetallic element at room temperature and one of only six elements on the periodic table that are liquid at or close to room temperature. The pure chemical element has the physical form of a diatomic molecule, Br₂. It is a dense, mobile, reddish-brown liquid, that evaporates easily at standard temperature and pressures to give a red vapor (its color resembles nitrogen dioxide) that has a strong disagreeable odor resembling that of chlorine. Bromine is a halogen, and is less reactive than chlorine and more reactive than iodine. Bromine is slightly soluble in water, and highly soluble in carbon disulfide, aliphatic alcohols (such as methanol), and acetic acid. It bonds easily with many elements and has a strong bleaching action.

Certain bromine-related compounds have been evaluated to have an ozone depletion potential or bioaccumulate in living organisms. As a result many industrial bromine compounds are no longer manufactured, are being restricted, or scheduled for phasing out.

Bromine is a powerful oxidizing agent. It reacts vigorously with metals, especially in the presence of water, as well as most organic compounds, especially upon illumination.

Bromine has no known role in human health. Organobromine compounds do occur naturally, a famous example being Tyrian purple. Most organobromine compounds iature arise via the action of vanadium bromoperoxidase.

Occurrence and production. The diatomic element Br₂ does not occur naturally. Instead, bromine exists exclusively as bromide salts in diffuse amounts in crustal rock. Due to leaching, bromide salts have accumulated in sea water, but at a lower concentration than chloride. Bromine may be economically recovered from bromide-rich brine wells and from the Dead Sea waters. The bromide-rich brines are treated with chlorine gas, flushing through with air. In this treatment, bromide anion is oxidized to bromine by the chlorine gas.

2 Br⁻ + Cl₂ → 2 Cl⁻ + Br₂

2KBr + Cl₂ = 2KCl + Br₂

Because of its commercial availability and long shelf-life, bromine is not typically prepared. Small amounts of bromine can however be generated through the reaction of solid sodium bromide with concentrated sulfuric acid (H₂SO₄). The first stage is formation of hydrogen bromide (HBr), which is a gas, but under the reaction conditions some of the HBr is oxidized further by the sulfuric acid to form bromine (Br₂) and sulfur dioxide (SO₂).

NaBr (s) + H₂SO₄ (aq) → HBr (aq) + NaHSO₄ (aq)

2 HBr (aq) + H₂SO₄ (aq) → Br₂ (g) + SO₂ (g) + 2 H₂O (l)

Similar alternatives, such as the use of dilute hydrochloric acid with sodium hypochlorite, are also available. The most important thing is that the anion of the acid (in the above examples, sulfate and chloride, respectively) be more electronegative than bromine, allowing the substitution reaction to occur.

Compounds

Inorganic chemistry. Bromine is an oxidizer, and it will oxidize iodide ions to iodine, being itself reduced to bromide:

Br₂ + 2 I⁻ → 2Br⁻ + I₂

Bromine will also oxidize metals and metaloids to the corresponding bromides. Anhydrous bromine is less reactive toward many metals than hydrated bromine, however. Dry bromine reacts vigorously with aluminium, titanium, mercury as well as alkaline earths and alkali metals.

Br₂ + H₂O Û HBr + HBrO

Br₂ + 2NaOH Û NaBr + NaBrO + H₂O

BrF₃ + 2HOH ® HBrO₂ + 3HF

3HBrO₂ = 2HBrO₃ + HBr

Br₂ + 5Cl₂ + 6H₂O = 2HBrO₃ + 10HCl

2KBrO₃ ® 2KBr + 3O₂­

 

Applications. A wide variety of organobromine compounds are used in industry. Some are prepared from bromine and others are prepared from hydrogen bromide, which is obtained by burning hydrogen in bromine.

Illustrative of the addition reaction is the preparation of 1,2-Dibromoethane, the organobromine compound produced in the largest amounts:

C₂H₄ + Br₂ → CH₂BrCH₂Br

Ethylene bromide is an additive in gasolines containing lead anti-engine knocking agents. It scavenges lead by forming volatile lead bromide, which is exhausted from the engine.

The bromides of calcium, sodium, and zinc account for a sizable part of the bromine market. These salts form dense solutions in water that are used as drilling fluids.

Miscellaneous uses:

· Several dyes, agrichemicals, and pharmaceuticals are organobromine compounds. 1-Bromo-3-chloropropane, 1-bromoethylbenzene, and 1-bromoalkanes are prepared by the antimarkovnikov addition of HBr to alkenes. Ethidium bromide, EtBr, is used as a DNA stain in gel electrophoresis.

· Bromine is also used in for the production of brominated vegetable oil, which is used as an emulsifier in many citrus-flavored soft drinks.

· Water purification compounds.

· Disinfectants

Safety. Elemental bromine is toxic and causes burns. As an oxidizing agent, it is incompatible with most organic and inorganic compounds. Care needs to taken when transporting bromine, it is commonly carried in steel tanks lined with lead, supported by strong metal frames.

When certain ionic compounds containing Bromine are mixed with Potassium permanganate (KMnO₄), they will form a pale brown cloud of Bromine gas. This gas smells like bleach and is very irritating to the mucus membranes. This form of Bromine will appear to diffuse slowly, but it will suddenly disappear. Upon exposure, one should move to fresh air immediately. If symptoms arise, medical attention is needed.

Bromine is corrosive to human tissue in a liquid state and its vapors irritate eyes and throat. Bromine vapors are very toxic with inhalation.

Humans can absorb organic bromines through the skin, with food and during breathing. Organic bromines are widely used as sprays to kill insects and other unwanted pests. But they are not only poisonous to the animals that they are used against, but also to larger animals. In many cases they are poisonous to humans, too.

The most important health effects that can be caused by bromine-containing organic contaminants are malfunctioning of the nervous system and disturbances in genetic materials.

But organic bromines can also cause damage to organs such as liver, kidneys, lungs and milt and they can cause stomach and gastrointestinal malfunctioning. Some forms of organic bromines, such as ethylene bromine, can even cause cancer.

Inorganic bromines are found iature, but whereas they occur naturally humans have added too much through the years. Through food and drinking water humans absorb high doses of inorganic bromines. These bromines can damage the nervous system and the thyroid gland.

Environmental effects of bromine. Organic bromines are often applied as disinfecting and protecting agents, due to their damaging effects on microorganisms. When they are applied in greenhouses and on farmland they can easily rinse off to surface water, which has very negative health effects on daphnia, fishes, lobsters and algae.

Organic bromines are also damaging to mammals, especially when they accumulate in the bodies of their preys. The most important effects on animals are nerve damage and next to that DNA damage, which can also enhance the chances of development of cancer. The uptake of organic bromine takes place through food, through breathing and through the skin. Organic bromines are not very biodegradable; when they are decomposed inorganic bromines will consist. These can damage the nerve system when high doses are absorbed.

It has occurred in the past that organic bromines ended up in the food of cattle. Thousands of cows and pigs had to be killed in order to prevent contagion of humans. The cattle suffered from symptoms such as liver damage, loss of sight and depletion of growth, decrease of immunity, decreasing milk production and sterility and malformed children.

 

IODINE I₂

Characteristics

Structure of solid iodine

Iodine under standard conditions is a bluish-black solid. It can be seen apparently sublimating at standard temperatures into a violet-pink gas that has an irritating odor. This halogen forms compounds with many elements, but is less reactive than the other members of its Group VII (halogens) and has some metallic light reflectance.

In the gas phase, iodine shows its violet color.

Elemental iodine dissolves easily in most organic solvents such as hexane or chloroform due to its lack of polarity, but is only slightly soluble in water. However, the solubility of elemental iodine in water can be increased by the addition of potassium iodide. The molecular iodine reacts reversibly with the negative ion, generating the triiodide anion I₃⁻ in equilibrium, which is soluble in water. This is also the formulation of some types of medicinal (antiseptic) iodine, although tincture of iodine classically dissolves the element in aqueous ethanol.

Crystalline iodine

The colour of solutions of elemental iodine change depends on the polarity of the solvent. Ion-polar solvents like hexane, solution are violet; in moderately polar dichloromethane, the solution is dark crimson, and, in strongly polar solvents such as acetone or ethanol, it appears orange or brown. This effect is due to the formation of adducts.

Iodine melts at the relatively low temperature of 113.7 °C, although the liquid is often obscured by a dense violet vapor of gaseous iodine.

Production

2 HI + Cl₂ → I₂↑ + 2 HCl

I₂ + 2 H₂O + SO₂ → 2 HI + H₂SO₄

2 HI + Cl₂ → I₂↓ + 2 HCl

2 CuI₂ → 2 CuI + I₂

Disinfectant and water treatment

Elemental iodine is used as a disinfectant in various forms. The iodine exists as the element, or as the water-soluble triiodide anion I₃^– generated in situ by adding iodide to poorly water-soluble elemental iodine (the reverse chemical reaction makes some free elemental iodine available for antisepsis). In alternative fashion, iodine may come from iodophors, which contain iodine complexed with a solubilizing agent (iodide ion may be thought of loosely as the iodophor in triiodide water solutions). Examples of such preparations include:

·        Tincture of iodine: iodine in ethanol, or iodine and sodium iodide in a mixture of ethanol and water.

·        Lugol’s iodine: iodine and iodide in water alone, forming mostly triiodide. Unlike tincture of iodine, Lugol’s has a minimized amount of the free iodine (I₂) component.

·        Povidone iodine (an iodophor)

Redox reactions

In everyday life, iodides are slowly oxidized by atmospheric oxygen in the atmosphere to give free iodine. Evidence for this conversion is the yellow tint of certain aged samples of iodide salts and some organoiodine compounds. The oxidation of iodide to iodine in air is also responsible for the slow loss of iodide content in iodized salt if exposed to air. Some salts use iodate to prevent the loss of iodine.

Iodine is easily oxidized and easily reduced. Most common is the interconversion of I^– and I₂. Molecular iodine can be prepared by oxidizing iodides with chlorine:

2 I⁻ + Cl₂ → I₂ + 2 Cl⁻

or with manganese dioxide in acid solution:

2 I⁻ + 4 H⁺ + MnO₂ → I₂ + 2 H₂O + Mn²⁺

Iodine is reduced to hydroiodic acid by hydrogen sulfide and hydrazine:

I₂ + H₂S → 2 HI + 1/8 S₈

2 I₂ + N₂H₄ → 4 HI + N₂

When dissolved in fuming sulfuric acid (65% oleum), iodine forms an intense blue solution. The blue color is due to I+ 2 cation, the result of iodine being oxidized by SO₃:

2 I₂ + 2 SO₃ + H₂SO₄ → 2 I₂ ⁺ + SO₂ + 2 HSO₄⁻

Oxides of iodine

The best-known oxides are the anions, IO₃⁻ and IO₄⁻, but several other oxides are known, such as the strong oxidant iodine pentoxide.

By contrast with chlorine, the formation of the hypohalite ion (IO^–) ieutral aqueous solutions of iodine is negligible.

I₂ + H₂O H⁺ + I⁻ + HIO   (K = 2.0×10⁻¹³)

 In basic solutions (such as aqueous sodium hydroxide), iodine converts in a two stage reaction to iodide and iodate:

Iodic acid (HIO₃), periodic acid (HIO₄) and their salts are strong oxidizers and are of some use in organic synthesis. Iodine is oxidized to iodate by nitric acid as well as by chlorates:

I₂ + 10 HNO₃ → 2 HIO₃ + 10 NO₂ + 4 H₂O

I₂ + 2 ClO₃⁻ → 2 IO₃⁻ + Cl₂

Inorganic iodine compounds

Iodine forms compounds with all the elements except for the noble gases. From the perspective of commercial applications, an important compound is hydroiodic acid, used as a co-catalyst in the Cativa process for the production of acetic acid. Titanium and aluminium iodides are used in the production of butadiene, a precursor to rubber tyres.

Alkali metal salts are common colourless solids that are highly soluble in water. Potassium iodide is a convenient source of the iodide anion; it is easier to handle than sodium iodide because it is not hygroscopic. Both salts are mainly used in the production of iodized salt. Sodium iodide is especially useful in the Finkelstein reaction, because it is soluble in acetone, whereas potassium iodide is less so. In this reaction, an alkyl chloride is converted to an alkyl iodide. This relies on the insolubility of sodium chloride in acetone to drive the reaction:

R-Cl (acetone) + NaI (acetone) → R-I (acetone) + NaCl (s)

Despite having the lowest electronegativity of the common halogens, iodine reacts violently with some metals, such as aluminium:

3 I₂ + 2 Al → 2 AlI₃

This reaction produces 314 kJ per mole of aluminum, comparable to thermite’s 425 kJ. Yet the reaction initiates spontaneously, and if unconfined, causes a cloud of gaseous iodine due to the high temperature.

 

IODINE USES

As the element iodine plays important roles within the human body and human health, lists of iodine uses often start with biological uses of iodine within the human body. This leads easily onto uses of iodine in health-related contexts, such as its use as an antiseptic and disinfectant, then to other medical uses of iodine e.g. as a radiocontrast agent for medical imaging such as CT scans and X-ray imaging. There are also various other chemical uses of iodine – both historical and current. Common examples include the traditional photographic chemical silver iodide.

Iodine Uses within the Human Body

Iodine is a “trace element” meaning that modern science and medicine consider iodine to be essential for the healthy maintenance of life. The single most important use of the element iodine in animal biology – which of course includes human biology, is for the formation of two hormones in the thyroid gland (which is an endocrine gland). The hormones are thyroxine (sometimes called T4) and triiodothyronine (sometimes called T3). T4 and T3 contain four and three atoms of iodine per molecule, of hormone respectively.

Formation of Thyroxine and Triiodothyronine. The thyroid gland absorbs iodide from the blood to form these hormones from the amino acid tyrosine. They are then stored prior to release into the bloodstream in an iodine-containing protein called thyroglobulin. The production and release of hormones T3 and T4 by the thyroid gland is regulated by another hormone called thyroid stimulating hormone (TSH), which is produced by the pituitary gland located in the head.

Why do these matter? The thyroid hormones thyroxine and triiodothyronine perform several important functions including regulating the body’s basal metabolic rate (BMR), which determines the amount of energy the body uses, just to ‘tick over’. Therefore insufficient or excess quantities of these hormones can lead to weight-related as well as many other health-issues.

 

Iodine Uses as an Antiseptic and Disinfectant

·                   Disinfectant (e.g. for Water Treatment): Iodine has been used to disinfect water for almost 100 years. Use of iodine for water purifcation has advantages and disadvantages compared with use of chlorine for disinfecting water, e.g. comparisons re. convenience, effect on the taste of the water and short/long-term safety. Neither of these chemicals kills all harmful bacteria and iodine should not be used to treat water for use by anyone with an allergy to iodine, with active thyroid disease or who is or may be pregnant. Examples of iodine-based preparations used to disinfect water incl. iodine topical solution, iodine tincture, Lugol’s solution, povidone-iodine and tetraglycine hydroperiodide – some of which are better known by their commercial registered tradenames.

·                   Domestic cleaning products: Iodine is used in many household cleaning products available from well-known supermarkets.

·                   Medical / Pharmaceutical: Iodine has been used in topical disinfectant preparations for cleaning wounds (see picture above-right), sterilizing skin before surgical/invasive procedures and similar for many years.

·                   Examples of the historical uses of iodine for medical applications include its issue to military personnel in WW1 and WW2. Iodine was made available in phials (see picture below-right) and used in field hospitals.

Iodine Uses in Modern Medicine

·                   Medical / Pharmaceutical: Iodine is still used in topical medical disinfectants in modern hospitals, though generally as an ingredient within commercially prepared products in order to control the concentrations of the chemicals involved.

·                   Lugol’s Solution: An example of an iodine-based product that has been widely used in medicine is Lugol’s iodine (developed by French physician Jean Guillaume Auguste Lugol in 1829 ). This consists of 5 g iodine (I2), 10 g potassium iodide (KI) and enough distilled water to form a brown solution total volume of 100 mL. Uses of Lugol’s Solution have included testing for starches in organic compounds, as a cell stain to make cell nuclei more visible, application to the vagina and cervix during colposcopy – to distinguish normal from “suspicious” tissue – called Schiller’s Test, to stain/indicate the mucogingival junction in the mouth, to observe how a cell membrane uses osmosis and diffusion, and to help rid the animals of unwanted parasites and harmful bacteria.

·                   X-ray Radiocontrast: Radiocontrast agents are chemicals used to improve the quality and hence usefulness of images – usually of internal bodily structures – obtained using X-ray based imaging techniques e.g. Computed Tomography (CT) or Radiography (X-ray imaging). Radiocontrast agents are usually compounds of either barium or iodine.

·                   Examples of iodine-based radiocontrast agents incl. iopamidol (Isovue 370), iohexol (Omnipaque 350), ioxilan (Oxilan 350), iopromide (Ultravist 370) and iodixanol (Visipaque 320).

·                   Food supplement (nutrient): Due to the human body’s need for a certain amount of iodine – see “Iodine Uses within the Human Body” above, iodine is sometimes added to food products e.g. some tablesalts to increase the likelihood of consumers receiving sufficient iodine through their diet. Note that serious ill-health effects can also result from excessive iodine in the body; the body needs to receive an ideal or “optimum” amount of iodine- not as much as possible.

Other Uses of Iodine in Biology.

·        Lugol’s Solution: As mentioned above.

·        To test for starch: A standard test for starch uses iodine.

Because iodine is not very soluble in water the first step is to form an iodine reagent by dissolving iodine in water in the presence of potassium iodide, resulting in a linear triiodide ion complex – which is soluble and yellow/orange in colour.

To use this to test an unknown sample to find out if it contains starch simply add a drop ofthe orange triiodide ion complex to a small volume of the other sample / suspected starch (usually in solution in a test-tube or directly onto a moist surface e.g. of a potato). If starch is present in the sample it reacts with the triiodine complex to forming a product that has a deep blue/black colour. If no starch is present then there is no colour change so the yellow/organe linear triiodide ion complex is usually still visible.

Other Iodine Uses (incl. uses of iodine combined with other elements to form iodine compounds)

·                   Photographic Films: Iodine was used the manufacture of chemical compounds used in traditional photography (e.g. silver iodide which is a light sensitive material used in film). This use relates only to photographs taken using old-fashioned “film” techniques and not to modern digital photography e.g. as used to take photographs or save images using mobile telephones or webcams.

·                   Food Colouring: The organoiodine compound erythrosine (C20H6I4Na2O5) is also known as Red No.3, E127 and by other synonyms. It is used as a food colouring agent as well as in printing inks, as a biological stain, and as a dental plaque disclosing agent.

 

Toxicity of iodide ion

Excess iodine has symptoms similar to those of iodine deficiency. Commonly encountered symptoms are abnormal growth of the thyroid gland and disorders in functioning and growth of the organism as a whole. Iodides are similar in toxicity to bromides.

Excess iodine can be more cytotoxic in the presence of selenium deficiency. Iodine supplementation in selenium-deficient populations is, in theory, problematic, partly for this reason.

 

ASTATINE

Astatine (pronounced /æstətin/) is a radioactive chemical element with the symbol At and atomic number 85. It is the heaviest of the halogens.

Notable characteristics. This highly radioactive element has been confirmed by mass spectrometers to behave chemically much like other halogens, especially iodine (it would probably accumulate in the thyroid gland like iodine), though astatine is thought to be more metallic than iodine. Following the color trends of the halogens, you will notice that the elements get darker in color. Following the trends, astatine is expected to be a nearly black solid, which, when heated, sublimes into a dark, purplish vapor (darker than iodine). Astatine is expected to form ionic bonds with metals such as sodium, like the other halogens, but it can be displaced from the salts by lighter, more reactive halogens. Astatine can also react with hydrogen to form hydrogen astatide, which when dissolved in water, forms hydroastatic acid. Astatine is the least reactive of the halogens, being less reactive than iodine.

History. The existence of “eka-iodine” had been predicted by Mendeleev. Astatine (after Greek αστατος astatos, meaning “unstable”) was first synthesized in 1940 by Dale R. Corson, K. R. MacKenzie, and Emilio Segrè at the University of California, Berkeley by barraging bismuth with alpha particles. An earlier name for the element was alabamine (Ab).

The name Dakin was proposed for this element in 1937 by chemist Rajendralal De working in Dhaka.

Occurrence. Astatine occurs naturally from uranium-235 and uranium-238 decay, but because of its short half-life is only found in minute amounts.

Astatine is the rarest naturally-occurring element, with the total amount in Earth’s crust estimated to be less than 1 oz (28 g) at any given time. This amounts to less than one teaspoon of the element.

Compounds. Multiple compounds of astatine have been synthesized in microscopic amounts and studied as intensively as possible before their inevitable radioactive disintegration. While these compounds are primarily of theoretical interest, they are being studied for potential use iuclear medicine.

Isotopes. Astatine has 33 known isotopes, all of which are radioactive; the range of their mass numbers is from 191 to 223. There exist also 23 metastable excited states. The longest-lived isotope is ²¹⁰At, which has a half-life of 8.1 hours; the shortest-lived known isotope is ²¹³At, which has a half-life of 125 nanoseconds.

 

Video Halogens & Noble Gases http://www.youtube.com/watch?v=lUQpnnS1ZA4&feature=related

http://www.youtube.com/watch?v=YWaiEDK0kZE&feature=BF&list=ULwx2gSk2-sz4&index=4

 

Noble gases

The noble gases are a group of chemical elements with very similar properties: under standard conditions, they are all odorless, colorless, monatomic gases, with very low chemical reactivity. The six noble gases that occur naturally are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and the radioactive radon (Rn).

For the first six periods of the periodic table, the noble gases are exactly the members of group 18 of the periodic table. However, this no longer holds in the seventh period (due to relativistic effects): the next member of group 18, ununoctium, is probably not a noble gas. Instead, group 14 member ununquadium exhibits noble-gas-like properties.

The properties of the noble gases can be well explained by modern theories of atomic structure: their outer shell of valence electrons is considered to be “full”, giving them little tendency to participate in chemical reactions, and it has only been possible to prepare a few hundred noble gas compounds. The melting and boiling points for each noble gas are close together, differing by less than 10 °C (18 °F); consequently, they are liquids over only a small temperature range.

Neon, argon, krypton, and xenon are obtained from air using the methods of liquefaction of gases and fractional distillation. Helium is typically separated from natural gas, and radon is usually isolated from the radioactive decay of dissolved radium compounds. Noble gases have several important applications in industries such as lighting, welding, and space exploration. A helium-oxygen breathing gas is often used by deep-sea divers at depths of seawater over 180 feet (55 m) to keep the diver from experiencing oxygen toxemia, the lethal effect of high-pressure oxygen, and nitrogearcosis, the distracting narcotic effect of the nitrogen in air beyond this partial-pressure threshold. After the risks caused by the flammability of hydrogen became apparent, it was replaced with helium in blimps and balloons.

History

Noble gas is translated from the German noun Edelgas, first used in 1898 by Hugo Erdmann to indicate their extremely low level of reactivity. The name makes an analogy to the term “noble metals”, which also have low reactivity. The noble gases have also been referred to as inert gases, but this label is now deprecated as many noble gas compounds are now known. Rare gases is another term that was used, but this is also inaccurate because argon forms a fairly considerable part (0.94% by volume, 1.3% by mass) of the Earth’s atmosphere.

Properties of the Group VIII Elements

 

 

Chemical properties

Neon, like all noble gases, has a full valence shell. Noble gases have eight electrons in the outermost shell, except in the case of helium, which has two.

The noble gases are colorless, odorless, tasteless, and nonflammable under standard conditions. They were once labeled group 0 in the periodic table because it was believed they had a valence of zero, meaning their atoms cannot combine with those of other elements to form compounds. However, it was later discovered some do indeed form compounds, causing this label to fall into disuse.

Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior:

The noble gases have full valence electron shells. Valence electrons are the outermost electrons of an atom and are normally the only electrons that participate in chemical bonding. Atoms with full valence electron shells are extremely stable and therefore do not tend to form chemical bonds and have little tendency to gain or lose electrons. However, heavier noble gases such as radon are held less firmly together by electromagnetic force than lighter noble gases such as helium, making it easier to remove outer electrons from heavy noble gases.

As a result of a full shell, the noble gases can be used in conjunction with the electron configuratiootation to form the noble gas notation. To do this, the nearest noble gas that precedes the element in question is written first, and then the electron configuration is continued from that point forward. For example, the electrootation of carbon is 1s²2s²2p², and the noble gas notation is [He]2s²2p². This notation makes it easier to identify elements, and is shorter than writing out the full notation of atomic orbitals.

Compounds

Structure of XeF₄, one of the first noble gas compounds to be discovered.

The noble gases show extremely low chemical reactivity; consequently, only a few hundred noble gas compounds have been formed. Neutral compounds in which helium and neon are involved in chemical bonds have not been formed (although there is some theoretical evidence for a few helium compounds), while xenon, krypton, and argon have shown only minor reactivity. The reactivity follows the order Ne < He < Ar < Kr < Xe < Rn.

In 1933, Linus Pauling predicted that the heavier noble gases could form compounds with fluorine and oxygen. He predicted the existence of krypton hexafluoride (KrF₆) and xenon hexafluoride (XeF₆), speculated XeF₈ might exist as an unstable compound, and suggested xenic acid could form perxenate salts. These predictions were shown to be generally accurate, except XeF₈ is now thought to be both thermodynamically and kinetically unstable.

Xe + F₂ = XeF₂;

Xe + 2F₂ = XeF₄;

Xe + 3F₂ = XeF₆.

XeF₆ + H₂O = XeOF₄ + 2HF;

XeOF₄ + 2H₂O = XeO₃ + 4HF;

6XeF₄ + 12H₂O = 2XeO₃ + 4Xe + 3O₂ + 24HF

Xenon compounds are the most numerous of the noble gas compounds that have been formed. Most of them have the xenon atom in the oxidation state of +2, +4, +6, or +8 bonded to highly electronegative atoms such as fluorine or oxygen, as in xenon difluoride (XeF₂), xenon tetrafluoride (XeF₄), xenon hexafluoride (XeF₆), xenon tetroxide (XeO₄), and sodium perxenate (Na₄XeO₆). Some of these compounds have found use in chemical synthesis as oxidizing agents; XeF₂, in particular, is commercially available and can be used as a fluorinating agent. In theory, radon is more reactive than xenon, and therefore should form chemical bonds more easily than xenon does. However, due to the high radioactivity and short half-life of radon isotopes, only a few fluorides and oxides of radon have been formed in practice.

Krypton is less reactive than xenon, but several compounds have been reported with krypton in the oxidation state of +2. Krypton difluoride is the most notable and easily characterized. Compounds in which krypton forms a single bond to nitrogen and oxygen have also been characterized, but are only stable below −60 °C (−76 °F) and −90 °C (−130 °F) respectively).

Krypton atoms chemically bound to other nonmetals (hydrogen, chlorine, carbon) as well as some late transition metals (copper, silver, gold) have also been observed, but only either at low temperatures ioble gas matrices, or in supersonic noble gas jets.

The noble gases‑including helium‑can form stable molecular ions in the gas phase. The simplest is the helium hydride molecular ion, HeH⁺, discovered in 1925. Because it is composed of the two most abundant elements in the universe, hydrogen and helium, it is believed to occur naturally in the interstellar medium, although it has not been detected yet. In addition to these ions, there are many knoweutral excimers of the noble gases. These are compounds such as ArF and KrF that are stable only when in an excited electronic state; some of them find application in excimer lasers.

In addition to the compounds where a noble gas atom is involved in a covalent bond, noble gases also form non-covalent compounds. The clathrates, first described in 1949, consist of a noble gas atom trapped within cavities of crystal lattices of certain organic and inorganic substances. The essential condition for their formation is that the guest (noble gas) atoms must be of appropriate size to fit in the cavities of the host crystal lattice. For instance, argon, krypton, and xenon form clathrates with hydroquinone, but helium and neon do not because they are too small or insufficiently polarizable to be retained. Neon, argon, krypton, and xenon also form clathrate hydrates, where the noble gas is trapped in ice.

An endohedral fullerene compound containing a noble gas atom

Bonding in XeF₂ according to the 3-center-4-electron bond model

Noble gas compounds such as xenon difluoride (XeF₂) are considered to be hypervalent because they violate the octet rule. Bonding in such compounds can be explained using a 3-center-4-electron bond model. This model, first proposed in 1951, considers bonding of three collinear atoms. For example, bonding in XeF₂ is described by a set of three molecular orbitals (MOs) derived from p-orbitals on each atom. Bonding results from the combination of a filled p-orbital from Xe with one half-filled p-orbital from each F atom, resulting in a filled bonding orbital, a filled non-bonding orbital, and an empty antibonding orbital. The highest occupied molecular orbital is localized on the two terminal atoms. This represents a localization of charge which is facilitated by the high electronegativity of fluorine.

The chemistry of heavier noble gases, krypton and xenon, are well established. The chemistry of the lighter ones, argon and helium, is still at an early stage, while a neon compound is still yet to be identified.

Occurrence and production

The abundances of the noble gases in the universe decrease as their atomic numbers increase. Helium is the most common element in the universe after hydrogen, with a mass fraction of about 24%. Most of the helium in the universe was formed during Big Bang nucleosynthesis, but the amount of helium is steadily increasing due to the fusion of hydrogen in stellar nucleosynthesis. Abundances on Earth follow different trends; for example, helium is only the third most abundant noble gas in the atmosphere. The reason is that there is no primordial helium in the atmosphere; due to the small mass of the atom, helium cannot be retained by the Earth’s gravitational field. Helium on Earth comes from the alpha decay of heavy elements such as uranium and thorium found in the Earth’s crust, and tends to accumulate iatural gas deposits. The abundance of argon, on the other hand, is increased as a result of the beta decay of potassium-40, also found in the Earth’s crust, to form argon-40, which is the most abundant isotope of argon on Earth despite being relatively rare in the Solar System. This process is the base for the potassium-argon dating method. Xenon has an unexpectedly low abundance in the atmosphere, in what has been called the missing xenon problem; one theory is that the missing xenon may be trapped in minerals inside the Earth’s crust. Radon is formed in the lithosphere as from the alpha decay of radium. It can seep into buildings through cracks in their foundation and accumulate in areas that are not well ventilated. Due to its high radioactivity, radon presents a significant health hazard; it is implicated in an estimated 21,000 lung cancer deaths per year in the United States alone.

Neon, argon, krypton, and xenon are obtained from air using the methods of liquefaction of gases, to convert elements to a liquid state, and fractional distillation, to separate mixtures into component parts. Helium is typically produced by separating it from natural gas, and radon is isolated from the radioactive decay of radium compounds. The prices of the noble gases are influenced by their natural abundance, with argon being the cheapest and xenon the most expensive.

Applications

Liquid helium is used to cool the superconducting magnets in modern MRI scanners.

Noble gases have very low boiling and melting points, which makes them useful as cryogenic refrigerants.

Helium is used as a component of breathing gases to replace nitrogen, due its low solubility in fluids, especially in lipids. Gases are absorbed by the blood and body tissues when under pressure like in scuba diving, which causes an anesthetic effect known as nitrogearcosis. Due to its reduced solubility, little helium is taken into cell membranes, and when helium is used to replace part of the breathing mixtures, such as in trimix or heliox, a decrease in the narcotic effect of the gas at depth is obtained. Helium’s reduced solubility offers further advantages for the condition known as decompression sickness, or the bends. Helium is also used as filling gas iuclear fuel rods for nuclear reactors. The reduced amount of dissolved gas in the body means that fewer gas bubbles form during the decrease in pressure of the ascent. Another noble gas, argon, is considered the best option for use as a drysuit inflation gas for scuba diving.

Since the Hindenburg disaster in 1937, helium has replaced hydrogen as a lifting gas in blimps and balloons due to its lightness and incombustibility, despite an 8.6% decrease in buoyancy.

In many applications, the noble gases are used to provide an inert atmosphere. Argon is used in the synthesis of air-sensitive compounds that are sensitive to nitrogen. Solid argon is also used for the study of very unstable compounds, such as reactive intermediates, by trapping them in an inert matrix at very low temperatures. Helium is used as the carrier medium in gas chromatography, as a filler gas for thermometers, and in devices for measuring radiation, such as the Geiger counter and the bubble chamber. Helium and argon are both commonly used to shield welding arcs and the surrounding base metal from the atmosphere during welding and cutting, as well as in other metallurgical processes and in the production of silicon for the semiconductor industry.

15,000-watt xenon short-arc lamp used in IMAX projectors

Noble gases are commonly used in lighting because of their lack of chemical reactivity. Argon, mixed with nitrogen, is used as a filler gas for incandescent light bulbs. Krypton is used in high-performance light bulbs, which have higher color temperatures and greater efficiency, because it reduces the rate of evaporation of the filament more than argon; halogen lamps, in particular, use krypton mixed with small amounts of compounds of iodine or bromine. The noble gases glow in distinctive colors when used inside gas-discharge lamps, such as “neon lights“. These lights are called after neon but often contain other gases and phosphors, which add various hues to the orange-red color of neon. Xenon is commonly used in xenon arc lamps which, due to their nearly continuous spectrum that resembles daylight, find application in film projectors and as automobile headlamps.

The noble gases are used in excimer lasers, which are based on short-lived electronically excited molecules known as excimers. Some noble gases have direct application in medicine. Helium is sometimes used to improve the ease of breathing of asthma sufferers. Xenon is used as an anesthetic because of its high solubility in lipids, which makes it more potent than the usual nitrous oxide, and because it is readily eliminated from the body, resulting in faster recovery. Xenon finds application in medical imaging of the lungs through hyperpolarized MRI. Radon, which is highly radioactive and is only available in minute amounts, is used in radiotherapy.

Discharge color

 

The color of gas discharge emission depends on several factors, including the following:

·        discharge parameters (local value of current density and electric field, temperature, etc. – note the color variation along the discharge in the top row);

·        gas purity (even small fraction of certain gases can affect color);

·        color balance and saturation level of the image recording medium;

·        material of the discarge tube envelope – note suppression of the UV and blue components in the bottom-row tubes made of thick household glass.

Practical skills

Qualitative tests on VIIA group elements

1. Ca²⁺ + 2F^– = CaF₂↓ – white precipitation

2. Ag⁺ + Cl^– = AgCl↓ – white precipitation

AgCl↓ + 2NH₃^.H₂O = [Ag(NH₃)₂]⁺ + OH^– + Cl^–

3. Ag⁺ + Br^– = AgBr↓ – green-yellow precipitation

AgBr↓ + 2NH₃^.H₂O (concentrated) = [Ag(NH₃)₂]⁺ + OH^– + Br^–

4. Ag⁺ + I^– = AgI↓ – yellow precipitation

AgI↓ + 2Na₂S₂O₃ = Na₃[Ag(S₂O₃)₂] + NaI

References:

1. The abstract of the lecture.

2. intranet.tdmu.edu.ua/auth.php

3. Atkins P. W. Physical chemistry / P.W. Atkins. – New York, 1994. – P.299‑307.

4. Cotton F. A. Chemical Applications of Group Theory / F. A. Cotton. ‑ John Wiley & Sons : New York, 1990.

5. Girolami G. S. Synthesis and Technique in Inorganic Chemistry / G. S. Girolami, T. B. Rauchfuss, R. J. Angelici. ‑ University Science Books : Mill Valley, CA, 1999.

6. Russell J. B. General chemistry / J B. Russell. New York.1992. – P. 550‑599.

7. Lawrence D. D. Analytical chemistry / D. D. Lawrence. –New York, 1992. – P. 218–224.

8. http://www.lsbu.ac.uk/water/ionish.html

9. http://en.wikipedia.org/wiki

The following website shows the reaction of VIIA group elements. It’s cool stuff! Check it out!

www.youtube.com/watch?v=u2ogMUDBaf4

www.youtube.com/watch?v=mY7o28-l_WU

www.youtube.com/watch?v=cbFCWFksYoM

www.youtube.com/watch?v=gpqgvSkoXmE

www.youtube.com/watch?v=V7LcQMGbDpE

 

Prepared by PhD Falfushynska H.