List of aqueous ions by element

When a salt of a metal ion, with the generic formula MX_n, is dissolved in water, it will dissociate into a cation and anions.

${\displaystyle MX_{n}\rightarrow M^{n+}(aq)+nX^{-}(aq)}$

(aq) signifies that the ion is aquated, with cations having a chemical formula [M(H₂O)_p]^q+ and anions whose state of aquation is generally unknown. For convenience (aq) is not shown in the rest of this article as the number of water molecules that are attached to the ions is irrelevant in regard to hydrolysis. This reaction occurs quantitatively with salts of the alkali-metals at low to moderate concentrations.

With salts of divalent metal ions, the aqua-ion will be subject to a dissociation reaction, known as hydrolysis, a name derived from Greek words for water splitting. The first step in this process can be written as

${\displaystyle M^{n+}+OH^{-}}$ ⇌ ${\displaystyle M(OH)^{(n-1)+}}$

When the pH of the solution is increased by adding an alkaline solution to it, the extent of hydrolysis increases. Measurements of pH or colour change are used to derive the equilibrium constant for the reaction. Further hydrolysis may occur, producing dimeric, trimeric or polymeric species containing hydroxy- or oxy- groups. The next step is to determine which model for the chemical processes best fits the experimental data.

Beryllium hydrolysis. Water molecules are omitted in this diagram

The model is defined in terms of a list of those complex species which are present in solutions in significant amounts. In the present context the complex species have the general formula [M_pO_q(OH)_r]^n±. where p, q and r define the stoichiometry of the species and n± gives the electrical charge of the ion. The experimental data are fitted to those models which may represent the species that are formed in solution. The model which gives the best fit is selected for publication. However, the pH range in which data may be collected is limited by the fact that an hydroxide with formula M(OH)_n will be formed at relatively low pH, as illustrated at the right. This will make the process of model selection difficult when monomers and dimers are formed. and virtually impossible when higher polymers are also formed. In those cases it must be assumed that the species found in solids are also present in solutions.

The formation of an hydroxo-bridged species is enthalpically favoured over the monomers, countering the unfavourable entropic effect of aggregation. For this reason, it is difficult to establish models in which both types of species are present.

The extent of hydrolysis can be quantified when the values of the hydrolysis constants can be determined experimentally. The first hydrolysis constant refers to the equilibrium

${\displaystyle M^{n+}(aq)}$ ⇌ ${\displaystyle M(OH)^{(n-1)+}(aq)+H^{+}}$

The association constant for this reaction can be expressed as

${\displaystyle K={\frac {[M(OH)]}{[M][OH]}}}$ (electrical charges are omitted from generic expressions)

Numerical values for this equilibrium constant can be found in papers concerned only with metal ion hydrolysis. However, it is more useful, in general, to use the acid dissociation constant, K_a.

${\displaystyle K_{a}={\frac {[M][H]}{[M(OH)]}}}$

and to cite the cologarithm, pK_a, of the value of this quantity in books and other publications. The two values are constrained by the relationship

log K(association) * log K(dissociation) = pK_w

pK_w refers to the self-ionization of water: pK = log (1/K) = -log(K).

Further monomeric complexes may be formed in a stepwise manner.

${\displaystyle [M(OH)_{n}]^{m+}+OH^{-}}$ ⇌ ${\displaystyle [M(OH)_{n+1}]^{(m-1)+}}$

Hydrolysed species containing two metal ions, with the general formula M₂(OH)_n, may be formed from pre-existing monomeric species. The stepwise reaction

${\displaystyle 2M(OH)^{n+}}$⇌ ${\displaystyle M_{2}(OH)_{2}^{2n+}}$

illustrates the process. An alternative stepwise reaction

${\displaystyle 2M(OH)^{n+}}$⇌ ${\displaystyle M_{2}O^{2n+}+H_{2}O}$

may also occur. Unfortunately it is not possible to distinguish between these two possibilities using data from potentiometric titrations because both of these reactions have no effect on the pH of the solution.

The concentration of a dimeric species decreases more rapidly with metal ion concentration than does the concentration of the corresponding monomeric species. Therefore, when determining the stability constants of both species it is usually necessary to obtain data from 2 or more titrations, each with a different metal salt concentration. Otherwise the stability constant non-linear least squares refinement may fail without providing the desired values, due to there being 100% mathematical correlation between the refinement parameters for the monomeric and dimeric species.

The principal problem when determining the stability constant for a polymeric species is how to select the "best" model to use from a number of possibilities. An example that illustrates the problem is shown in Baes & Mesmer, p. 119.[1]

A trimeric species must be formed from a chemical reaction of a dimer with a monomer, with the implication that the value of the stability constant of the dimer must be "known", having been determined using separate experimental data. In practice this extremely difficult to achieve. Instead, it is generally assumed that the species in solution are the same as the species that have been identified in crystal structure determinations. There is no way to establish whether or not the assumption is justified. Furthermore, species that are required as intermediaries between the monomer and the polymer may have such low concentrations as to be "undetectable".

An extreme example concerns the species with a cluster of 13 aluminium(III) ions, which can be isolated in the solid state; there must be at least 12 intermediate species in solution, which have not been characterized. It follows that the published stoichiometry of the polymeric species in solution may well be correct, but it is always possible that other species are actually present in solution. In general, the omission of intermediary species will affect the reliability of the published speciation schemes.

Some hydroxides of non-metallic elements are soluble in water; they are not included in the following table. Examples cited by Baes and Mesmer (p. 413) include hydroxides of Gallium(III), Indium(III), Thallium(III), Arsenic(III), Antimony(III) and Bismuth(III). Most hydroxides of transition metals are classified as being "insoluble" in water. Some of them dissolve, with reaction, in alkaline solution.

M(OH)_n + OH⁻ → [M(OH)_{n+1}]⁻

For some highly radioactive elements, such as astatine and radon, only tracer quantities have been experimented on. As such, unambiguous characterisation of the species they form is impossible, and so their species have been excluded from the table below. Some theoretical speculations as to what they might be are present in the literature; more information can be found at the main articles of the elements involved.

The occurrence of the different kinds of ions of the elements is shown in this periodic table:

• Aqueous solution
• Metal ions in aqueous solution

• Baes, CE; Mesmer, RE (1976). The Hydrolysis of Cations. Malabar: Krieger. pp. xvi+489. ISBN 0-89874-892-5.
• Brown, PW; Ekberg, C (2016). Hydrolysis of Metal Ions. Weinheim: Wiley‐VCH. pp. xvi+918. ISBN 978-3-527-33010-2.
• Richens, David T (1997). The Chemistry of Aqua Ions. New York: John Wiley & Sons. pp. xi+592. ISBN 978-0-471-97058-3.
• Turova, N (2011). Inorganic Chemistry in Tables. Berlin: Springer. pp. iv+157. ISBN 978-3-642-20487-6.
• Schweitzer, GK; Pesterfield, LL (2010). The Aqueous Chemistry of the Elements. Oxford: Oxford University Press. pp. x+448. ISBN 978-0-19-539335-4.
• Sanderson, Robert Thomas (1960). Chemical Periodicity. New York: Reinhold. pp. 330+illust.
• Greenwood, Norman, N.; Earnshaw, Alan (1984). "Chapter 2, Chemical Periodicity and the Periodic Table". Chemistry of the Elements (2nd ed.). Oxford: Butterworth. ISBN 0-7506-3365-4.{{cite book}}: CS1 maint: multiple names: authors list (link)