Electronegativity (Edexcel A Level Chemistry): Revision Note

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

Bond Polarity

• When two atoms in a covalent bond have the same electronegativity the covalent bond is nonpolar

The two chlorine atoms have the same electronegativities so the bonding electrons are shared equally between the two atoms

• The difference in electronegativities will dictate the type of bond that is formed

• When the electronegativities are very different (difference of more than 1.7) then ions will be formed and the bond will be ionic

• When two atoms in a covalent bond have a difference in electronegativities of 0.3 to 1.7 a covalent bond is formed and the bond will be polar

  • The electrons will be drawn towards the more electronegative atom

• As a result of this: 

  • The negative charge centre and positive charge centre do not coincide with each other

  • This means that the electron distribution is asymmetric

  • The less electronegative atom gets a partial charge of δ+ (delta positive)

  • The more electronegative atom gets a partial charge of δ- (delta negative)

• The greater the difference in electronegativity the more polar the bond becomes

Cl has a greater electronegativity than H causing the electrons to be more attracted towards the Cl atom which becomes delta negative and the H delta positive

Assigning polarity to molecules

• To determine whether a molecule with more than two atoms is polar, the following things have to be taken into consideration:

  • The polarity of each bond

  • How the bonds are arranged in the molecule

• Some molecules have polar bonds but are overall not polar because the polar bonds in the molecule are arranged in such way that the individual dipole moments cancel each other out

There are four polar covalent bonds in CH₃Cl which do not cancel each other out causing CH₃Cl to be a polar molecule; the overall dipole is towards the electronegative chlorine atom

Though CCl₄ has four polar covalent bonds, the individual dipole moments cancel each other out causing CCl₄ to be a nonpolar molecule

• Further examples of molecules with no net dipole:

Carbon dioxide and boron trifluoride have polar bonds but no net dipole