How do the oxidation states of the elements vary in the transition series?

(N/A) The variation of oxidation states in the first transition series ($3d$-series) is as follows:
$1$. At the beginning of the series,fewer $d$-electrons are available for chemical bonding,resulting in fewer oxidation states. For example,Scandium ($Sc$,$d^1$) shows only $(+3)$,while Titanium ($Ti$,$d^2$) shows $(+2, +3, +4)$.
$2$. At the end of the series,$d$-orbitals are nearly or completely filled,leaving fewer orbitals available for bonding. For example,Zinc ($Zn$,$d^{10}$) shows only $(+2)$,and Copper ($Cu$,$d^9$) shows $(+1, +2)$.
$3$. The greatest number of oxidation states is observed in the middle of the series,where both $s$ and $d$ electrons are available for bonding. Manganese $(Mn)$ shows the widest range,from $(+2)$ to $(+7)$.
$4$. Elements in the transition series often exhibit multiple oxidation states differing by unity (e.g.,$V^{II}, V^{III}, V^{IV}, V^{V}$),unlike non-transition elements where oxidation states typically differ by two.
$5$. Down a group,higher oxidation states become more stable. For example,$Mo(VI)$ and $W(VI)$ are more stable than $Cr(VI)$.
$6$. Low oxidation states (e.g.,zero) are observed in complexes where ligands act as $\pi$-acceptors,such as in $Ni(CO)_4$ and $Fe(CO)_5$.