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Glacier iceberg: pure water

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Pure water and solubility

Water is difficult to purify as it dissolves many materials.

V Pure water

V Hydrophobicity

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.


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).


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Solubility of solids

The excellent solvent properties of water, together with its non-toxic nature, make water a preferred solvent for many chemical reactions [1566]. 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.


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 attactive 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]



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This page was established in 2017 and last updated by Martin Chaplin on 12 July, 2018

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