Electronegativity & Bonding of the Period 3 Elements (CIE A Level Chemistry)

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

Electronegativity of the Period 3 elements increases from Na to Cl

• 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

• The table shows that going from Al to S the bonding changes from metallic to covalent and the structure changes from giant to simple structure
• Na, Mg and Al are metallic elements which form positive ions arranged in a giant lattice in which the ions are held together by a ‘sea’ of delocalised electrons around them
• Since Al donates three electrons into the sea of delocalised electrons to form an ion with +3 charge, the electrostatic forces between the electrons and the aluminium ion will be very strong
• The electrons in the ‘sea’ of delocalised electrons are those from the valence shell of the atoms
• Na will donate one electron into the ‘sea’ of delocalised electrons, Mg will donate two and Al three electrons
• As a result of this, the metallic bonding in Al is stronger than in Na
• This is because the electrostatic forces between a 3+ ion and the larger number of negatively charged delocalised electrons are much larger compared to a 1+ ion and the smaller number of delocalised electrons in Na
• Since there are more electrons in a metallic lattice of aluminium compared to sodium and magnesium, aluminium is a better electrical conductor

A giant metallic lattice

Metal cations form a giant lattice held together by electrons that can freely move around

• 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

The diagram shows the giant molecular structure of silicon where silicon atoms are held together by strong covalent bonds

• 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 diagram shows the simple molecular structure of phosphorus with covalent bonds between the atoms

The simple molecular structure of sulfur

The diagram shows the simple molecular structure of sulfur with covalent bonds between the atoms