Electronegativity And Bond Polarity In Covalent Molecules
Electronegativity is the ability of a covalently bonded atom to attract the bond pair of electrons towards itself.
The greater the electronegativity value of an atom, the greater the power with which it attracts the electrons in a covalent bond towards itself.
Electronegativity depends on:
- the nuclear charge
- the distance between the nucleus and the outer shell electrons
- the shielding of the nuclear charge by electrons in inner shells.
The smaller the atom, the closer the nucleus is to the shared outer main level electrons and the greater its electronegativity. The larger the nuclear charge ([or a given shielding effect), the greater the electronegativity.
For Groups 1 to 17 (Group I to VII) the pattern of electronegativity is:
- electronegativity increases across a period from Group 1 to Group 17
- electronegativity increases up each group.
What this means is that fluorine is the most electronegative element known since it is on top of Group 17. Since electronegativity values increase across a period from Group 1 to Group 17, non-metals are more electronegative than metals.
The order of electronegativity of the most reactive metals, in order of increasing electronegativity, is:
Carbon and hydrogen have electronegativities that are lower than those of most other non-metallic elements.
Pauling scale of electronegativity
The Pauling scale is used as a measure or electronegativity on a scale of 0 to 4. The greater the number on the Pauling scale, the more electronegative the atom is. Noble gases have no electronegativity value because they do not, in general, form covalent bonds.
Bond polarity in molecules with two atoms
In a covalent bond, an element with the highest electronegativity attract bonding pair electrons towards itself with the following result:
- the electron distribution becomes asymmetric
- the two atoms are partially charged with the less electronegative atom having a partial charge δ+ (delta positive) and the more electronegative atom having a partial charge δ– (delta negative)
This type of covalent bond is said to be polar (or we can say it has a dipole).
However, if the electronegativity values of the covalently bonded atoms are the same, the bond pair electrons are equally shared. This kind of covalent bond is non-polar. For example, hydrogen (H2), chlorine (Cl2) and bromine (Br2) are non-polar molecules.
The diagram shows the polar bond in a hydrogen chloride molecule. Chlorine is more electronegative than hydrogen and it attracts the bond pair electrons towards itself.
The greater the difference in the electronegativity values of the atoms in a covalent bond, the more polar the bond becomes. The degree of polarity of a molecule is measured as a dipole moment.
The direction of the dipole is shown by the arrow sign with crossed tail. The arrow points to the partially negatively charged end of the dipole.
Bond polarity in molecules with more than two atoms
In molecules containing more than two atoms, we have to take the following into account:
- the polarity of each bond
- the arrangement of the bonds in the molecule.
Trichloromethane, CHCl3, is a polar molecule because the combined effect of three C-Cl dipoles pointing in a similar direction is not cancelled out by the polarity of the C-H bond. This is because the C-Cl bond is polar but the C-H bond is virtually non-polar and therefore the electron distribution is asymmetric. Thus, the CHCl3 molecule is polar, with the negative end towards the chlorine atoms.
There are molecules that contain polar bonds but with no overall polarity due to the dipole moments cancelling each other out. One such molecule is tetrachloromethane, CCl4. It has four polar C-Cl bonds pointing towards the four corners of a tetrahedron due to this arrangement, the dipoles in each bond cancel each other, making tetrachloromethane a non-polar molecule.