Transition metals
So far in the syllabus, transition metals have only been mentioned briefly and you need to know more about them for your HL syllabus.
By definition, transition metals are metals that have, or form, partially filled d-orbitals. Since these electrons are present in the outer shell, they are easily lost and participate in metallic bonding.
Thus, 3d transition metals have 3d electrons and 4s electrons in their metallic structure. The general trend is that when more electrons are involved, the stronger the bond becomes. This causes the main physical properties of transition metals:
- Higher melting and boiling points than other metals.
- Higher conductivity than other metals.
The periodic trend for melting point however is not evident. It should be the case that across the period of transition metals, melting and boiling point increase as more electrons are added. While this is generally true, there are exceptions to this rule so let's review these:
- Across a period, nuclear charge increases.
- Shielding effect does not increase.
- Thus, effective nuclear charge increases across a period.
- This, combined with the increased number of delocalized valence electrons should result in a stronger metallic bond.
However, since there are a significant number of electrons in the valence shell now, there is considerable electron repulsion across the period. This results in the decrease in melting point seen in Vanadium and Cobalt to Copper.
The sudden drops in Manganese and Zinc can be explained by the half full and full d-orbital, which is so stable that these electrons do not participate in metallic bonding.
Additionally, remember that Chromium and Copper sacrifice a 4s electron to create a half-full or full d-orbital. This thus also significantly decreases melting point, but the remaining 4s electron can participate in metallic bonding.