Q:

Nomenclature of MnO4^-: is it an oxyanion? 

Having a real pain with nomenclature.  Look at this substance:
$$\ce{MnO4-}$$
So, do we have an anion of manganese and oxygen, or is that just an anion of oxygen?  Anyway, my book says that $\ce{MnO4-}$ is "permanganate". 
Since it ends with "-ate", it seems my book thinks $\ce{MnO4-}$ is an oxyanion (because those are the ones that end with things like "-ate" and "-ite"). 
However, it was my understanding that an oxyanion is when you have oxygen and a non-metal.  But manganese is a metal!  What is going on here then? 
And not just that.  Let's assume that $\ce{MnO4-}$ is indeed an oxyanion called permanganate.  From what I've learned, if you remove one oxygen atom (getting  $\ce{MnO3-}$), it should become just manganate, right?  But, according to my book, manganate is $\ce{MnO2-}$. It had to lose two oxygen atoms and not one to become manganate.  Why? 

A:

As far as I know, an oxoanion just refers to an anion which has oxygen covalently bonded to the central atom; the central atom can be pretty much anything apart from the Group 1 and 2 metals, since the only oxygen compounds those form are ionic oxides/peroxides/superoxides.  For the 3d transition metals, you have $\ce{VO3^-}$ (vanadate(V)), $\ce{CrO4^{2-}}$ (chromate(VI)), $\ce{MnO4-}$ (manganate(VII)), and $\ce{FeO4^{2-}}$ (ferrate(VI)), amongst others. 
I would not recommend reading too much into the nomenclature, because a lot of it is just what chemists are comfortable with after using for a long time - it's like a language that you just get more familiar with over time.  But there are some general rules.  The 'per...ate' name, such as perchlorate $\ce{ClO4-}$, implies a higher oxidation state than the normal '...ate' compound, which is chlorate $\ce{ClO3-}$ in this case.  For chlorine, only certain oxidation states are available, namely the -1, 0, +1, +3, +5 and +7 oxidation states.  Apart from 0, which occurs in diatomic chlorine $\ce{Cl2}$, you will notice that these are in increments of 2 - which means that you can generate the whole list of chlorine anions by adding one more oxygen, but keeping the charge of the anion constant: $\ce{Cl-}$, $\ce{ClO-}$, $\ce{ClO2-}$, $\ce{ClO3-}$, $\ce{ClO4-}$. This tendency to form compounds 2 oxidation states apart is very pronounced for the p-block elements, i.e. Group 15, 16 and 17. 
However, for d-block elements such as $\ce{Mn}$, this isn't the case because single d electrons are easily lost and gained.  So after permanganate, $\ce{MnO4-}$, which features $\ce{Mn}$ in the +7 oxidation state, the oxoanion with the next highest oxidation state is the manganate ion $\ce{MnO4^2-}$ which has $\ce{Mn}$ in the +6 oxidation state. ( As far as I know the ion $\ce{MnO2-}$ is not called "manganate".)