Some physical properties or group 17 elements
| Element | Atomic Radius (Aº) | Ionic Radius (Aº) | Ionization Energy (kJ mol-1) | Melting point (K) | Boiling point (K) | Electron affinity | Electro negativity |
|---|---|---|---|---|---|---|---|
| F | 0.72 | 1.86 | 1681 | 53 | 85 | 332.6 | 4.0 |
| Cl | 0.99 | 1.81 | 1255 | 172 | 238 | 348.5 | 3.0 |
| Br | 1.14 | 1.95 | 1142 | 266 | 332 | 324.7 | 2.8 |
| I | 1.33 | 2.16 | 1007 | 386 | 456 | 295.5 | 2.5 |
1. Atomic and ionic radii
The halogens have the smallest atomic radii in their respective periods due to maximum effective nuclear charge. Among themselves, the atomic and ionic radii increase with increase in atomic number. This is due to increase in the number of electron shells.
2. Ionization energies
The ionization energies of halogens are very high. This indicates that they have very little tendency to lose electrons. However, on going down the group from fluorine to astatine, the ionization energy decreases. This is due to gradual increase in atomic size, which is maximum for iodine. Consequently, it has the least ionization energy in family.
3. Melting and boiling points
The melting and boiling points of halogens increase with increase in atomic number down the group.
Explanation:The forces existing between these molecules are weak Van der Waals forces, which increase down the group. This is also clear from the change of state from fluorine to iodine. At room temperature, fluorine and chlorine are gases; bromine is a liquid while iodine and astatine are solids.
4. Electron affinities
(i) All these have maximum electron affinities in their respective periods. This is due to the fact that the atoms of these elements have only one electron less than the stable noble gas (ns2np6) configurations. Therefore, may have maximum tendency to accept an additional electron.
(ii) In general, electron affinity decreases from top to bottom in a group. This is due to the fact that the effect of increase in atomic size is much more than the effect of increase in nuclear charge and thus, the additional electron feels less attraction by the large atom. Consequently, electron affinity decreases. .(iii) Fluorine has unexpectedly less electron affinity than chlorine. Therefore, chlorine has the highest electron affinity in this group. The lower electron affinity of fluorine as compared to chlorine is due to very small size of the fluorine atom. As a result, there are strong inter-electronic repulsions in the relatively small 2p subshell of fluorine and thus the incoming electron does not feel much attraction. Therefore, its electron affinity is small. Thus, electron affinity among halogens varies as: F < Cl > Br > I.
Chlorine has the highest electron affinity in the periodic table.5. Electronegativity
Halogens have large electronegativity values. The values decrease down the group from fluorine to iodine because the atomic size increases and the effective nuclear charge decreases. Fluorine is the most electronegative element in the periodic table.
6. Metallic or non-metallic character
Because of very high ionization energy values, all halogens are non-metallic in character. The non-metallic character decreases as we go down the group. Therefore, the last element, iodine is a solid with a metallic lustre and forms positive ions such as I+ and I3+.
7. Color
All the halogens are colored. The color of different halogens are given below:
| Halogen | Fluorine | Chlorine | Bromine | Iodine |
|---|---|---|---|---|
| Colour | Light yellow | Greenish yellow | Reddish brown | Dark violet |
Explanation:
The color of halogens is due to the fact that their molecules absorb radiations from visible light and the outer electrons are easily excited to higher energy levels. The amount of energy required for excitation depends upon the size of the atom. Fluorine atom is the smallest and the force of attraction between the nucleus and the outer electrons is very large. As a result, it requires large excitation energy and absorbs violet 1ight (high energy) and therefore, appears pale yellow. On the other hand, iodine needs very less excitation energy and absorbs yellow light of low energy. Thus it appears dark violet. Similarly, we can explain the greenish yellow color of chlorine and reddish brown color of iodine.