The elements fluorine (F), chlorine(Cl), bromine(Br), iodine(I) and astatine(At) are called halogens.These elements are called the halogens from Greek hals, “salt” and gentian, “to form or generate”, because they are literally the salt formers. The halogen elements form a group of very reactive nonmetals and are quite similar to each other in their chemical properties. The first four elements are the common elements of the halogen family but astatine is a rare halogen. It is radioactive and its most stable isotope has a half-life of 8.3 hrs

Halogens exist as discrete diatomic molecules in all phases (gas, liquid or solid ). Fluorine and chlorine are gases of pale yellow and greenish-yellow colours respectively at room temperature and pressure. Bromine is a liquid of red-brown colour and iodine is a metallic-appearing shiny greyish black solid. The halogens have irritating odours, and they attack the skin. Bromine in particular causes burns that heal slowly. The outer shell of halogens have the configuration ns2 np5 (one electron short of the stable octet of the noble gases)

The ionization energy data of the halogens show that the fluorine atom holds its electrons tightly whereas the electrons are least tightly bound in iodine. The trend can be correlated with the sizes of the halogen atoms as shown in Table 5.1. The electron affinity values of halogens are large and negative, that is why halogens gain electrons readily. They have large, positive standard electrode potentials and their electronegativities are also fairly high.

PECULIAR BEHAVIOUR OF FLUORINE

1. Small size of F atom and of F- ion.

2. High first ionization energy and electronegativity.

3. Low dissociation energy of F2 molecule as compared to Cl2 and Br2.

4. Restriction of the valence shell to an octet.

5. Direct combination with inert gases.

Due to the small size of the F atom (or F- ion), there will be a better overlap of orbitals and consequently leads to shorter and stronger bonds with elements other than O, N and itself. Ionic fluorides have higher lattice energies than the other halides and these values are responsible for the insolubility of the fluorides of Ca, Mg, Ba, Sr and lanthanides in water. Due to the low dissociation energy of fluorine molecule, it is highly reactive. Other halogens react not too fast under the same conditions.

For example, CF4 and SF6. Also due to this restriction, fluorine remains restricted to the oxidation state -1. Fluorine is the only element that directly binds with noble gases such as Kr, Xe, and Rn that make up the fluoride.

However, fluorides are more stable with respect to separation into elements. Due to the valence shell being constrained to octet, many fluoro compounds appear idle.

OXIDIZING PROPERTIES

Relative Reactivities of the Halogens as Oxidizing Agents All the free halogens act as oxidizing agents when they react with metals or nonmetals. The reactant elements acquire a positive oxidation state in the compounds formed. On forming ionic compounds with metals, the halogens gain electrons and are converted to negative halide ions.
The halogens form a homologous series but fluorine differs from the other halogens in many respects which is due to:

 2Na + Cl >> 2Na Cl+ gas

The Halogens and the Noble Gases

The oxidizing power of halogens decreases with increase in atomic number. Fluorine has the highest oxidizing power and iodine the least. The order of decreasing power as an oxidizing agent is F2 > Cl2 > Br2 > I2 The oxidizing power of halogens depends upon the following factors:

1. Energy of dissociation

2. Electron affinity of atoms

3. Hydration energies of ions

4. Heats of vaporization (for Br2 and I2)

If a halogen has low energy of dissociation, a high electron affinity and higher hydration energy of its ions, it will have high oxidizing power. The oxidizing power of F2 is higher because it has low energy of dissociation and higher hydration energy of its ions. Due to the relative strength as oxidizing agents, it is possible for each free halogen to oxidize the ions of other halogens next to it in the family. Standard electrode potential measures oxidizing power.

F2 CI2 Br2 I2

Standard reduction potential.

Eo (V) +2.87 +1.36 +1.07 +0.542 X 2e 2X +→X 2+2e- 2X

Fluorine can oxidize all the halide ions to molecular halogens. (A reaction will occur if the value of E°is positive) Iodine being a weak oxidizing agent cannot oxidize chloride or bromide ion.

— o -2 -22 F + 2e 2F E =2.87V 2Cl Cl +2e F +2Cl Cl +2F → → → o E = -1.36V o E =+1.51V

The Halogens and the Noble Gases

In a similar way, the chlorine will oxidize both bromide and iodide ions, while bromine can oxidize only iodide ion.
Fluorine and chlorine can oxidize various coloured dyes to colourless substances, e.g. litmus and universal indicator can be decolourized when exposed to fluorine or chlorine. When used for bleaching, chlorine acts as an oxidizing agent.

-22 Cl +2Br Br +2Cl Br +2l I +2Br → →

COMPOUNDS OF HALOGENS

 Hydrides (hydrogen halides, HX)

All halogens react with hydrogen forming hydrides. The reaction of molecular hydrogen and fluorine is very fast and explosive. With chlorine, molecular hydrogen reacts in the presence of sunlight. Bromine and iodine react with molecular hydrogen at a higher temperature. The reaction with iodine is very slow and reversible. The direct combination is used as a preparative method only for HCl and HBr. Hydrogen fluoride and hydrogen chloride can also be obtained by the action of concentrated sulphuric acid on fluorides and chlorides, but analogous reactions with bromides and iodides result in partial oxidation of the hydrogen halide to the free halogen.

Properties of Hydrogen Halides

HF is a colourless volatile liquid whereas other hydrogen halides (HCI, HBr, HI) are colourless gases at room temperature. They give fumes in moist air. They are strong irritants.

2NaCl(s)+H SO (conc.) Na SO (aq)+2HCl(g) 2NaBr(s)+2H SO (conc.) Na SO (aq)+Br (l)+SO (g)+2H O → →

The Halogens and the Noble Gases

Chain polymers may also exist under certain conditions. Chains and rings of HF are of various sizes, some of these persist in the vapour phase as well. Some of the physical properties of hydrogen halides are given in Table 5.2.
Hydrogen fluoride attacks glass and has found applications as a nonaqueous solvent. It can be handled in Teflon ( polytetrafluoroethylene) containers or if absolutely dry, in copper or stainless- steel vessels kept under vacuum. Pure liquid HF is strongly hydrogen-bonded and is a viscous liquid. Its viscosity is less than that of water due to the absence of a three-dimensional network of H-bonds which occur in H2O. Hydrogen bonding is also responsible for the association of HF molecules in the vapour phase. Various test results indicate that gaseous HF consists of an equilibrium mixture of monomers and cyclic hexamers.

Some Physical Properties of Hydrogen Halides Property HF HCI HBr HI Melting points(°C) -83.8 -114.2 -86.9 -50.8 Boiling points (°C) 19.5 -85.0 -66.7 -35.3 Heat of fusion at M.P. (kJ/mol) 4.58 1.99 2.41 2.87 Heat of vaporization at B.P. (kJ/mol) 30.3 16.2 17.6 19.7 Heat of formation /kJ mol-1(rHf) -270.0 -92.0 -36.0 +26.0 Bond energy (kJ /mol-1) 566 431 366 299 H-X Bond length (pm) 92 128 141 160 Dissociation into elements at 1000°C (%) 0 0.014 0.5 33 Dipole moment (Debye) 1.8 1.1 0.8 0.4

Since the dipole moment of molecules decreases from HCI to HI, probably dipole-induced dipole forces play an important role in the intermolecular binding of the heavier HX molecules.

Melting points, boiling points, heats of fusion and heats of vapourization generally increase regularly from HCl to HI. The HF has much higher values for these properties due to hydrogen bonding. A very high boiling point of hydrogen fluoride is major evidence of the presence of hydrogen bonding among its molecules.The relative volatility of HCl, HBr and HI reflects the strengthening of the van der Waal’s forces due to the increasing size of halogens.

Hydrogen iodide is a strong reducing agent. In redox reactions, the hydrogen halides are oxidized to elemental halogens, e.g. In water, hydrogen halides give hydrofluoric, hydrochloric, hydrobromic and hydroiodic acids. Hydrofluoric acid is a weak acid due to limited ionization. The other three acids are very strong acids. The acidic strength increases in the order.

HF<HCl<HBr<HI

The halogens do not react directly with oxygen. With the help of some indirect methods, following oxides of group VIIA elements have been made.

5.5.2 Oxides of Halogens

The strength of the hydrogen halogen bond is very high in HF. It decreases with increasing size of the halogen atom. The bond strength is reflected in the case of dissociation of hydrogen halides at elevated temperatures.
HF, HCl, HBr and HI act as reducing agents in the following order:

FLUORINE CHLORINE BROMINE IODINE

Oxygen difluoride, OF2 Dichlorine monoxide, CI2O Brominemonoxide, Br2O Iodine tetraoxide, I2O4 Dioxygen difluoride, O2F2 Chlorine dioxide, CIO2 Bromine dioxide, BrO2 Iodine iodate, I4O9 Trioxygen difluoride Chorine hexaoxide, CI2O6 Chlorine heptaoxide.Cl2O7 Bromine trioxide, BrO3,(Br3O8) Iodine pentoxide, I2O5

Oxides of Halogens

Trioxygen Difluoride, O3F2 This oxide can be prepared when a mixture of fluorine and oxygen is subjected to electric discharge. At 363 °C, it is a dark red viscous liquid but turns to reddish-brown solid at 350 °C. On decomposition, it gives oxygen and another oxide of fluorine.

Oxides of Fluorine

O3F2 reacts with F2 in the presence of electric discharge to produce O2F2 32O F 2O F +O →
3 2O F +F 3O F → OXIDES OF CHLORINE

The oxides of chlorine are generally unstable. It is not possible to synthesize them by direct combination of the elements Cl2 and O2. They have extensive industrial use as commercial bleaching agents for wood, paper- pulp a
1. Chlorine Heptaoxide, CI2O7 CI2O7 is an anhydride of perchloric acid (HCIO4). It can be obtained at -10oC by dehydration of HCIO4 with P2O5.
2. Chlorine dioxide, CIO2 It is a pale yellow gas. It is prepared by reducing NaCIO3, with NaCl or SO2 or CH3OH in a strongly acidic solution.

CIO2 can also be prepared by the action of concentrated H2SO4 on KCIO3, This reaction is violent. To control the reaction oxalic acid should be added.

CIO2 explodes into Cl2 and O2 on warming. It is soluble in water and is stable in dark. It decomposes slowly in H2O to HCI and HCIO3. It is a paramagnetic substance. It is used as an antiseptic, for purification of water and to bleach cellulose material.

Iodine Pentoxide I2O5 It can be prepared by heating iodic acid at 240°C.
It is a white crystalline solid, stable up to 300°C. It has a polymeric structure. It is insoluble in organic solvents. It forms iodic acid with water.
It reacts with H2S, HCI and CO as an oxidizing agent. It is used for the quantitative analysis of CO.

Reactions of Chlorine with Cold and Hot NaOH

The reactions of chlorine with cold and hot NaOH are examples of “Disproportionation reactions”. A reaction in which a species (molecule, atom or ion)is simultaneously oxidized and reduced, is called a “disproportionation reaction”. In cold (15°C) state chlorine will react with NaOH (aq) to form hypochlorite and a halide.

2NaOH(aq)+Cl (g) NaCl(aq)+NaClO(aq)+H O(aq) (a) → Sod.hypochlorite

The reaction is a disproportionation reaction, because the zero oxidation state of the chlorine atom in Cl2, is converted to -1 in chloride and +1 in hypochlorite.

Sodium hypochlorite which is produced in the cold state in the above reaction decomposes forming sodium chloride and sodium chlorate at 70°C.