Metal ions in aqueous solution

Metal ions in aqueous solution

A metal ion in aqueous solution is a cation, dissolved in water, of chemical formula [M(H2O)n]z+. The solvation number, n, determined by a variety of experimental methods is 4 for Li+ and Be2+ and 6 for elements in rows 3 and 4 of the periodic table. Lanthanide and actinide aqua ions have solvation number of 8 and 9. The strength of the bonds between the metal ion and water molecules in the primary solvation shell increases with the electrical charge, z, on the metal ion and decreases as its radius, r, increases. Aqua ions are subject to hydrolysis. The logarithm of the first hydrolysis constant is proportional to z2/r for most aqua ions. The aqua ion is associated, through hydrogen bonding with other water molecules in a secondary solvation shell. Water molecules in the first hydration shell exchange with molecules in the second solvation shell and molecules in the bulk liquid. The residence time of a molecule in the first shell varies among the chemical elements from about 100 ps to more than 200 y. Aqua ions are prominent in electrochemistry.


Introduction to metal aqua ions

The metallic elements
Li Be
Na Mg Al
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po
Fr Ra Ac
(La) Ce Pr Nd Pm  Sm Eu Gd Tb Dy Ho Er Tb Yb Lu
(Ac) Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

Most chemical elements are metallic and form simple aqua ions with the formula [M(H2O)n]z+ when the oxidation state is 1, 2 or 3. With the higher oxidation states the simple aqua ions dissociate losing hydrogen ions to yield complexes that contain both water molecules and hydroxide or oxide ions, such as the vanadium(IV) species [VO(H2O)5]2+. In the highest oxidation state only oxyanions, such as the permanganate(VII) ion, MnO4- are known.[1]

Schematic representation of the aqua ion [Na(H2O)6]+. The oxygen atoms are arranged at the vertices of a regular octahedron centered on the sodium ion
First and second solvation shells of an octahedral aqua ion. Up to 12 water molecules may be present in the second shell (only two are shown in this diagram) linked by hydrogen bonds to the molecules in the first shell.

In aqueous solution the water molecules directly attached to the metal ion are said to belong to the first coordination sphere, also known as the first, or primary, solvation shell. The bond between a water molecule and the metal ion is a dative covalent bond, with the oxygen atom donating both electrons to the bond. Each coordinated water molecule may be attached by hydrogen bonds to other water molecules. The latter are said to reside in the second coordination sphere. The second coordination sphere is not a well defined entity for ions with charge 1 or 2. In dilute solutions it merges into the water structure in which there is an irregular network of hydrogen bonds between water molecules.[2] With tripositive ions the high charge on the cation polarizes the water molecules in the first solvation shell to such an extent that they form strong enough hydrogen bonds with molecules in the second shell to form a more stable entity.[3]

The strength of the metal-oxygen bond can be estimated in various ways. The hydration enthalpy, though based indirectly on experimental measurements, is the most reliable measure. The scale of values is based on an arbitrarily chosen zero, but this does not affect differences between the values for two metals. Other measures include the M-O vibration frequency and the M-O bond length. The strength of the M-O bond tends to increase with the charge and decrease as the size of the metal ion increases. In fact there is a very good linear correlation between hydration enthalpy and the ratio of charge squared to ionic radius, z2/r.[4] For ions in solution Shannon's "effective ionic radius" is the measure most often used.[5]

Water molecules in the first and second solvation shells can exchange places. The rate of exchange varies enormously, depending on the metal and its oxidation state. Metal aqua ions are always accompanied in solution by solvated anions, but much less is known about anion solvation than about cation solvation.[6]

Understanding of the nature of aqua ions is helped by having information on the nature of solvated cations in mixed solvents[7] and non-aqueous solvents, such as liquid ammonia, methanol, dimethyl formamide and dimethyl sulphoxide to mention a few.[8]

Occurrence in nature

Aqua ions are present in most natural waters.[9] Na+, K+, Mg2+ and Ca2+ are major constiuents of seawater.

Aqua ions in seawater (Salinity = 35)
Ion Concentration
/mol kg−1
Na+ 0.469
K+ 0.0102
Mg2+ 0.0528
Ca2+ 0.0103

Many other aqua ions are present in seawater in concentrations ranging from ppm to ppt.[9] The concentrations of sodium, potassium, magnesium and calcium in blood are similar to those of seawater. Blood also has lower concentrations of essential elements such as iron and zinc. Sports drink is designed to be isotonic and also contains the minerals which are lost in perspiration.

Magnesium and calcium ions are common constituents of domestic water and are responsible for permanent and temporary hardness, respectively. They are often found in mineral water.

Experimental methods

Information obtained on the nature of ions in solution varies with the nature of the experimental method used. Some methods reveal properties of the cation directly, others reveal properties that depend on both cation and anion. Some methods supply information of a static nature, a kind of snapshot of average properties, others give information about the dynamics of the solution.

Nuclear magnetic resonance (NMR)

Ions for which the water-exchange rate is slow on the NMR time-scale give separate peaks for molecules in the first solvation shell and for other water molecules. The solvation number is obtained as a ratio of peak areas. Here it refers to the number of water molecules in the first solvation shell. Molecules in the second solvation shell exchange rapidly with solvent molecules, giving rise to a small change in the chemical shift value of un-coordinated water molecules from that of water itself. The main disadvantage of this method is that it requires fairly concentrated solutions, with the associated risk of ion-pair formation with the anion.

Solvation numbers determined by NMR[10]
ion number nucleus
Be2+ 4 1H 17O
Mg2+ 6 1H
Zn2+ 6 1H
Al3+ 6 1H 17O
Ga3+ 6 1H 17O
In3+ 6 1H
Fe2+ 6 17O
Co2+ 6 1H
Ni2+ 6 1H 17O
Th4+ 9 1H

X-ray diffraction (XRD)

A solution containing an aqua ion does not have the long-range order that would be present in a crystal containing the same ion, but there is short-range order. X-ray diffraction on solutions yields a radial distribution function from which the coordination number of the metal ion and metal-oxygen distance may be derived. With aqua ions of high charge some information is obtained about the second solvation shell.[11][12]

This technique requires the use of relatively concentrated solutions. X-rays are scatted by electrons, so scattering power increases with atomic number. This makes hydrogen atoms all but invisible to X-ray scattering.

Large angle X-ray scattering has been used to characterize the second solvation shell with trivalent ions such as Cr3+ and Rh3+. The second hydration shell of Cr3+ was found to have 13 ± 1 molecules at an average distance of 402 ± 20 pm. This implies that every molecule in the first hydration shell is hydrogen bonded to two molecules in the second shell.[13]

Neutron diffraction

Diffraction by neutrons also give a radial distribution function. In contrast to X-ray diffraction, neutrons are scatted by nuclei and there is no relationship with atomic number.[14] Indeed, use can be made of the fact that different isotopes of the same element can have widely different scattering powers. In a classic experiment, measurements were made on four nickel chloride solutions using the combinations of 58Ni, 60Ni, 35Cl and 37Cl isotopes to yield a very detailed picture of cation and anion solvation.[15] Data for a number of metal salts show some dependence on the salt concentration.

Cation hydration determined by neutron diffraction[16]
Cation Anion Molality
/mol kg-1 (a)
θ /deg (b)
Li+ Cl- 27.77 2.3 (2)(c) 75 (5)
9.95 3.0 (5) 52 (5)
3.57 5.5 (3) 40 (5)
Ca2+ Cl- 4.49 6.4 (3) 34 (9)
2.80 7.2 (2) 34 (9)
1.0 10.0 (6) 38 (9)
Ni2+ Cl- 3.05 5.8 (2) 42 (8)
0.85 6.6 (5) 27 (10)
0.46 6.8 (8) 17 (10)
0.086 6.8 (8) 0 (20)
Ni2+ ClO4- 3.80 5.8 (2) 42 (8)
Cu2+ ClO4- 2.00 4.9 (3) 38 (6)
Fe3+ NO3- 2.0 5.0 (2) 22 (4)
Nd3+ Cl- 2.85 8.5 (2) 24 (4)
Dy3+ Cl- 2.38 7.4 (5) 17 (3)
(a) moles solute per kg of solvent. (b) angle between M-O and the H-m lines, m= mid point between H or D atoms. (c) Figures in brackets are standard deviations on the last sgnificant figure of the value.

Most of these data refer to concentrated solutions in which there are very few water molecules that are not in the primary hydration spheres of the cation or anion, which may account for some of the variation of solvation number with concentration even if there is no contact ion-pairing. The angle θ gives the angle of tilt of the water molecules relative to a plane in the aqua ion. This angle is affected by the hydrogen bonds formed between water molecules in the primary and secondary solvation shells.

The measured solvation number is a time-averaged value for the solution as a whole. When a measured primary solvation number is fractional there are two or more species with integral solvation numbers present in equilibrium with each other. This also applies to solvation numbers that are integral numbers, within experimental error. For example, the solvation number of 5.5 for a lithium chloride solution could be interpreted as being due to presence of two different aqua ions with equal concentrations.

[Li(H2O)6]+ is in equilibrium with [Li(H2O)5]+ + H2O

Another possibility is that there is interaction between a solvated cation and an anion, forming an ion-pair.This is particularly relevant when measurements are made on concentrated salt solutions. For example, a solvation number of 3 for a lithium chloride solution could be interpreted as being due to the equilibrium

[Li(H2O)4]+ + Cl- is in equilibrium with [Li(H2O)3Cl] + H2O

lying wholly in favour of the ion-pair.

Vibrational spectra

Infrared spectra and Raman spectra can be used to measure the M-O stretching frequency in metal aqua ions. Raman spectroscopy is particularly useful because the Raman spectrum of water is weak whereas the infrared spectrum of water is intense. Interpretation of the vibration frequencies is somewhat complicated by the presence, in octahedral and tetrahedral ions, of two vibrations, a symmetric one measured in the Raman spectrum and an anti-symmetric one, measured in the infrared spectrum.

Symmetric M-O stretching vibrations of some aqua ions in solution[17]
metal ion wavenumber /cm−1
Be2+ 530-543
Mg2+ 360-365
Mn2+ 395
Fe2+ 389
Ni2+ 405
Cu2+ 440
Zn2+ 385-400
Hg2+ 380[18]
Al3+ 520-526
Ga3+ 475
In3+ 400

Although the relationship between vibration frequency and force constant is not simple, the general conclusion that can be taken from these data is that the strength of the M-O bond increases with increasing ionic charge and decreasing ionic size. The M-O stretching frequency of an aqua ion in solution may be compared with its counterpart in a crystal of known structure. If the frequencies are very similar it can be concluded that the coordination number of the metal ion is the same in solution as it is in a compound in the solid state.

Dynamic methods

Data such as conductivity, mobility and diffusion relate to the movement of ions through as solution. When an ion moves through a solution it tends to take both first and second solvation shells with it. Hence solvation numbers measured from dynamic properties tend to be much higher that those obtained from static properties.

Hydration numbers measured by dynamic methods[19]
Li+ Na+ Cs+ Mg2+ Ca2+ Ba2+ Zn2+ Cr3+ Al3+
Transport number 13-22 7-13 4 12-14 8-12 3-5 10-13
Ion mobility 3-21 2-10 10-13 7-11 5-9 10-13
Diffusion 5 3 1 9 9 8 11 17 13

Solvation numbers and structures

Alkali metals

Solvation numbers of 3 - 6 have been found for lithium aqua ions. Solvation numbers less than 4 may be the result of the formation of contact ion-pairs. The higher solvation numbers may be interpreted in terms of water molecules that approach [Li(H2O)4]+ through a face of the tetrahedron, though molecular dynamic simulations may indicate the existence of an octahedral aqua ion.[20] There are most probably six water molecules in the primary solvation sphere of the sodium ion. For potassium, rubidium and cesium even the primary solvation shell is not well defined.[21]

Alkaline earth metals

Molecular hydration data for group 2 metal ions[22]
M2+(aq) M-O distance /pm –ΔHStrikeO.png /kJ mol−1 ΔVStrikeO.png /cm3 mol−1
Be 167 2494 -12.0
Mg 209 1921 -21.2
Ca 242 1577 -17.9
Sr 263 1443 -18.2
Ba 281 1305 -12.5

[Be(H2O)4]2+ has a very well-defined primary solvation shell; The BeO4 part has tetrahedral symmetry, Td.[23] [Mg(H2O)6]2+ is also a well-characterized species, with an octahedral core. The situation for calcium is more complicated. Neutron diffraction data gave a solvation number for calcium chloride, CaCl2, which is strongly dependent on concentration: 10.0 ± 0.6 at 1 mol dm−3, decreasing to 6.4± 0.3 at 2.8 mol dm−3. The Shannon radius of the 6-coordinate calcium ion is 100 pm compared to 72 pm for magnesium, a 28% increase in size. This allows for higher aquation numbers and the lower charge density (z2/r) results in weaker M-O bonds, which makes ion-pairing easier. Various solid hydrates are known with 8-coordination in square antiprism and dodecahedral geometry.[24]

Aluminium and Group 3 metals

By convention aluminium is placed in group 13, with gallium, indium and thallium. Nevertheless a comparison of aluminium and scandium aqua ions illustrates a trend between the congeners Na+/K+, Mg2+/Ca2+ and Al3+/Sc3+ which depends on the increase in size on going from row 3 to row 4 in the periodic table. Shannon radii for 6-coordinate Al3+ and Sc3+ are 54 and 74.5 pm. Ga3+ has a Shannon radius of 62 pm, only about 13% larger than that of Al3+. This is due to the presence of the ten elements between scandium and gallium, which make an additional contribution to the general decrease in size across all rows of the periodic table. The aluminium aqua ion, [Al(H2O)6]3+ is very well characterized in solution and the solid state. The AlO6 core has octahedral symmetry, point group Oh.

There is now a substantial body of indirect evidence that seven molecules of water are present in the aqua ion of scandium, Sc3+.[25] Yttrium(III) has approximately the same Shannon radius as holmium(III) so Y3+ has very similar properties to those of Ho3+, so its aqua ion is probably 8-coordinate. The Lanthanum aqua ion is probably 9-coordinate, as are those of the lighter lanthanides.

Transition metals and group 12 metals

Jahn-Teller distorted octahedral structure of [Cu(H2O)6]2+

The ions of these metals in the +2 and +3 oxidation states have a solvation number of 6. All have a regular octahedral structure except the aqua ions of chromium(II) and copper(II) which are subject to Jahn-Teller distortion. In the copper case the two axial Cu−O distances are 238 pm, whereas the four equatorial Cu−O distances are 195 pm.[26] Silver(I) is probably 4-coordinate, [Ag(H2O)4]+.[27]

A solvation number of 6 with an octahedral structure is well established for zinc(II) and cadmium(II) in dilute solutions. In concentrated solutions the Zn2+ ion may adopt a 4-coordinate, tetrahedral, structure, but the evidence is not conclusive because of the possibility of ion-pairing and/or hydrolysis.[28] The solvation number of mercury(II)) is most likely to be 6.[29]

The bis aqua structure of the mercury(I) ion,[(H2O)-Hg-Hg-(OH2]+, found in solid compounds, is also found in solution.[30] Another aqua species in which there is a metal-metal bond is the molybdenum(II) species formulated as [(H4O)-Mo-Mo-(OH4]4+ on the basis of theoretical considerations of the nature of the Mo-Mo bond. Each molybdenum is surrounded by four water molecules in a square-planar arrangement, in a structure similar to that of the known structure of the chloro complex [Mo2Cl8]4-. The presence of a fifth water molecule in the axial position is not ruled out, however.[31]

There are a few divalent and trivalent aqua ions of transition metals in the second and third transition series. With oxidation state 4, however, only hydrolyzed species exist.

Group 13-15 metals

The aqua ions of gallium(III) and indium(III) and thallium(III)have a solvation number of 6. The aqua ion of thallium(I) is often assumed to be 6-coordinate, but this assumption is not based on strong experimental evidence.[32] The Shannon radius of Tl+, at 150 pm, is not very different from that of K+, at 138 pm, so some similarity between Tl+ and K+ chemistry is both expected and observed.[33]

The solvation number of the tin(II) aqua ion, [Sn(H2O)n]2+ is not known with any certainty due to the presence of hydrolysis in the concentrated solutions needed for X-ray scattering measurements.[34] The same is true for the lead(II) aqua ion. With bismuth(III) there is indirect evidence for a solvation number of 9. A nonahydrate is been characterized in the solid state, with a face-capped trigonal prism structure. The Shannon radius for 9-coordinate bismuth (115 pm) is comparable to that of neodymium (116.3 pm) for which a solvation number of 9 is well-established.[35]

Lanthanides and actinides

face-capped trigonal prism structure

The trivalent lanthanide ions decrease steadily in size from lanthanum to lutetium, an effect known as the lanthanide contraction. Nevertheless there is strong evidence that the hydration number changes from 9 to 8 at around gadolinium.[36] Structures of aqua ions in the solid state include the 9-coordinate tricapped trigonal prism structure with the lighter lanthanide ions and the 8-coordinate square antiprism structure with the heavier lanthanide ions.There are no experimental data for the solvation state of cerium(IV) or europium(II).[37]

Solvation numbers of 9 or more a believed to apply the actinide ions in the +3 and +4 oxidation states, but an experimental value is available only for thorium(IV).


Some elements in oxidation states higher than 3 form stable, aquated, oxo ions. Well known examples are the vanadyl(IV) and uranyl(VI) ions. They can be viewed as particularly stable hydrolysis products in a hypothetical reaction such as

[V(H2O)6]4+ → [VO(H2O)5]2+ + 2H+

The vanadium has a distorted octahedral environment (point group C4v) of one oxide ion and 5 water molecules.[38] Vanadium(V) is believed to exist as the dioxo-ion [VO2(H2O)4]+ at pH less than 2, but the evidence for this ion depends on the formation of complexes, such as oxalate complexes which have been shown to have the VO2+ unit, with cis-VO bonds, in the solid state.[39] The chromium(IV) ion [CrO(H2O)5]2+ , similar to the vanadium ion has been proposed on the basis of indirect evidence.[40]

The uranyl ion, UO22+, has a trans structure. The aqua ion UO22+(aq) has been assumed, on the basis of indirect evidence, to have 5 water molecules in the plane perpendicular to the O-U-O axis in a pentagonal bipyramid structure, point group D5h. However it is also possible that there are 6 water molecules in the equatorial plane, hexagonal bipyramid, point group D6h, as many complexes with this structure are known.[41] The solvation state of the plutonyl ion, PuO22+, is not known.


The main goal of thermodynamics in this context is to derive estimates of single-ion thermodynamic quantities such as hydration enthalpy and hydration entropy. These quantities relate to the reaction

Mz+ (gas) + solvent → Mz+ (in solution)

The enthalpy for this reaction is not directly measurable, because all measurements use salt solutions that contain both cation and anion. Most experimental measurements relate to the heat evolved when a salt dissolves in water, which gives the sum of cation and anion solvation enthalpies. Then, by considering the data for different anions with the same cation and different cations with the same anion, single ion values relative to an arbitrary zero, are derived.

Minus hydration enthalpy for (octahedral) divalent transition metal M2+ ions[42]
Hydration enthalpies of trivalent lanthanide Ln3+ ions[42]
Single ion standard hydration enthalpy /kJ mol−1[42]
... Ga3+
... In3+
... Tl3+

Other values include Zn2+ -2044.3, Cd2+ -1805.8 and Ag+ -475.3 kJ mol−1.

There is an excellent linear correlation between hydration enthalpy and the ratio of charge squared, z2, to M-O distance, reff.[43]

ΔHStrikeO.png = - 69500 z2 / reff

Values for transition metals are affected by crystal field stabilization. The general trend is shown by the magenta line which passes through Ca2+, Mn2+ and Zn2+, for which there is no stabilization in an octahedral crystal field. Hydration energy increases as size decreases. Crystal field splitting confers extra stability on the aqua ion. The maximum crystal field stabilization energy occurs at Ni2+. The agreement of the hydration enthalpies with predictions provided one basis for the general acceptance of crystal field theory.[44]

The hydration enthalpies of the trivalent lanthanide ions show an increasingly negative values at atomic number increases, in line with the decrease in ionic radius known as the lanthanide contraction.

Single ion hydration entropy can be derived. Values are shown in the following table. The more negative the value, the more there is ordering in forming the aqua ion. It is notable that the heavy alkali metals have rather small entropy values which suggests that both the first and second solvation shells are somewhat indistinct.

Single ion standard hydration entropy at 25°C /J deg−1 mol−1[45]
... Ga3+
... In3+

Hydrolysis of aqua ions

There are two ways of looking at an equilibrium involving hydrolysis of an aqua ion. Considering the dissociation equilibrium

[M(H2O)n]z+ - H+is in equilibrium with [M(H2O)n-1(OH)](z-1)+

the activity of the hydrolysis product, omitting the water molecules, is given by

{[M(OH)](z − 1) + } = K1, − 1{Mz + }{H + } − 1

The alternative is to write the equilibrium as a complexation or substitution reaction

[M(H2O)n]z+ +OH- is in equilibrium with :[M(H2O)n-1(OH)](z-1)+ + H2O

In which case

{[M(OH)](z − 1) + } = K1,1{Mz + }{OH }

The concentration of hydrogen and hydroxide ions are related by the self-ionization of water, Kw = {H+} {OH-} so the two equilibrium constants are related as

K_{1,-1} = K_{1,1} \times K_w

In practice the first definition is more useful because equilibrium constants are determined from measurements of hydrogen ion concentrations. In general,

[Mx(OH)y] = βx, − y * [M]x[H] y

charges are omitted for the sake of generality and activities have been replaced by concentrations. β * are cumulative hydrolysis constants.

Modeling the hydrolysis reactions that occur in solution is usually based on the determination of equilibrium constants from potentiometric (pH) titration data. The process is far from straightforward for a variety of reasons.[46] Sometimes the species in solution can be precipitated as salts and their structure confirmed by X-ray crystallography. In other cases, precipitated salts bear no relation to what is postulated to be in solution, because a particular crystalline substances may have both low solubility and very low concentration in the solutions.

First hydrolysis constant

The logarithm of hydrolysis constant, K1,-1, for the removal of one proton from an aqua ion

[M(H2O)n]z+ - H+ is in equilibrium with [M(H2O)n-1(OH)](z-1)+
[ [M(OH)]{(z-1)+ ] = K1,-1 [Mz+] [H+] -1

shows a linear relationship with the ratio of charge to M-O distance, z/d. Ions fall into four groups. The slope of the straight line is the same for all groups, but the intercept, A, is different.[47]

log K1,-1 = A + 11.0 z/d
cation A
Mg2+, Ca2+, Sr2+, Ba2+
Al3+, Y3+, La3+
-22.0 ± 0.5
Li+, Na+, K+
Be2+, Mn2+, Fe2+, Co2+, Ni2+, Cu2+, Zn2+, Cd2+
Sc3+, Ti3+, V3+, Cr3+, Fe3+, Rh3+, Ga3+, In3+
Ce4+, Th4+, Pa4+, U4+, Np4+, Pu4+,
-19.8 ± 1
Ag+, Tl+
Ti3+, Bi3+,
-15.9 ± 1
Sn2+, Hg2+, Sn2+, Pd2+ ca. 12

The cations most resistant to hydrolysis for their size and charge are hard pre-transition metal ions or lanthanide ions. The slightly less resistant group includes the transition metal ions. The third group contains mostly soft ions ion of post-transition metals. The ions which show the strongest tendency to hydrolyze for their charge and size are Pd2+, Sn2+ and Hg2+.[47]

The standard enthalpy change for the first hydrolysis step is generally not very different from that of the dissociation of pure water. Consequently, the standard enthalpy change for the substitution reaction

[M(H2O)n]z+ +OH- is in equilibrium with :[M(H2O)n-1(OH)](z-1)+ + H2O

is close to zero. This is typical of reactions between a hard cation and a hard anion, such as the hydroxide ion.[48] It means that the standard entropy charge is the major contributor to the standard free energy change and hence the equilibrium constant.

ΔGStrikeO.png = -RT ln K = ΔHStrikeO.png - TΔSStrikeO.png

The change in ionic charge is responsible for the effect as the aqua ion has a greater ordering effect on the solution than the less highly charged hydroxo complex.

Multiple hydrolysis reactions

Beryllium hydrolysis as a function of pH. Water molecules attached to Be are omitted
trimeric hydrolysis product of beryllium
The molybdenum(IV) oxo-aqua ion [Mo3O4(H2O)9]4+

The hydrolysis of beryllium shows many of the characteristics typical of multiple hydrolysis reactions. The concentrations of various species, including polynuclear species with bridging hydroxide ions, change as a function of pH up to the precipitation of an insoluble hydroxide. Beryllium hydrolysis is unusual in that the concentration of [Be(H2O)3(OH)]+ is too low to be measured. Instead a trimer ([Be3(H2O)6(OH3))3+ is formed, whose structure has been confirmed in solid salts. The formation of polynuclear species is driven by the reduction in charge density within the molecule as a whole. The local environment of the beryllium ions approximates to [Be(H2O)2(OH)2]+. The reduction in effective charge releases free energy in the form of a decrease of the entropy of ordering at the charge centers.[49]

Some polynuclear hydrolysis products[50]
Species formula cations structure
M2(OH)+ Be2+, Mn2+, Co2+, Ni2+
Zn2+, Cd2+, Hg2+, Pb2+
single hydroxide bridge between two cations
M2(OH)2(2z-2)+ Cu2+, Sn2+
Al3+, Sc3+, Ln3+, Ti3+, Cr3+
VO2+, UO22+, NpO22+, PuO22+
double hydroxide bridge between two cations
M3(OH)33+ Be2+, Hg2+ six-membered ring with alternate Mn+ and OH- groups
M3(OH)4(3z-4)+ Sn2+, Pb2+
Al3+, Cr3+, Fe3+, In3+
Cube with alternate vertices of Mn+ and OH- groups, one vertex missing
M4(OH)44+ Mg2+, Co2+, Ni2+, Cd2+, Pb2+ Cube with alternate vertices of Mn+ and OH- groups
M4(OH)88+ Zr4+, Th4+ Square of Mn+ ions with double hydroxide bridges on each side of the square

The hydrolysis product of aluminium formulated as [Al13O4(OH)24(H2O)12]7+ is very well characterized and may be present in nature in water at pH ca. 5.4.[51]

The overall reaction for the loss of two protons from an aqua ion can be written as

[M(H2O)n]z+ - 2 H+is in equilibrium with [M(H2O)n-2(OH)2](z-2)+

However, the equilibrium constant for the loss of two protons applies equally well to the equilibrium

[M(H2O)n]z+ - 2 H+is in equilibrium with [MO(H2O)n-2](z-2)+ + H2O

because the concentration of water is assumed to be constant. This applies in general: any equilibrium constant is equally valid for a product with an oxide ion as for the product with two hydroxyl ions. The two possibilities can only be distinguished by determining the structure of a salt in the solid state. Oxo bridges tend to occur when the metal oxidation state is high.[52] An example is provided by the molybdenum(IV) complex [Mo3O4(H2O)9]4+ in which there is a triangle of molybdenum atoms joined by σ- bonds with an oxide bridge on each edge of the triangle and a fourth oxide which bridges to all three Mo atoms.[53]


There are very few oxo-aqua ions of metals in the oxidation state +5 or higher. Rather, the species found in aqueous solution are monomeric and polymeric oxyanions. Oxyanions can be viewed as the end products of hydrolysis, in which there are no water molecules attached to the metal, only oxide ions.

Exchange kinetics

A water molecule in the first solvation shell of an aqua ion may exchange places with a water molecule in the bulk solvent. It is usually assumed that the rate-determining step is a dissociation reaction.

[M(H2O)n]z+ → [M(H2O)n-1]z+* + H2O

The * symbol signifies that this is the transition state in a chemical reaction. The rate of this reaction is proportional to the concentration of the aqua ion, [A].

 \mathrm {rate} = - \left( \frac{d[A]}{dt} \right) _T = k[A].

The proportionality constant, k, is called a first-order rate constant at temperature T. The unit of the rate constant for water exchange is usually taken as mol dm−3s−1. The half-life for this reaction is equal to loge2 / k. This measure is useful because it is independent of concentration. It has the dimension of time. The quantity 1/k, equal to the half life divided by 0.6932, is known as the residence time or time constant.[54]

The residence time for water exchange varies from about 10−10 s for Cs+ to about 10+10 s (more than 200 y) for Ir3+. It depends on factors such as the size and charge on the ion and, in the case of transition metal ions, crystal field effects. Very fast and very slow reactions are difficult to study. The most information on the kinetics a water exchange comes from systems with a residence time between about 1 μs and 1 s. The enthalpy and entropy of activation, ΔH and ΔS can be obtained by observing the variation of rate constant with temperature.

Selected kinetic parameters for water exchange at 25°C[55]
ion residence
time /μs
/kJ mol-1
/J deg−1mol−1
Be2+ 1.0×103
Mg2+ 2.0 43 8
V2+ 1.3×104 69 21
Cr2+ 0.0032 13 -13
Mn2+ 0.0316 34 12
Fe2+ 0.32 32 -13
Co2+ 0.79 33 -17
Ni2+ 40 43 -22
Cu2+ 5.0×10−4 23 25
Zn2+ 0.032
Al3+ 6.3×106 11 117
Ti3+ 16 26 -63
Cr3+ 2.0×1012 109 0
Fe3+ 316 37 -54
Ga3+ 501 26 -92
Rh3+ 3.2×1013 134 59
In3+ 50 17
La3+ 0.050
UO22+ 1.3

The kinetic parameters for aqua ions of the transition metals are affected by crystal field stabilization energy (CFSE) on both the aqua ion and its dissociation product which has one less water molecule in the primary solvation shell. This explains the inertness (long residence time) of Cr3+, which has an octahedral structure and a d3 electronic configuration and Rh3+ and Ir3+ which are also octahedral and have a low-spin d6 configuration. Trivalent ions have longer residence time than divalent ions, except for the very large lanthanide ions. The values in the table show that this is due to both activation enthalpy and entropy factors.[56]


The assignment of mechanisms to the water exchange reaction has not yet been made unambiguously. The simple dissociation mechanism needs to be modified to take into account the presence of a second solvation shell. In all cases exchange between molecules in the second shell and bulk water is fast compared to the rate-determining step. Three possible mechanism arise.[57]

  • Id (dissociative interchange) mechanism. A water molecule leaves the first solvation shell and enters the second shell.
  • Ia (associative interchange) mechanism. The same exchange reaction takes place, but there is a significant interaction, in the transition state, between the incoming water molecule and the metal ion.
  • A mechanism. This is an associative mechanism involving a molecule from the bulk solvent.

This classification represents extreme cases; the actual mechanism may have something of an intermediate nature. The most useful parameter for distinguishing between possible mechanism is the activation volume change, ΔV, which is obtained from the pressure variation of the rate constant.[58]

\left( \frac{\partial ln k}{\partial P} \right) _T = \frac {\Delta V^\ddagger}{RT}

The activation volume increases for dissociative mechanisms and decreases for associative mechanisms.

Formation of complexes

Metal aqua ions are often involved in the formation of complexes. The reaction may be written as

pMx+(aq) + qLy- → [MpLq](px-qy)+

In reality this is a substitution reaction in which one or more water molecules from the first hydration shell of the metal ion are replaced by ligands, L. The complex is described as an inner-sphere complex. A complex such as [ML](p-q)+ may be described as a contact ion pair.

When the water molecule(s) of the second hydration shell are replaced by ligands, the complex is said to be an outer-sphere complex, or solvent-shared ion pair. The formation of solvent-shared or contact ion pairs is particularly relevant to the determination of solvation numbers of aqua ions by methods that require the use of concentrated solutions of salts, as ion-pairing is concentration dependent. Consider, for example, the formation of the complex [MgCl]+ in solutions of MgCl2. The formation constant, K, of the complex is about 1 but varies with ionic strength.[59] The concentration of the rather weak complex increases from about 0.1% for a 10mM solution to about 70% for a 1M solution (1M = 1 mol dm−3).


The standard electrode potential for the half-cell equilibrium Mz+ + z e is in equilibrium with M(s) has been measured for all metals except for the heaviest trans-uranium elements.

Standard electrode potentials /V for couples Mz+/M(s)[60]
... Zn2+
... Cd2+
... Hg2+
Standard electrode potentials /V for 1st. row transition metal ions[60]
Couple Ti V Cr Mn Fe Co Ni Cu
M2+ / M −1.63 −1.18 −0.91 −1.18 −0.473 −0.28 −0.228 +0.345
M3+ / M −1.37 −0.87 −0.74 −0.28 −0.06 +0.41
Miscellaneous standard electrode potentials /V[60]
Ag+ / Ag Pd2+ / Pd Zr4+ / Zr Hf4+ / Hf Au3+ / Au Ce4+ / Ce
+0.799 +0.915 −1.53 −1.70 +1.50 −1.32

As the standard electrode potential is more negative the aqua ion is more difficult to reduce. For example, comparing the potentials for zinc (-0.75 V) with those of iron (Fe(II) -0.47 V, Fe(III) -0.06 V) it is seen that iron ions are more easily reduced than zinc ions. This is the basis for using zinc to provide anodic protection for large structures made of iron or to protect small structures by galvanization.


  1. ^ Burgess, Section 1.2.
  2. ^ Burgess, p. 20.
  3. ^ Richens, p. 25.
  4. ^ Burgess, p. 181.
  5. ^ Shannon, R.D. (1976). "Revised effective ionic radii and systematic studies of interatomic distances in halides and chalcogenides". Acta Cryst. A 32 (5): 751–767. Bibcode 1976AcCrA..32..751S. doi:10.1107/S0567739476001551. . Richens, Appendix 2.
  6. ^ Burgess, chapter 11.
  7. ^ Burgess, Chapter 6.
  8. ^ Chipperfield, John (1999). Non-Aqueous Solvents. Oxford: OUP. ISBN 0198502591. 
  9. ^ a b Stumm&Morgan
  10. ^ Burgess, p. 53.
  11. ^ Sykes, A.G. (Editor); Johansson, Georg (1992). Structures of Compexes in Solution Derived from X-ray Diffraction Measurements. Advances in Inorganic Chemistry. 39. San Diego: Academic Press. pp. 161–232. ISBN 0120236397. 
  12. ^ Ohtaki, H.; Radnai, T. (1993). "Structure and dynamics of hydrated ions". Chem. Rev. 93 (3): 1157–1204. doi:10.1021/cr00019a014. 
  13. ^ Magini, M.; Licheri, G.; Paschina, G.; Piccaluga,G. Pinna, G. (1988.). X-ray diffraction of ions in aqueous solutions : hydration and complex formation. Boca Raton, Fla: CRC Press. ISBN 0849369452. 
  14. ^ Sykes, A.G. (Editor); Enderby, J.E.; Nielson, G.W. (1989). The Coordination of Metal Ions. Advances in Inorganic Chemistry. 34. San Diego: Academic Press. pp. 195–218. ISBN 0120236346. 
  15. ^ Nieson, G.W.; Enderby, J.E. (1983). "The Structure of an Aqueous Solution of Nickel Chloride". Proc. Roy. Soc., A 390: 353–371. Bibcode 1983RSPSA.390..353N. doi:10.1098/rspa.1983.0136. 
  16. ^ Enderby, J.E. (1987). "Diffraction Studies of Aqueous Ionic Solutions". In Bellisent-Funel, M-C.; Neilson, G.W.. The Physics and Chemistry of Aqueous Solutions. NATO ASI Series. Reidel. pp. 129–145. ISBN 9027725349.  p. 138.
  17. ^ Burgess, p. 85.
  18. ^ Adams, D.M. (1967). Metal-Ligand and Related Vibrations. London: Edward Arnold.  p.254.
  19. ^ Richens, p. 40.
  20. ^ Richens, p.126.
  21. ^ Richens, p. 127.
  22. ^ Richens, p. 140. Data from Persson, I.; Sandstrom, M.; Yokohama, H.; Chaudhry, M. (1995). "?". Z. Naturforsch. A 50: 21. doi:?. 
  23. ^ Richens, p. 129.
  24. ^ Richens, section 2.3.
  25. ^ Richens, p. 176.
  26. ^ Richens, chapters 4-12
  27. ^ Richens, p. 521
  28. ^ Richens, p. 544.
  29. ^ Richens, p. 555.
  30. ^ Richens, p 551.
  31. ^ Richens, p. 282.
  32. ^ Richens, section 2.4.
  33. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. p. 241. ISBN 0080379419. .
  34. ^ Richens, pp152-153.
  35. ^ Richens, p. 157.
  36. ^ Richens, p. 185.
  37. ^ Richens, p. 198.
  38. ^ Richens, p. 236
  39. ^ Richens, p. 240
  40. ^ Richens, p. 278.
  41. ^ Richens, p.202.
  42. ^ a b c Data from Burgess, p. 182.
  43. ^ Richens, Figure 1.2.
  44. ^ Orgel, Lesie E. (1966). An introduction to transition-metal chemistry. Ligand field theory (2nd. ed.). London: Methuen. p. 76. 
  45. ^ Burgess, p. 187.
  46. ^ Baes&Mesmer, chapter 3.
  47. ^ a b Baes&Mesmer, p407
  48. ^ Baes&Mesmer, p 409.
  49. ^ Baes&Mesmer, section 18.2.
  50. ^ Baes&Mesmer, Table 18.3.
  51. ^ Richens, p. 145.
  52. ^ Baes&Mesmer, p. 420.
  53. ^ Richens, Figure 6.26, p. 295
  54. ^ *Atkins, P.W.; de Paula, J. (2006). Physical Chemistry (8th. ed.). Oxford University Press. ISBN 0198700725.  Chapter 22.
  55. ^ Adapted from Burgess, Tables 11.4 and 11.5
  56. ^ Burgess, p. 326.
  57. ^ Burgess, p. 319.
  58. ^ Richens, p. 74. Burgess, p. 323.
  59. ^ IUPAC SC-Database A comprehensive database of published data on equilibrium constants of metal complexes and ligands
  60. ^ a b c Burgess, Table 8.1


  • Baes, C.F.; Mesmer, R.E. (1986) [1976]. The Hydrolysis of Cations. Malabar, FL: Robert E. Krieger. ISBN 0-89874-892-5. 
  • Burgess, John (1978). Metal Ions in Solution. Chichester: Elles Horwood. ISBN 0-85312-027-7. 
  • Richens, David. T. (1997). The Chemistry of Aqua Ions. Wiley. ISBN 0-471-97058-1. 
  • Stumm, Werner; Morgan, James J. (1995). Aquatic Chemistry - Chemical Equilibria and Rates in Natural Waters (3rd. ed.). Wiley-Blackwell. ISBN 0-471-51185-4. 
  • Schweitzer, George K.; Pesterfield, Lester L. (2010). The Aqueous Chemistry of the Elements. Oxford: OUP. ISBN 019539335X. 

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