Polyatomic ion
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A polyatomic ion, also known as a molecular ion, is a covalently bonded set of two or more atoms, or of a metal complex, that can be considered to behave as a single unit and that has a net charge that is not zero. Unlike a molecule, which has a net charge of zero, this chemical species is an ion. (The prefix poly- carries the meaning "many" in Greek, but even ions of two atoms are commonly described as polyatomic.)
In older literature, a polyatomic ion may instead be referred to as a radical (or less commonly, as a radical group). (In contemporary usage, the term radical refers to various free radicals, which are species that have an unpaired electron and need not be charged.)
A simple example of a polyatomic ion is the hydroxide ion, which consists of one oxygen atom and one hydrogen atom, jointly carrying a net charge of −1; its chemical formula is OH−. In contrast, an ammonium ion consists of one nitrogen atom and four hydrogen atoms, with a charge of +1; its chemical formula is NH+4.
Polyatomic ions often are useful in the context of acid–base chemistry and in the formation of salts.
Often, a polyatomic ion can be considered as the conjugate acid or base of a neutral molecule. For example, the conjugate base of sulfuric acid (H2SO4) is the polyatomic hydrogen sulfate anion (HSO−4). The removal of another hydrogen ion produces the sulfate anion (SO2−4).
Nomenclature of polyatomic anions[]
There are two "rules" that can be used for learning the nomenclature of polyatomic anions. First, when the prefix bi is added to a name, a hydrogen is added to the ion's formula and its charge is increased by 1, the latter being a consequence of the hydrogen ion's +1 charge. An alternative to the bi- prefix is to use the word hydrogen in its place: the anion derived from H+ + CO2−3, HCO−3, can be called either bicarbonate or hydrogen carbonate.
Most of the common polyatomic anions are oxyanions, conjugate bases of oxyacids (acids derived from the oxides of non-metallic elements). For example, the sulfate anion, SO2−4, is derived from H2SO4, which can be regarded as SO3 + H2O.
The second rule is based on the oxidation state of the "main" atom in the ion, which in practice is often (but not always) directly related to the number of oxygen atoms in the ion, following the pattern shown below. Consider the chlorine oxyanion family:
| oxidation state | −1 | +1 | +3 | +5 | +7 |
|---|---|---|---|---|---|
| anion name | chloride | hypochlorite | chlorite | chlorate | perchlorate |
| formula | Cl− | ClO− | ClO−2 | ClO−3 | ClO−4 |
| structure | 
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As the chlorine atom's oxidation state becomes more positive, the number of oxygen atoms attached increases. This gives rise to the following common pattern: first, think of the -ate ion as being the "base" name; adding a per- prefix adds an oxygen, while changing the -ate suffix to -ite will reduce the oxygens by one, and keeping the suffix -ite and adding the prefix hypo- reduces the number of oxygens by one more -- all without changing the charge. The naming pattern follows within many different oxyanion series based on a standard root for that particular series. The -ite has one less oxygen than the -ate, but different -ate anions might have different numbers of oxygen atoms.
These rules do not work with all polyatomic anions, but they do apply to several of the more common ones. The following table shows how these prefixes are used for some of these common anion groups.
| bromide | hypobromite | bromite | bromate | perbromate |
| Br− |
BrO− |
BrO− 2 |
BrO− 3 |
BrO− 4 |
| Iodide | hypoiodite | Iodite | Iodate | periodate |
| I− |
IO− |
IO− 2 |
IO− 3 |
IO− 4 or IO5− 6 |
| sulfide | hyposulfite | sulfite | sulfate | persulfate |
| S2− |
S 2O2− 2 |
SO2− 3 |
SO2− 4 |
SO2− 5 |
| selenide | selenite | selenate | ||
| Se2− |
Se 2O2− 2 |
SeO2− 3 |
SeO2− 4 |
|
| telluride | tellurite | tellurate | ||
| Te2− |
TeO2− 2 |
TeO2− 3 |
TeO2− 4 |
|
| nitride | hyponitrite | nitrite | nitrate | |
| N3− |
N 2O2− 2 |
NO− 2 |
NO− 3 |
|
| phosphide | hypophosphite | phosphite | phosphate | |
| P3− |
H 2PO− 2 |
PO3− 3 |
PO3− 4 |
PO3− 5 |
| arsenide | arsenite | arsenate | ||
| As3− |
AsO3− 2 |
AsO3− 3 |
AsO3− 4 |
Other examples of common polyatomic ions[]
The following tables give additional examples of commonly encountered polyatomic ions. Only a few representatives are given, as the number of polyatomic ions encountered in practice is very large.
| Tetrahydroxyborate | B(OH)−4 |
| Acetylide | C2−2 |
| Ethoxide or ethanolate | C2H5O− |
| Acetate or ethanoate | CH3COO− or C2H3O−2 |
| Benzoate | C6H5COO− or C7H5O−2 |
| Citrate | C6H5O3−7 |
| Carbonate | CO2−3 |
| Oxalate | C2O2−4 |
| Cyanide | CN− |
| Chromate | CrO2−4 |
| Dichromate | Cr2O2−7 |
| Bicarbonate or hydrogencarbonate | HCO−3 |
| Hydrogen phosphate | HPO2−4 |
| Dihydrogen phosphate | H2PO−4 |
| Hydrogen sulfate or bisulfate | HSO−4 |
| Manganate | MnO2−4 |
| Permanganate | MnO−4 |
| Azanide or amide | NH−2 |
| Peroxide | O2−2 |
| Hydroxide | OH− |
| Bisulfide | SH− |
| Thiocyanate | SCN− |
| Silicate | SiO2−4 |
| Thiosulfate | S2O2−3 |
| Onium ions | Carbenium ions | Others | |||
|---|---|---|---|---|---|
| Guanidinium | C(NH2)+3 | Tropylium | C7H+7 | Mercury(I) | Hg2+2 |
| Ammonium | NH+4 | Triphenylcarbenium | (C6H5)3C+ | ||
| Phosphonium | PH+4 | cyclopropenium | C3H+3 | ||
| Hydronium | H3O+ | ||||
| Fluoronium | H2F+ | ||||
| Pyrylium | C5H5O+ | ||||
See also[]
External links[]
- Ions