Variability and relative stability of oxidation states
In this essay the concentration will be on the oxidation states of the transition metals rather than the entire periodic table.
An oxidation state is a number assigned to an element in a compound according to some rules. This number enables us to describe oxidation-reduction reactions, and balancing redox chemical reactions. Compared to Group II elements such as calcium, transition elements form ions with a wide variety of oxidation states. Calcium ions typically don't lose more than two electrons, whereas transition metals can lose up to nine.
The energies required to remove electrons from calcium are low until you try to remove electrons from below its outer two s orbitals. In fact Ca3+ has an ionisation enthalpy so high that it rarely occurs naturally. However a transition element like vanadium has roughly linear increasing ionisation enthalpies throughout its s and d orbitals, due to the close energy difference between the 3d and 4s orbitals. Transition metal ions are therefore commonly found in very high states.
Certain patterns can be seen to emerge across the period of transition elements:
The number of oxidation states of each ion increases up to Mn, after which they start to drop. This drop is due to the stronger pull from the protons in the nucleus towards the electrons, making them harder to remove.
When the elements are in lower oxidation states, they can be found as simple ions. However elements in higher oxidation states are usually bonded covalently to electronegative compounds such as O or F, often in an anion.
Properties with respect to the stability of oxidation states:
Higher oxidation state ions become less stable across the period.
Ions in higher oxidation states tend to make good oxidising agents, whereas elements in low oxidation states become reducing agents.
The 2+ ions across the period start as strong reducing agents, and become more stable.