Electronegativity

In this lesson, we will learn:

To define the term electronegativity.
To understand the relationship between electronegativity and type of (ionic/covalent) bonding.
To understand the relationship between electronegativity and bond polarity.
To be able to explain the trend in electronegativity in the periodic table.

Notes:

Electronegativity is the ability of an atom to attract bonding electrons to its outer shell. It is measured using the Pauling scale – fluorine is highest at 4.0 on the scale, the most electronegative element, whilst francium is the lowest at 0.7 and is the least electronegative element.
The general trend of electronegativity in the periodic table shows the most electronegative elements in the top-right of the table, and the bottom-right corner has the least electronegative, or most electropositive elements.

Electronegativity is the bridge between the opposite types of bonding: covalent and ionic. Whether a bond is ionic or covalent is determined by the gap in electronegativity of the elements

With a very large gap in electronegativity between atoms, a highly ionic bond will form. The very electronegative element (e.g. fluorine, oxygen) will have a much greater ability to attract bonding electrons into its outer shell than the less electronegative element (e.g. Na, K), which effectively ‘donates’ any outer shell electrons it holds. This forms an ionic bond, or more accurately, a bond with large ionic character.
With a small gap in electronegativity between atoms, a covalent bond will form. The gap means neither atom will ‘win’ the pair of electrons, they will instead share both at the same time – which is what a covalent bond is!

high ionic character or largely covalent character

Bonds between atoms with a large gap in electronegativity are called polar bonds

A gap in electronegativity between molecules is what creates polar bonds, but not all compounds with polar bonds are polar compounds!
For a molecule to be polar, it must have an asymmetric arrangement of polar bonds to create a dipole moment.

Take carbon dioxide, CO₂ for example. While the individual C=O bonds are polar, the bonds are symmetrical, arranged facing in the opposite direction to each other. This leads to their permanent dipoles cancelling each other out and leaves the overall molecule non-polar.
Dichloromethane (CH₂Cl₂) is tetrahedral shaped, having two polar bonds (C-Cl) and two non-polar bonds (C-H).
The permanent dipoles of the C-Cl bonds are not cancelled out by the two C-H bonds so a ‘dipole moment’ exists, where a - exists towards the two C-Cl bonds. Because of this, we call dichloromethane a polar molecule.

Electrostatic theory explains the trends in electronegativity in the periodic table:

Going from left to right across the periodic table, electronegativity increases. This is because each further element has one extra proton in its nucleus (increasing overall nuclear charge) and one extra outer shell electron that it is attracting (increasing negative charge while shielding stays the same). This greater charge difference produces a stronger electrostatic force and therefore the ability to attract electrons to/in the outer shell gets stronger when going from left to right in the table.
Going down any group in the periodic table, electronegativity decreases.
This is because despite each element down the group having more protons, each further element has an entire extra shell of electrons, so the outer shell and the nucleus has one more ‘in-between’ shell of electrons shielding them. This much greater shielding effect means the nucleus’s ability to attract electrons to the outer shell gets poorer as you go down a group in the periodic table.

Looking at the trend more carefully, the effect of electron shielding down a group is more influential than the increased nuclear charge across a period, so oxygen is the second most electronegative element (around 3.5 on the Pauling scale), followed by nitrogen and chlorine (both around 3.0).