Electronegativity is the relative tendency of a bonded atom to attract electrons to itself.
An atom with extremely low electronegativity, like a Group 1 (IA) metal, is said to be electropositive since its tendency is to lose rather than to gain, or attract, electrons
Electronegativity decreases down a Group in the Periodic Table as the atomic radius and number of inner electron shells both increase.
Electronegativity increases across a Period of the Periodic Table, in general, due to increasing nuclear charge and decreasing atomic radius.
For the commonly encounted atoms in High School Science, the order in decreasing electronegativity is
F > O > N ~ Cl > Br > C ~ S ~ I > P ~ H > Si
1 H 2.1
The number on top of the element's symbol is its atomic number, Z (the number of protons in the nucleus). In the middle of the box there is a letter (or letters), the symbol for the element. The number underneath the symbol is the electronegativity of the element using Pauling's method.
Identical non-metallic atoms have identical electronegativities and form non-polar covalent bonds since the bonding electrons will be shared equally between the atoms in the molecule.
When atoms of similar, but different, electronegativities (a difference < ~1.7) bond, the more electronegative atom has a greater share of the bonding electrons than the less electronegative atom. The more electronegative atom has a partial negative charge, and the less electronegative atom has a partial positive charge. The resulting covalent bond is called a polar covalent bond.
Non-metals are more electronegative than metals. When an extremely electronegative atom, like fluorine, bonds with an electropositive atom, like sodium, the resulting bond is ionic due to the huge difference in electronegativity (difference > ~1.7). The electronegative atom's pull on the bonding electrons is so strong that it pulls the bonding electron off the electropositive atom resulting in two oppositely charged ions which are held together by electrostatic attraction (an ionic bond).
Non-polar covalent bonds
Identical non-metallic atoms have identical elctronegativity
H-H both H atoms have the same electronegativity so the bond is non-polar
Different atoms with same electronegativity
N-Cl Both N and Cl have an electronegativity of 3.0 so the bonding electrons will be shared equally between the two atoms resulting in a non-polar covalent bond.
Polar covalent bonds
O-H O has an electronegativity of 3.5, H has an electronegativity of 2.1.
The difference in electronegativity is 3.5 - 2.1 = 1.4
1.4 is less than 1.7, so the resulting bond is polar covalent.
Oxygen is the more electronegative so it will have a greater share of the bonding electrons and therefore a partial negative charge, Oδ- Hydrogen is less electronegative so it will have a lesser share of the bonding electrons and therefore a partial positive charge, Hδ+ Since the bond has two 'poles' or 'ends' it is sometimes referred to as a dipole.
The polar covalent bond can be represented as:
δ-O - Hδ+
O - H ←
Where the arrow head points towards the most electronegative atom.
Na - F Na has an electronegativity of 0.9 and F has an electronegativity of 4.0.
The difference in electronegativity is 4.0 - 0.9 = 3.1
3.1 is greater than 1.7, so the resulting bond is ionic.
F has such a strong attraction for electrons that it pulls the electron off the Na resulting in a negative charge for fluorine, F-, and a positive charge for sodium, Na+ The bond between Na+ and F- is NOT covalent since the bonding electrons are not shared between the 'atoms'. Rather the bond is ionic since there has been a transfer of electrons from the least electronegative atom to the more electronegative atom resulting in two ions of opposite charge.
The best representation of this ionic compound is: