Defining 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