• Elements in the periodic table are arranged in order of increasing atomic number and are placed in vertical columns (groups) and horizontal rows (periods)
• The elements across the periods show repeating patterns in chemical and physical properties
• This is called periodicity

Atomic radius

• The atomic radius is the distance between the nucleus and the outermost electron of an atom
• The atomic radius is measured by taking two atoms of the same element, measuring the distance between their nuclei and then halving this distance
• In metals this is also called the metallic radius and in non-metals, the covalent radius

Atomic radii of Period 3 elements table

• Across the period, the atomic radii decrease
• Note that radii is the plural of radius

Ionic radius

• The ionic radius is the distance between the nucleus and the outermost electron of an ion
• Metals produce positively charged ions (cations) whereas nonmetals produce negatively charged ions (anions)
• The cations have lost their valence electrons which causes them to be much smaller than their atoms
• Because there are fewer electrons, this also means that there is less shielding of the outer electrons
• Going across the period from Na⁺ to Si⁴⁺ the ions get smaller due to the increasing nuclear charge attracting the outer electrons in the second principal quantum shell nucleus (which has an increasing atomic number)
• The anions are smaller than their original atoms because each atom has gained one or more electrons in their third principal quantum shell
• This increases the repulsion between electrons while the nuclear charge is still the same
• Going across P^3- to Cl^– the ionic radii decreases as the nuclear charge increases across the period and less electrons are gained by the atoms (P gains 3 electrons, S 2 electrons and Cl 1 electron)

Ionic radii of ions of Period 3 elements table

Melting point

Melting points of the elements across Period 3 table

• The above trends can be explained by looking at the bonding and structure of the elements which are summarised in the table below

Bonding & structure of Period 3 elements table

• The table shows that 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:
  • The electrons are free to move around and are not bound to an atom 

• 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
• 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
• Because of this, the melting points increase going from Na to Al

• Si has the highest melting point due to its giant molecular structure in which each Si atom is held to its neighbouring Si atoms by strong covalent bonds

• P, S, Cl and Ar are non-metallic elements and exist as simple molecules (P₄, S₈, Cl₂ and Ar 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
• Therefore, the melting points decrease going from P to Ar (note that the melting point of S is higher than that of P as sulphur exists as larger S₈ molecules compared to the smaller P₄ molecule)

Electrical conductivity

• Electrical conductivity refers to how well a substance can conduct electricity

Trends in electrical conductivity across Period 3 table

• As mentioned before, going from Na to Al, there is an increase in the number of valence electrons that are donated to the ‘sea’ of delocalised electrons
• Because of this, there are more electrons available in Al to move around through the structure when it conducts electricity, making Al a better electrical conductor than Na
• Due to the giant molecular structure of Si, there are no delocalised electrons that can freely move around within the structure
• Si is therefore not a good electrical conductor and is classified as a semimetal (metalloid)
• The lack of delocalised electrons is also why P and S cannot conduct electricity