Determining Charges of Polyatomic ions

[Tclack]

Does there exist a fool-proof way of doing it?

I've scoured the forums, and I've already found this question, but the answers are not satisfying.

Here for example is exactly my question:
https://www.physicsforums.com/showthread.php?t=525291&highlight=polyatomic+ions

I've already memorized a bunch of -ate's (nick the camel... and I've made my own to cover more)

HERE'S WHAT I ORIGINALLY THOUGHT:
My assumption is that you look at each atom's ion state and add them up. This works for some, but not others. Nitrate for example (NO3-).
A Nitrogen ion is -3 (or +5)
And each Oxygen is -2 (or +6)
So I can see +5 and -6 makes: -1, Which works out!

It doesn't work for Nitrite however (NO2-)
+5 plus -4 leaves a +1

So, Why is what I originally thought wrong?
and how does one determine the charges of polyatomic ions? (specifically oxyanions... for now)

 

[Borek]

There is a good new, and a bad news.

Bad news is - there is no universal method.

Good news is: it doesn't matter (much).

As you have already found, final result depends on the oxidation state on the central atom, and we can determine it knowing the polyatomic charge - so using it the other way around won't work.

However, in reality there are just a few polyatomic ions that are commonly used, so it is easier to remember them, than to look for elaborate schemes. NO₃^-, NO₂^-, SO₃^2-, SO₄^2-, PO₄^3-, CO₃^2-, perhaps BO₃^3-. All those containing chlorine or bromine will have a charge of -1. There are some tricky phosphoric acids, but they are rarely used and I always just remembered they exist and I have to check the details once I need them.

 

[Tclack]

I FOUND SOMETHING!

It's on wikipedia and it's not sourced, but...

http://en.wikipedia.org/wiki/Oxyanion#Naming

Does this make any sense? Sulfate is SO4^2−
But, let's assume we don't know that charge, and someone told us to write sulfate. So, it's not in group 7, so the central atom of S will have a 6+ charge. With 4 oxygens at -2 giving a -8 ox. number, the total is -2!

Superficially, this makes sense (I haven't exhaustedly checked every one), but if it is true, am I correct in assuming that all* oxyanions of the same central atom have the same charge?

So far I've found:
nitrate and nitrite have a -1 charge
phosphate and phosphite have a -3
perchlorate, chlorate, chlorite, hypochlorite all have -1 charge
Sulfate, sulfite, hyposulfite have a -2 charge
Arsenate, arsenite have a -3 charge
perbromate, bromate, bromite, hypobromite all have a -1 charge
Perselenate, Selenate and selenite have -2
*I assume that transition metals do their own thing, because I've found
Permanganate (MnO4^1-) and Manganate(MnO4^2-) apparently, Permanganate is formed from the Mn7+ ion and Manganate from the Mn6+ ion. Stupid transition metals ruin predictive properties.*One thing I'm trying to resolve is the determination of the -ate suffix.
There's two possible origins:
1. The naming rule from wikipedia. (Halogens get per_ate, the others get _ate if oxy #'s = grp #'s)
2. What I've always heard, that the most common oxyanion gets the -ate suffix

It seems unlikely that these are both correct. I've read #2 in at least two textbooks and I've only found #1 on wikipedia. This makes #2 more likely. However, checking a bunch of them gives merit to #1. Does anyone know of another source? I'd like to cite it if possible.

 

[Borek]

Don't blame just transition metals, throw in periodic acid with two different forms depending on pH: IO₄^- and IO₆^5-

 

[LudusRex]

I can tell you that there is indeed a fool-proof way of determining the charges of polyatomic ions. It involves understanding the concept of oxidation states and applying the rules of assigning oxidation states to each atom in the ion.

First, we need to understand that the charge of a polyatomic ion is determined by the sum of the oxidation states of all the atoms in the ion. The oxidation state of an atom is the hypothetical charge that it would have if all its bonds were ionic.

To determine the oxidation state of an atom in a polyatomic ion, we follow these rules:

1. The oxidation state of an atom in a pure element is always 0.
2. The oxidation state of a monoatomic ion is equal to its charge.
3. For oxygen, the oxidation state is usually -2, except in peroxides (such as H2O2) where it is -1.
4. For hydrogen, the oxidation state is usually +1, except in metal hydrides (such as NaH) where it is -1.
5. For fluorine, the oxidation state is always -1.
6. The sum of the oxidation states of all the atoms in a neutral molecule is 0, and in a polyatomic ion is equal to the ion's charge.

Let's apply these rules to the examples you mentioned. In nitrate (NO3-), we have three oxygen atoms with a -2 oxidation state each, and one nitrogen atom with a +5 oxidation state. This gives us a total of -6 + 5 = -1, which is the charge of the ion.

In nitrite (NO2-), we have two oxygen atoms with a -2 oxidation state each, and one nitrogen atom with a +4 oxidation state. This gives us a total of -4 + 4 = 0, which is the charge of the ion. However, since the ion has a -1 charge, it means that one of the oxygen atoms must have a different oxidation state than -2. In this case, that oxygen atom has a -1 oxidation state, giving us a total of -1 + 4 = -1, which is the charge of the ion.

So, to determine the charges of polyatomic ions, we need to follow the rules of assigning oxidation states to each atom in the ion and make sure that the sum of the oxidation states is equal to the ion's charge. This method is