Bond Polarity (AQA A Level Chemistry)

Electronegativity

• Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond towards itself

• The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical

• This phenomenon arises from the positive nucleus’s ability to attract the negatively charged electrons, in the outer shells, towards itself

• The Pauling scale is used to assign a value of electronegativity for each atom

First three rows of the periodic table showing electronegativity values

• Fluorine is the most electronegative atom on the Periodic Table, with a value of 4.0 on the Pauling Scale

• It is best at attracting electron density towards itself when covalently bonded to another atom

Electron distribution in the C-F bond of fluoromethane

Nuclear charge

• Attraction exists between the positively charged protons in the nucleus and negatively charged electrons found in the energy levels of an atom

• An increase in the number of protons leads to an increase in nuclear attraction for the electrons in the outer shells

• Therefore, an increased nuclear charge results in an increased electronegativity

Atomic radius

• The atomic radius is the distance between the nucleus and electrons in the outermost shell

• Electrons closer to the nucleus are more strongly attracted towards its positive nucleus

• Those electrons further away from the nucleus are less strongly attracted towards the nucleus

• Therefore, an increased atomic radius results in a decreased electronegativity

Shielding

• Filled energy levels can shield (mask) the effect of the nuclear charge causing the outer electrons to be less attracted to the nucleus

• Therefore, the addition of extra shells and subshells in an atom will cause the outer electrons to experience less of the attractive force of the nucleus

  • Sodium (period 3, group 1) has higher electronegativity than caesium (period 6, group 1) as it has fewer shells and therefore the outer electrons experience less shielding than in caesium

• Thus, an increased number of inner shells and subshells will result in a decreased electronegativity

Trends in Electronegativity

• Electronegativity varies across periods and down the groups of the periodic table

Down a group

• There is a decrease in electronegativity going down the group

• The nuclear charge increases as more protons are being added to the nucleus

• However, each element has an extra filled electron shell, which increases shielding

• The addition of the extra shells increases the distance between the nucleus and the outer electrons resulting in larger atomic radii

• Overall, there is decrease in attraction between the nucleus and outer bonding electrons

Electronegativity decreases going down the groups of the periodic table

Across a period

• Electronegativity increases across a period

• The nuclear charge increases with the addition of protons to the nucleus

• Shielding remains relatively constant across the period as no new shells are being added to the atoms

• The nucleus has an increasingly strong attraction for the bonding pair of electrons of atoms across the period of the periodic table

• This results in smaller atomic radii

Electronegativity increases going across the periods of the Periodic Table