Trends in Physical Properties of Group 17 Elements

  • During the period 1811-1886, the halogens were discovered. They are non-metallic elements such as: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I) and astatine (At) which are the members of the halogen family.
  • This family constitutes the group 17 of the periodic table.
  • The word halogens have been derived from the Greek word “Halo” that means sea salt and “gen” means producer.
  • Among them, the last member At is radioactive in nature. The isotopes of At are short lived that is why its chemistry is not well known to much.
  • None of the halogens occur in the free elemental form in nature. This is due to their high reactivity.
  • Chlorine is the most found halogens. Due to their high electro-negativities, halogens form negative ions and invariably are found in nature as fluorides, chlorides, bromides and iodides.
  • These salts occur in sea water and are present in vast salt-beds, which have probably been formed by the evaporation of the salt-water.
  • Some important physical properties of various members of this family are given below:

i) Physical state

  • Fluorine and chlorine exist as gases, bromine as a liquid and iodine as a solid.
  • The strength of intermolecular forces increases as one goes down the group.
  • This is the reason that there is a gradual change from gaseous state to solid state as we go down the group.

ii) Atomic size and ionic radii

a) Halogens have the smallest atomic radii in their respective period which is due to the highest effective nuclear charge in the halogen atoms.

b) The atomic radii and the radii of the uni-negative ion increase in going from fluorine to iodine. Thus, the atomic and ionic radii follow the order

Atomic radius:         I    >   Br    >    Cl   >    F

Ionic radius:             I  >    Br  >   Cl  >     F

  • The atomic and ionic radii increase in the going from F to I (i.e., down the group) because at each element down the group, a new electronic shell is added to the atom/ion.

iii) Melting and boiling points

  • The halogens are low melting and low boiling. The melting and boiling points of halogens increase in going from F to I.
  • Thus, the melting and boiling points of halogens are as follow:

F  <   Cl  <  Br   <  I

Halogen exists as diatomic molecules. The forces that exist between the molecules are weak van der Waals’ type. That is why the lower members are gases. The strength of these intermolecular forces goes on increasing while going from F to I in the group. This is why the melting and boiling points of these elements increase in going from top to the bottom of the group.

iv) Dissociation energy

  • The energy change when 1 mole of halogen molecules are broken into atoms in the gaseous phase is called the dissociation energy, or atomization energy.
  • Thus, the dissociation energy of X2 is the energy change in the reaction,

X2 (g) → 2X (g)    ΔHdiss  

ΔHdiss is equal to the dissociation energy of X2.

a) All halogens have very low dissociation energies. Fluorine has unusually low dissociation energy. The dissociation energies of some diatomic molecules are given below:

Element:                              F2           Cl2        Br2            I2         H2           O2            N2

Dissociation energy:     159        243        193            151      458        495         941

(kJ mol-1)

  • The dissociation energy decreases while going from chlorine to iodine. However, fluorine has dissociation energy much lower than that of chlorine.
  • The lower dissociation energy of F2 molecule in comparison to Cl2 molecule is due to its small atomic size. F atom is very small. So, F – F bond length is small in F2 molecule. As a result, the charge density and the repulsion between the non-bonding electrons on fluorine atoms is very high as compared to that in Cl2, Br2 and I2 molecules.
  • This makes the F atoms in F2 molecules repel each other, and make the dissociation of F2 molecule into F atoms easier.

v) Enthalpy of hydration of X ions

  • The enthalpy of hydration of an anion X is the enthalpy change for the reaction,

X+ nH2O → X(aq)

The hydration energy of the halide ions becomes smaller as we go down the group.

  • The hydration enthalpy is inversely related to the ionic radius. In going from down the group, the ionic radii increase. So, the hydration enthalpy becomes smaller.

vi) Electronegativity

  • Halogens have very high electro-negativities.
  • Fluorine is the most electronegative element. The electronegativity of the halogens decrease in going from fluorine to iodine i.e., the electronegativity follows the order

F  >  Cl   >  Br  >  I

→→→→→→Electronegativity decreases→→→→→→→→

  • High values of electro-negativities for halogens are because of their small atomic radii and high effective nuclear charge.
  • Decrease in the electronegativity in going from fluorine to iodine is because of the increase in the size of the atoms while going from F to I.

vii) Ionization energy

  • Halogens have very high ionization energies. This is due to their small atomic size and high effective nuclear charge. The ionization energy decreases as we go down the group i.e., in going from F to I.
  • This is because of an increase in the atomic radii in going from fluorine to iodine.

viii) Electron affinity

  • The electron affinity values of halogens are given below:

Halogen:                        Fluorine    chlorine   Bromine   Iodine

Electron affinity:               -333         -348         -324         -295

(kJ mol-1)

a) Electron affinity values of halogens are highly negative. This means energy is released when an electron is added to an atom of these elements.

Halogens have very high electron affinities because of the following reasons:

  1. Atoms of all halogens require only one electron to attain the stable configurations of the nearest noble gases. Completely filled orbitals are more stable.
  2. The atomic radii of halogens are small.
  3. In any period, the effective nuclear charge of halogens is the highest.

b) Fluorine has unexpectedly smaller electron affinity than chlorine.

  • Fluorine atom is very small. So, the electron charge density on fluorine atom is large. Therefore, the electron experiences more repulsive force from the electron cloud of the F atom that comes to the atom.
  • As a result the net energy released during the reaction, F (g) + e → F(g), becomes slightly lesser, and the electron affinity of fluorine becomes smaller than that of chlorine.

c) The electron affinity becomes smaller in going from Cl to I. Since the electron affinity of F is smaller than that of Cl, hence Cl has the highest electron affinity.

  • When we go from Cl to I, the atomic size increase. So, there is reduction in the force with which the added electron is attracted towards the atom.
  • This leads to the lowering of electron affinity as we go from Cl to I.

ix) Metallic or non-metallic character

  • Halogens are highly electronegative elements. So, these elements are typical non-metals.
  • However, the non-metallic character goes on decreasing while going down the group from F to At. But, there is no evidence of any metallic behavior even for astatine (At).

x) Oxidation states

  • Halogens attain stable configuration of the nearest noble gas by acquiring one electron.
  • This may be done either by gain of one electron or by sharing of one electron pair with some other atom.
  • Fluorine is the most electronegative element and shows only -1 oxidation state.
  • Other halogens show both the negative and positive oxidation states which are given below:

Fluorine    -1

Chlorine  -1, +1, +3, +5, +7

Bromine  -1, +1,+3, +5

Iodine -1, +1, +3, +5, +7

Bromine and chlorine also exhibit the oxidation states of +4 and +6 in oxo-acids and oxides.

  • Fluorine does not show higher or variable oxidation states because it does not have d-orbitals in its valence shell.
  • Due to the availability of the d-orbitals in higher halogens, the paired electrons in s- and p-orbitals can be unpaired by promoting them to the vacant d-orbitals. This opens the possibilities of variable oxidation states.

xi) Colour

  • All halogens are coloured i.e., Fluorine is light yellow, Chlorine is greenish yellow, Bromine is reddish brown and iodine is violet (purple).
  • Thus, the colour of halogens gets darkened as one goes from F to I.
  • When molecules or atoms absorb energy (may be in the form of light), its outer electrons get excited to higher energy levels. The amount of energy or wavelength of light absorbed depends upon the nature of the molecule. When the atoms or molecules absorb light in the visible region, the complementary colour is emitted out. Then, the light absorbing material appears to have that complementary colour.
  • Halogens absorb light in the visible region. Fluorine absorbs the blue light and emits pale – yellow light. So, fluorine appears light yellow in colour.
  • Iodine absorbs the yellow light, so it appears violet in visible light.




Trends in Physical Properties of Group 17 Elements