Electronegativity & Bonding of the Period 3 Elements (Cambridge (CIE) A Level Chemistry): Revision Note

Trends in Period 3 Electronegativity & Bonding

Electronegativity

• Electronegativity is the power of an element to draw the electrons towards itself in a covalent bond

• Going across the period, the electronegativity of the elements increases

Electronegativity across Period 3 table

Graph of the electronegativity of the Period 3 elements

• As the atomic number increases going across the period, there is an increase in nuclear charge

• Across the period, there is an increase in the number of valence electrons however the shielding is still the same as each extra electron enters the same shell

• As a result of this, electrons will be more strongly attracted to the nucleus causing an increase in electronegativity across the period

Bonding & structure of Period 3 elements table

Bonding and structure from Al to S

• As you move across the Periodic Table from aluminium (Al) to sulfur (S), both bonding and structure change:

  • Bonding changes from metallic (in Al) to covalent (in Si, P, S, etc.)

  • Structure changes from giant lattices to simple molecular structures

• Sodium (Na), magnesium (Mg), and aluminium (Al) are all metals

• They form a giant metallic lattice:

  • Positive metal ions are arranged in a regular lattice

  • They are surrounded by a ‘sea’ of delocalised electrons

  • These electrons come from the outer shell (valence shell) of each atom

Delocalised electrons and bond strength

• Na donates 1 electron per atom

• Mg donates 2 electrons per atom

• Al donates 3 electrons per atom

• As a result:

  • More delocalised electrons = stronger electrostatic forces between the metal ions and the electron cloud

  • Al³⁺ forms stronger metallic bonds than Na⁺, due to:

    • Higher ionic charge

    • Greater number of delocalised electrons

Electrical conductivity

• Because aluminium contributes more delocalised electrons, it has:

  • More charge carriers

  • Stronger metallic bonding

  • Therefore, aluminium is a better conductor of electricity than sodium or magnesium

A giant metallic lattice

• Si is a non-metallic element and has a giant molecular structure in which each Si atom is held to its neighbouring Si atoms by strong covalent bonds

• There are no delocalised electrons in the structure of Si which is why silicon cannot conduct electricity and is classified as a metalloid

 The giant molecular structure of silicon

Bonding from P to Ar

• Phosphorous, sulfur, chlorine and argon are non-metallic elements

  • Phosphorous, sulfur and chlorine exist as simple molecules (P₄ , S₈ , Cl₂)

  • Argon exists as single atoms

• The covalent bonds within the molecules are strong, however, between the molecules there are only weak instantaneous dipole-induced dipole forces

• It doesn’t take much energy to break these intermolecular forces

• The lack of delocalised electrons means that these compounds cannot conduct electricity

The simple molecular structure of phosphorous

The simple molecular structure of sulfur