Introduction:

Group 17 consisting non-metals, namely, fluorine, chlorine, bromine, iodine and astatine, collectively known as the halogens.  Halogens derive their name from the Greek words, halos + gens meaning salt producers, as they form salts in combination by metals. The most common of such salts is sodium chloride or the common salt. Halogens find out a broad variety of uses in daily life.

Occurrence, Extraction and Uses:

Chemistry of halogens is extremely interesting and has varied applications in our everyday life. Due to their high reactivity, halogens don't take place free in nature. In the combined form, though, they are fairly abundant.

Occurrence:

Fluorine is the first element of this group; it comprises around 0.054% of earth's crust, where it takes place as fluorspar (CaF₂), cryolite (AlF₃·3NaF) and fluorapatite [CaF₂-3Ca₃(PO₄)₂].

Chlorine, that forms 0.013% of earth's crust, takes place mostly as chlorides of sodium and other alkali and alkaline earth metals in salt mines and also in sea water. Sea water is nearly a 3% solution of different salts, of which sodium chloride forms ~8.3%.  Bromine, as bromides, takes place in sea water and dry salt lakes and comprises about 2.0 x 10^-4 % of earth's crust. Iodine is the rarest of all the halogens, making only 4.6 x 10^-5% of earth's crust. Its major source used to be keep, or the ash obtained on burning sea weeds, and Chile saltpetre in which it takes place as iodates. Though, now it is generally extracted from brine. Astatine is a radioactive 85e element. The naturally occurring isotopes of astatine encompass half-lives of less than one minute. Thus, it takes place in negligible amounts in nature. The isotope ²¹¹At having the longest half-life of 7.21 hours, is formed by bombarding bismuth by α-particles.

²⁰⁹₈₃Bi + ⁴₂He → ²¹¹₈₅At + 2 ¹₀n

Extraction:

The only practicable procedure of preparing fluorine gas is Moissan's original method based on the electrolysis of KF dissolved in anhydrous HF. We are familiar that chlorine is obtained as a byproduct all along with hydrogen in the manufacture of sodium hydroxide via electrolysis of brine. Electrolysis of molten sodium chloride as well provides chlorine and sodium. In several, parts of the world, it is generated by the electrolysis of aqueous HCl. Bromine is made on an industrial scale through reaction of bromides with chlorine. A mixture of air and chlorine is blown via an aqueous solution of a bromide at a pH of 3.5. Chlorine relocates bromine and air blows it out of the solution.

Iodine can as well be prepared likewise by the oxidation of iodides via chlorine. Iodine is as well made by treating brine by AgNO₃ to precipitate AgI. This is possible as AgI is the least soluble of all the silver halides. Precipitated AgI is treated by clean scrap iron or steel to form metallic silver and a solution of FeI2. The solution is then treated by chlorine to discharge N. Precipitated silver is re-dissolved in diluted HNO3 to give AgNO₃ that is used again to precipitate the AgI.

I¯ (brine) + (AgNO₃) → AgI↓ Ag ↓ + FeI₂

FeI₂ + (Cl₂) → FeCl₂ + I₂

3Ag + 4HNO → 3AgNO₃ + NO + 2H₂O

Preparation of Fluorine: 

Isolation of fluorine represented a tough problem to chemists for around a century. However its existence was first illustrated by Davy in the year 1813, yet it could not be isolated before the year 1886. All efforts at isolation of fluorine failed due to the reasons shown below:

1) Very high chemical reactivity of fluorine towards the other elements.

2) It attacked the vessels whether made up of glass, platinum, carbon or any other metal in which its preparation was tried.

3) In view of fluorine being the most prevailing oxidant, no oxidising agent could be available that could bring about the oxidation of HF to F₂. Thus, the only procedure available was that of electrolysis.

4) The procedure of electrolysis was not fruitful. Aqueous HF on electrolysis provided hydrogen and oxygen. On other hand anhydrous HF was found to be the non-conductor of electricity.

5) Very poisonous and corrosive character of anhydrous HF confirmed fatal to early chemists.

All the above stated reasons were very disheartening. Though, Moissan picked up courage and entered this field.  He electrolyzed a cooled solution of KF in the anhydrous liquid HF at 250 K by using platinum-iridium electrodes sealed by fluorspar caps in the platinum U-tube. In this reaction, the real electrolyte is KF whereas HF acts as the ionizing solvent, F₂ is evolved at anode and H, at the cathode as represented below:

KF → K⁺ + F^-

At anode,

F^- → F + e

F + F → F₂

At cathode

K⁺ + e → K

2K + 2HF → 2KF + H₂ ↑

Potassium fluoride therefore formed again experiences electrolysis. As the hydrogen fluoride is used up, more is added to prevent the melting point of the mixture from rising. The outgoing gases, F₂ and H₂, are not allowed to mix-up in the electrolytic cell. The fluorine gas is collected in the plastic receivers.

Moissan's original process has been modified. In place of expensive Pt/Ir alloy, cells are made up of copper, steel or Monel metal, which is a nickel-copper alloy, has been employed. These get covered through a thin protective film of the fluoride just as aluminium is protected via the thin film of oxide. Anode is a carbon rod impregnated by copper to render it inert and cathode is made up of steel or copper. A mixture of KF and HF in the molar ratio of 1:1 or 1:2 is employed as electrolyte providing a working temperature of 515 K or 345 K, correspondingly.

Uses:

The major use of halogens is in the halogenation of organic and inorganic compounds. You should have heard of or employed the tincture of iodine (as iodine dissolved in alcohol) as an antiseptic. Iodine is present in the thyroid hormone. A lack of iodine causes goitre and leads to the stunted growth and cretinism. To prevent this, common salt is regularly iodized. You should have utilized toothpastes having fluorides in order to prevent tooth decay by dental caries. You are aware that naturally occurring uranium is a mixture of two isotopes ²³⁸U (99.3%) and ²³⁵U (0.7%). Of the two, the later is fissionable and is employed for the generation of nuclear power. Fluorine is employed for the production of uranium hexafluoride, the compound employed for the separation of ²³⁵U and ²³⁸U isotopes by gaseous diffusion process. Besides this, liquid fluorine was employed as an oxidant in rocket fuels but this has now been discontinued. Teflon, so familiar to modern house-wife in the form of its coating on kitchenware to make them non-sticking, is a polymer of completely fluorinated ethylene. We are familiar with the use of chlorofluorocarbons or CFCs as refrigerants, in aerosol sprays and in micro-electronics. Freons, example: CCl₂F₂ (Freon-12) and CCl₃F (Freon-11) employed as refrigerants as well have fluorine. Bleaching powder, CaOCl₂, is employed for bleaching paper pulp and textiles. The bleaching powder or liquid chlorine is employed for disinfection of water on a huge scale. Chlorine was employed in the chemical warfare in World War-I. The most significant use of chlorine is in the manufacture of PVC or polyvinylchloride, which due to its non-inflammability and insulating properties is employed as an electrical insulator, for covering the electric wires, making conduit pipes and so on.

Dichlorodiphenyltrichloroethane (DDT) is employed broadly as an insecticide. Methyl bromide is the most efficient nematocide acknowledged. It is as well employed as a general pesticide. The use of silver bromide in preparing photographic plates or films is a common knowledge.

General Characteristics:

Each and every elements of Group 17 have seven electrons in their outermost shells, having configuration ns², np⁵. Therefore, they are just one electron short of the electronic configuration of noble gases. The single unpaired electron in p-orbital is responsible for the chemical bonding by other elements.

Physical Properties:

Halogens exist as the non-polar diatomic molecules which are colored. Fluorine is pale yellow, chlorine is yellowish green, bromine is brown and iodine is violet in the gaseous state. Solid iodine is approximately black having a shiny metallic lustre. Apart from iodine which consists of some helpful biological applications, halogens are extremely hazardous and toxic, fluorine being the most. Their vapors generate a choking sensation whenever inhaled.

Several physical properties of the halogens are illustrated in table shown below:

Table: Physical properties of halogens

Physical properties such as melting and boiling points are associated to the sizes and masses of the molecules and also intermolecular attraction.

The effects of mass and size that steadily increase as we go down the group, is simple to understand. As for the intermolecular attraction in non-polar homo-nuclear diatomic molecules such as halogens, which don't encompass any permanent polarity, the only forces of attraction are the weak vander Waals forces. The Polarisability of halogens rises as we go down the group, and its maximum in iodine and least in fluorine. Therefore, Van der Waals forces of attraction are maximum in iodine and least in the fluorine, with bromine and chlorine coming in between. As a consequence of this, fluorine and chlorine are gases at ordinary temperatures; bromine is a liquid and iodine a solid. This is as well reflected in the trends examined in their enthalpies of fusion and vaporization.

As we move all along a period, the effective nuclear charge rises reaching a maximum at the noble gases. Halogens that instantly precede the noble gases encompass a very high effective nuclear charge coupled by small sizes and therefore encompass the highest ionization energies in the corresponding periods, next only to the noble gases. Similar to the trend in other groups, utilisation energy of halogens as well reduces in going down the group from fluorine to iodine.

Halogens have 7 (seven) electrons in their valence shells, they encompass a very strong tendency of gaining an electron to get a stable noble gas configuration. Thus, they encompass very high electron affinities. However, their electron affinities are highest in their corresponding periods. Their electron affinity follows the order Cl > F > Br > I.

As we are familiar that as we go across the p-block elements in a period, the electronegativity increases reaching a maximum at the halogen group. Therefore, halogens are the most electronegative elements in their corresponding periods. Electronegativity reduces on moving down a group, making fluorine the most and iodine the least electronegative among the halogens.

In going down the group from chlorine to iodine, the X-X bond dissociation energy steadily reduces. This is simply illustrated by considering once again the size factor. In chlorine molecule; which is the smallest of the three, namely, Cl₂, Br₂, I₂, the two bonding electrons are nearer to both the nuclei and are held strongly, whereas in bromine and iodine, the distance of bonding electrons from the nuclei gradually raises resultant in lesser attraction and resulting weakening of the bond. Moreover, as the size of the atom increases, it yields in a less effective overlap of the orbitals and thus, progressively weaker bonds are formed as we go down the group. The bond dissociation energy rises in the order of I₂ < Br₂ < Cl₂ and if this trend was to carry on, you must anticipate the F-F bond dissociation energy to be greater than the bond energy of chlorine, 244 kJ mol^-1. However this is not so. The actual bond dissociation energy of fluorine molecule is, though, surprisingly low and consists of the value of 158 kJ mol^-1 only. The anomalously low bond dissociation energy of fluorine molecule is attributed to the fact that fluorine atom is extremely small and the non-bonding electrons on fluorine are closer to one other, resultant in a much greater lone pair-lone pair repulsion that weakens the covalent bond and lowers its dissociation energy. This repulsion is not so great in relatively bigger halogen molecules such as chlorine, bromine and iodine where the lone pairs are at a greater average distance from one other.

Oxidation States:

Fluorine is for all time univalent. As it is the most electronegative element, it always consists of the oxidation number, 1. Fluorine consists of no d-orbital in its valence shell, therefore it can't encompass any excited states or any other oxidation number.

Oxidation state of -1 is the most common and stable one for other halogens as well. Though, consistent by the decreasing electronegativity, -1 oxidation state becomes steadily less stable in going down the group. Since chlorine, bromine and iodine are less electronegative than fluorine and oxygen, they show an oxidation state of +1 in their fluorides and oxides. Moreover, apart from fluorine all the other halogens show oxidation states of +3, +5 and +7 due to the availability of vacant d-orbitals as illustrated below:

Fig: Oxidation states

Chlorine and bromine as well show oxidation state +4 (ClO₂ and BrC₂) and + 6 (Cl₂O₆ and BrO₃). Iodine shows an oxidation state + 4 in I₂O₄.

Oxidation Power:

Oxidation might be considered as the removal of electrons, in such a way that the oxidising agent gains electrons. As halogens encompass a greater tendency to pick up electrons, they act as strong oxidising agents. Their oxidising power, though, reduces on moving down the group. The strength of an oxidising agent or its capability to accept the electrons based on some energy terms. The reaction, 

1/2 X₂ (standard) + e → X^-(aq)

Representing the oxidising action of a halogen is in reality a complex process. It includes the given steps:

1/2 X₂ (standard) → ΔH_v → 1/2 X₂ (g)

1/2 X₂(g) → ΔH_d → X(g)

X(g) + e → E_A → X^-(g)

X^-(g) + aq → ΔH_hyd → X^-(aq)

The above changes are symbolized in the form of Born-Haber cycle, as illustrated below:

Fig: Born-Haber cycle

Obviously, as energy is invariably absorbed in steps (II) and (III), enthalpy of vapourisation, ΔH_v and enthalpy of dissociation, ΔH_d, always encompass positive values. Though, energy is discharged in steps (IV) and (V), therefore, electron affinity, E_A, and enthalpy of hydration, ΔH_hyd are negative. As a result, from Hess's law, the total enthalpy change, ΔE for the reduction reaction (I) is represented by the expression:

ΔE = ΔH_v + ΔH_d + E_A + ΔH_hyd

For fluorine and chlorine that are gases at room temperature, the enthalpy of evaporation is omitted. Enthalpy changes related with each of the above steps and the total enthalpy change, ΔE are given in the table above.

We can observe from the table, that ΔE, or the total enthalpy change related with the reaction (I) reduces from fluorine to iodine. Considering that the difference in entropy changes is small and mostly enthalpy changes find out the free energy change, it can be inferred that the free energy change for reaction (I) becomes less negative on descending the group. In another words, fluorine is the strongest oxidising agent of the four. Therefore, you might note that in spite of the electron affinity of chlorine being highest, fluorine is the strongest oxidising agent due to its low enthalpy of dissociation and high enthalpy of hydration.

Therefore, it is the total enthalpy change and not the electron affinity that controls the strength of an oxidising agent. 

Table: Halogen-element Bond Energy, kJ mol^-1

The oxidising power of halogens in the solid stale reactions as well exhibits the same order. This is due to the lattice energies of the ionic halides follow the similar order as the hydration energy; fluorides having the highest and iodides the lowest lattice energy.

The other significant factor that makes fluorine the strongest oxidising agent is the high element-fluorine covalent bond energy as illustrated in the table shown below:

We will notice that the values in the first row representing fluorine-element bond energy are highest apart from in the case of fluorine-fluorine bond energy. Therefore fluorine makes a very strong bond with nearly all other elements. An effect of this is that fluorine is capable to form compounds by other elements in their higher oxidation states. The order of the capability of the halogens to join with elements in higher oxidation states is F > Cl > Br > I.

Comparison of the reduction potential of halogens with that of oxygen can point out which of them would oxidize water to oxygen. The standard electrode potential for the reduction of oxygen to water in acid solution is around 1.23 volts:

1/2 O₂ + 2H⁺ + 2e → H₂O

The standard reduction potential of the halogen half cell reaction,

1/2 X₂(g) + e → X¯ (aq)

is around +2.87 volts in case of fluorine and +1.36 volts in case of chlorine. These are more positive than that for reduction of oxygen to water. Such two halogens can, thus, oxidize water to oxygen. Fluorine does so readily however chlorine reacts instead slowly, at first providing HClO that later decomposes to oxygen and HCl. The electrode potentials of bromine and iodine are less than that of oxygen, thus, they are not capable to oxidize water to oxygen.

To sum up, the major thermodynamic factors responsible for making fluorine so exclusively highly oxidising and reactive are:    

• High hydration energy of the fluoride ion.
• High lattice energy of ionic fluorides.
• Low F-F bond energy.
• High element fluorine bond energy.

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