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Glacier iceberg: pure water

Glacier ice floating

Pure water and solubility

Water is difficult to purify as it dissolves many materials.

V Pure water

V Hydrophobicity

V Solubility of gasses

V Solubility; organic; inorganic

< Ion-pairs

< Moles, molarity, and molality

Water is often called the 'Universal solvent'     

Pure water

Water may be purified by distillation, deionization, reverse osmosis, and carbon filtering [3067], It is very difficult to obtain really pure water (for example, < 5 ng ˣ g-1 solute; also even 'pure' water contains tiny dust particles, colony forming units (CFU) and nanobubbles. Nanobubbles may be removed by extensive degassing by freeze-thaw cycling under reduced pressure or helium washing [1825]. Fine dust (≈ ag) and nanoparticles are difficult to remove but it is possible using several distillations in wax coated vessels (not naked glass or silica as they release fine particles). There are standard specifications for the purest water ( ASTM D 1193). Ultra-pure water has >18.0 MΩ-cm resistivity, < 50 µg ˣ L-1 total organic carbon (TOC), < 1 µg ˣ L-1 Na+, < 1 µg ˣ L-1 Cl-, < 3 µg ˣ L-1 silica, < 10 CFU ˣ L-1. The purest water is for use in manufacturing semiconductors ( ASTM D 5127).; < 5 µg ˣ L-1 total organic carbon (TOC), < 0.05 µg ˣ L-1 Na+, < 0.1 µg ˣ L-1 Cl-. Typical general laboratory 'pure' water has >10.0 MΩ-cm resistivity, < 50 ng g-1 total organic carbon (TOC), < 50 CFU ˣ ml-1. Ultra-pure water may contain 12C, 14N, 16O, 40Ar, 12CmHn, and 14N mH n from atmospheric gasses plus Na, Mg, Cl, Ca, Cr, Fe, Ge, Br, Sb, Ag, and I as detected by ICP-MS [3068]. Pure water should be protected from atmospheric contamination. Care should be taken on storage as ions may leach from glass or polypropylene. For a review of aqueous solubility prediction, see [744.Characterization of a 'pure' water sample should include electrical conductivity, concentration of silica and common ions, total organic carbon concentration, evaporation residue, pH and optical absorbance at 254 nm.

Note that (hexagonal) ice, in contrast to liquid water, is a very poor solvent and this may be made use of when purifying water (for example, degassing) using successive freeze-thaw cycles.

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An important factor in the solubility of organic molecules is their hydrophobicity. Compounds with lower polarity (i.e., greater hydrophobicity) are less able to disrupt the structure of the water molecules. The best measure of polarity is the logarithm of the partition coefficient (LogP) of the organic molecule between n-octanol and water [3073]; the higher the LogP, the more hydrophobic (non-polar) is the compound.

 LogP = Log10(concentration in octanol/concentration in water)


LogP values of some organic compounds






< 0.3



Ethyl acetate


Carbon tetrachloride




Dibutyl ether


Diethyl ether




Methylene chloride




Butyl acetate


Petroleum ether (60-80)


Di-isopropyl ether


Petroleum ether (80-100)




Dipentyl ether








Petroleum ether (100-120)







LogP values increase by about 0.52 for every methylene group (-CH2-) added in a homologous series. Thus, the LogP of hexanol is that of butanol (0.8) plus 2 ˣ 0.52 (i.e., approximately 1.8).


In order to gather a more efficient use of lipophilicity in the intramolecular hydrogen bonding of potential drugs, a more apolar organic partition system utlizing toluene rather than octanol may be used [3425].

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Solubility of gases

the amount of dissolved gas is proportional to its partial pressure in the gas phase

  William Henry, 1803  c  


Equilibrium solubility of oxygen under pressure

from [ IAPWS ]

Solubility of oxygen under pressure

Dissolved gasses are often mistakenly ignored in aqueous solutions although they may impart properties different to those of pure water. At equilibrium, molecules of the solute in the gas phase enter the liquid phase at the same rate as molecules of the solute in the liquid phase escape to the gas phase. Non-polar gases are poorly soluble in water with their equilibrium solubility proportional to their partial pressure (p, fugacity), see right.


Equiilibrium solubilities are described by the use of the Henry's Law constant (Henry's volatility constant, high Henry's constant = high volatility= low solubility; often a cause of confusion as other definitions occur in the Literature. It has been recommended that the Henry's constant used here, and in much of the historical literature, should be alternatively be called Henry's volatility) where Henry's Law states that

Henrys constant= partial pressure/mole fraction


More precisely Henry's volatility constant is defined as

KH= the limit as x2 tends to zero of fugacity/mole fraction

where f2 and x2 is the fugacity and mole fraction of the solute, respectively (IAPWS). In this case the units used for Henry's constant are Pa. Conversion factors for other (equally valid) Henry's constant definitions have been tabulated, together with a large number of Henry's constants for various gasses [3406]. Whatever constants are used, the solubility of the gas can be calculated given its partial pressure,


Equilibrium solubility of gasses with temperature

from [ IAPWS ]

Solubility of gasses with temperature

Most solid solutes dissolve more in water as the temperature is raised. However, whilst most gaseous solutes also dissolve more in most solvents as the temperature is raised, non-polar gases are much more soluble in water at lower temperatures than would be expected from their solubility behavior at high temperatures (see right and anomaly M7).


It may also be seen from the solubility profiles (right) that the gasses are relatively somewhat soluble even up to 100 °C, in contrast to the common mistaken belief that aqueous solutions are efficiently degassed at high temperatures.


Rather surprisingly, no inflections have been found in the solubility data around the density maximum at ≈ 4 °C [3403].


Solubilities for the noble gasses, from [IAPWS, 1166]


Solubilities for the noble gasses in liquid water, http://www.iapws.org/relguide/HenGuide.pdf, Radon data from refs. 1166, 678


The solubilities of the noble gases are shown opposite [IAPWS, 1166]. Their hydration may be considered as the sum of two processes: (A) the endothermic opening of a clathrate pocket in the water, and (B) the exothermic placement of a molecule in that pocket, due to the multiple van der Waals dispersion interactions (for example, krypton dissolved in water is surrounded by a clathrate cage with 20 Kr···OH2 such interactions [1357]). In water at low temperatures, the energy required by the process (A) is very small as such pockets may be easily formed within the water clustering (by CS -> ES).


Using the noble gases to investigate the solvation of non-polar gases is useful as they are spherically symmetrical and have low polarizability, whereas shape and polarizability may confuse the hydration of other gases. The solubility of the noble gases increases considerably as the temperature is lowered. Their enthalpy and entropy of hydration become more negative as their fit into the water dodecahedral clathrate improves.


Equilibrium solubility of gasses from the air, ml, STP

from [ IAPWS ]

Equilibrium solubility of gasses from the air, ml, STP


Under pressure of 101,325 Pa of each gas, the solubilities of the following atmospheric gasses at 0 °C are: nitrogen 1.11 mM, oxygen 2.31 mM, carbon dioxide 77.6 mM, argon 2,51 mM, neon 0.603 mM, helium 0,457 mM, methane 2.61 mM, krypton 5.05 mM, hydrogen 1.07 mM, carbon monoxide 1.71 mM, xenon 10.32 mM. At equilibrium with air at 25 °C and under the atmospheric pressure of 101.325 kPa, the following concentrations of the atmospheric gasses are present in water: nitrogen 0.549 mM, oxygen 0.288 mM, carbon dioxide 14.3 µM, argon 14.1 µM, neon 9.05 nM, helium 2.25 nM, methane 2.71 nM, krypton 3.10 nM, hydrogen 0.454 pM, carbon monoxide 0.158 pM, xenon 4.07 pM (see left, IAPWS ). The solubilities of the inert gasses are given in more detail elsewhere as are those of carbon dioxide and carbon monoxide.


Solubility of methane under high pressure,

from [3407]

Solubility of methane under high pressure, from [3407]




The solubility of gases diverges from Henry's Law above about a MPa. At very high pressures (see that for methane left) there may be a transformation into clathrate hydrates of filled ices [3407]. It is expected that other gasses such as O2 and N2 behave similarly.











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Solubility of solids

Solubility is the capacity of a solute, to dissolve in a solvent. The equilibrium solubility depends in a complex manner on the molecular properties of the solute, any cosolute(s) including their concentration(s), solvent and any cosolute(s) including their concentration(s), as well as the temperature, pressure, pH, ionic strength, and sometimes on other factors. Solubilities are difficult to predict in a quantitative manner.


The excellent solvent properties of water, together with its non-toxic nature, make water a preferred solvent for many chemical reactions [1566]. It has been shown that aqueous solvation takes place in two stages; a rapid partial water rearrangement (< ps) counteracting the polarity of the solute but building up a strain within the water's hydrogen bonding, followed by a slower relaxation (> ns) of this hydrogen bonding involving reorientation of some of the water molecules [3389].


Aqueous solubility is determined by a number of factors: (1) the crystallinity of the solute, (2) the interactions between the solute and water, (3) any ionization, dissociation, and stability issues, (4) the temperature and (5) the ionic strength and cosolutes. Hydrophilic organic compounds containing several oxygen atoms and/or nitrogen atoms are generally soluble in water, particularly if they possess positive and/or negative charges. This is due to their strong interactions with water molecules. At high concentrations, the solute may form soluble aggregates to various extent that can be compared by the use of molecular dynamics and graph theory [3371].


Organic compounds


The solubility of organic molecules may be estimated from the general solubility equation [3078].,


logS = 0.5 - 0.01 ˣ (MPt - 25) - LogP


where S is the molar aqueous solubility, MPt is the melting point in °C, and LogP is derived from the octanol-water partition coefficient of the solute [3077]. If MPt < 25 °C, the term (MPt - 25) is set to zero. This equation can also be used for weak acids (so long as pKa + logS ≤ 0) and bases (so long as pKa - logS ≤ 14). This covers most weak electrolytes [3079].


Often, changes to the solubility of a pharmaceutical are of real benefit; particularly to allow dissolution of the targeted dose. The general solubility equation can be used together with the known changes in LogP and MPt with small structural changes to predict the most useful changes in the structure required [3074].


Inorganic salts and common ion effects


The small molar volume and high permittivity of water contribute towards water's high dissolving power for salts as they reduce the attractive Coulombic forces between oppositely charged ions and allow multiple stabilizing interactions between the dissolved ions and the water molecules.


Inorganic salts are often classified as soluble, sparingly soluble or insoluble, although all are soluble in water to some extent, even if it may be very small. There are no strict limits for the solubility nomenclature but, generally, soluble salts have solubilities above about 0.1 M whereas insoluble salts have solubilities below about 1 mM.


The solubility product (Ksp) is the product of the molar concentrations of the ions in a saturated solution. It depends on the crystal structure of the undissolved salt. Due to the common ion effect, increasing the anion concentration will reduce the cation concentration and vice versa. As an example, the Ksp of iron(III) hydroxide Fe(OH)3 is 2.79 ˣ 10-39 M4 = [Fe3+] ˣ [OH-] ˣ [OH-] ˣ [OH-] and that of iron(II) hydroxide Fe(OH)2 is 1.4 ˣ 10-15 M3 = [Fe2+] ˣ [OH-] ˣ [OH-]. Thus, the solubility depends on the pH. The cations ion-pair to hydroxide, however, with both FeOH2+(aq) 6% and Fe(OH)2+(aq) being formed.
Without considering this ion-pairing, at pH = 7 there would be 2.79 ˣ 10-18 M iron(III) or 0.14 M iron(II) in saturated solutions with the ferrous ions appearing far more soluble than the ferric ions. However, due to ion-pair formation, hydrolysis, complex ion formation, and ion-water complexes, there are only a few cases in which solubility and Ksp are related in such a simple way [3076].


With metal hydroxides having widely different solubilities, they can often be separated from each other by changing the pH, with one cation precipitating at a particular pH whilst the other remains in solution. Many metal hydroxides are amphoteric, with the precipitated solid hydroxides redissolving in excess hydroxide ion,


Al(OH)3 (s, ppt) + OH-(aq, pH >12) → Al(OH)4-(aq)                b


Some rules concerning inorganic salt solubilities are


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a These Ksp values are from Stefan Franzen; all are at 25 °C. Generally, Ksp values may depend on the way they are determined and there are discrepancies between different sources [3076]. These values ignore any possible ion-pair formation; for example, 0.1 M FeCl3 contains only 10% Fe3+ along with a mixed solution of 42% FeCl2+, 40% FeCl2+, 6% FeOH2+, and 2% Fe(OH)2+ [3075]. See also the calcium carbonate equilibria. Ion-pair formation is particularly relevant at high concentrations. [Back]


b s = solid; aq = aqueous solution; ppt = precipitate. [Back]


b W. Henry, Experiments on the quantity of gases absorbed by water, at different temperatures, and under different pressures, Philosophical Transactions of the Royal Society of London, 93 (1803) 29-42,274-276. [Back]



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