Electronegativity

• Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The higher the electronegativity of an atom, the greater its attraction for bonding electrons.

• Electronegativity varies across the periodic table. Non-metals generally have high electronegativity values as compared to metals.

• The electronegativity increases from left to right along a period due to increasing number of protons in the nucleus which attracts the bonding pair of electrons more strongly.

• The electronegativity usually decreases down a group due to the increase in atomic radius and shielding effect, which weakens the attraction for the bonding pair of electrons.

• The Pauling Scale is commonly used to quantify electronegativity, defined by Linus Pauling. On this scale, fluorine has the highest electronegativity (4.0), while elements like Francium and Cesium have the lowest (around 0.7).

• Differences in electronegativity among atoms in a bond leads to polarity. In a polar covalent bond, the shared electrons are drawn closer to the more electronegative atom, causing a partial negative charge on that atom and a partial positive charge on the other.

• Large differences in electronegativity can lead to ionic bonding. This happens when one atom is so electronegative that it completely takes the bonding electrons from the other atom, creating ions.

• Polarity is important in determining how substances interact with each other - in terms of solubility, melting and boiling point etc.

• Understanding electronegativity can also provide insight on reactivity trends of elements and even help predict the type of bonding in an unknown substance.

Please note that while values and trends can guide predictions, there can be exceptions due to other factors not considered in these basic definitions.