How does metallic bonding strength vary across the periodic table?

Metallic bonding strength generally increases across a period and decreases down a group in the periodic table.

In the periodic table, the strength of metallic bonds tends to increase as you move from left to right across a period. This is primarily due to the increase in the number of delocalised electrons per atom. As you move across a period, the number of valence electrons (which are free to move and contribute to the 'sea' of delocalised electrons) increases. These delocalised electrons are attracted to the positive metal ions, forming strong metallic bonds. For example, aluminium (in the third period) has stronger metallic bonds than sodium (in the same period) because aluminium has more delocalised electrons.

However, as you move down a group in the periodic table, the strength of the metallic bonds tends to decrease. This is because the atomic radius increases down a group, meaning the metal ions are further apart and the attraction between the positive ions and the delocalised electrons weakens. For instance, potassium (in the fourth period) has weaker metallic bonds than sodium (in the third period) due to its larger atomic radius. For more details on how atomic size affects bonding, see Atomic Radius.

It's also worth noting that the number of energy levels (shells) filled with electrons also plays a role in the strength of metallic bonds. More energy levels mean more shielding between the positive nucleus and the delocalised electrons, which weakens the metallic bond. This is another reason why metallic bonding strength decreases down a group, as atoms have more filled energy levels. Understanding the arrangement of groups and periods can further explain these changes, as discussed on Groups and Periods in the Periodic Table.

IB Chemistry Tutor Summary: In the periodic table, metallic bonding strength increases as you move from left to right across a period due to more delocalised electrons. However, it decreases as you move down a group because the atomic radius enlarges, which weakens the attraction between the electrons and metal ions. Additional electron shells also reduce bond strength by increasing electron shielding.