1 IUPAC definition 2 Qualifiers used to describe extent of solubility 3 Molecular view 4 Factors affecting solubility
4.1 Temperature 4.2 Pressure
9.1 Differential solubility
IUPAC definition According to the IUPAC definition, solubility is the analytical composition of a saturated solution expressed as a proportion of a designated solute in a designated solvent. Solubility may be stated in various units of concentration such as molarity, molality, mole fraction, mole ratio, mass (solute) per volume (solvent) and other units.
Qualifiers used to describe extent of solubility The extent of solubility ranges widely, from infinitely soluble (without limit) (fully miscible) such as ethanol in water, to poorly soluble, such as silver chloride in water. The term insoluble is often applied to poorly or very poorly soluble compounds. A number of other descriptive terms are also used to qualify the extent of solubility for a given application. For example, U.S. Pharmacopoeia gives the following terms:
Mass parts of solvent required to dissolve 1 mass part of solute
1 to 10
10 to 30
30 to 100
100 to 1000
Very slightly soluble
1000 to 10,000
Practically insoluble or insoluble
The thresholds to describe something as insoluble, or similar terms, may depend on the application. For example, one source states that substances are described as "insoluble" when their solubility is less than 0.1 g per 100 mL of solvent.
Factors affecting solubility
The solubility of a given solute in a given solvent typically depends
on temperature. Depending on the nature of the solute the solubility
may increase or decrease with temperature. For most solids and
liquids, their solubility increases with temperature. In
liquid water at high temperatures, (e.g., that approaching the
critical temperature), the solubility of ionic solutes tends to
decrease due to the change of properties and structure of liquid
water; the lower dielectric constant results in a less polar solvent.
Gaseous solutes exhibit more complex behavior with temperature. As the
temperature is raised, gases usually become less soluble in water (to
minimum, which is below 120 °C for most permanent
gases), but more soluble in organic solvents.
The chart shows solubility curves for some typical solid inorganic
salts (temperature is in degrees
The solubility of organic compounds nearly always increases with temperature. The technique of recrystallization, used for purification of solids, depends on a solute's different solubilities in hot and cold solvent. A few exceptions exist, such as certain cyclodextrins.
Pressure For condensed phases (solids and liquids), the pressure dependence of solubility is typically weak and usually neglected in practice. Assuming an ideal solution, the dependence can be quantified as:
i , a q
i , c r
displaystyle left( frac partial ln N_ i partial P right)_ T =- frac V_ i,aq -V_ i,cr RT
where the index i iterates the components, Ni is the mole fraction of the ith component in the solution, P is the pressure, the index T refers to constant temperature, Vi,aq is the partial molar volume of the ith component in the solution, Vi,cr is the partial molar volume of the ith component in the dissolving solid, and R is the universal gas constant. The pressure dependence of solubility does occasionally have practical significance. For example, precipitation fouling of oil fields and wells by calcium sulfate (which decreases its solubility with decreasing pressure) can result in decreased productivity with time.
displaystyle p=k_ rm H ,c
where kH is a temperature-dependent constant (for example, 769.2 L·atm/mol for dioxygen (O2) in water at 298 K), p is the partial pressure (atm), and c is the concentration of the dissolved gas in the liquid (mol/L). The solubility of gases is sometimes also quantified using Bunsen solubility coefficient. In the presence of small bubbles, the solubility of the gas does not depend on the bubble radius in any other way than through the effect of the radius on pressure (i.e., the solubility of gas in the liquid in contact with small bubbles is increased due to pressure increase by Δp = 2γ/r; see Young–Laplace equation). Henry's law is valid for gases that do not undergo speciation on dissolution. Sieverts' law shows a case when this assumption does not hold. The carbon dioxide solubility in seawater is also affected by temperature and by the carbonate buffer. The decrease of solubility of carbon dioxide in seawater when temperature increases is also an important retroaction factor (positive feedback) exacerbating past and future climate changes as observed in ice cores from the Vostok site in Antarctica. At the geological time scale, because of the Milankovich cycles, when the astronomical parameters of the Earth orbit and its rotation axis progressively change and modify the solar irradiance at the Earth surface, temperature starts to increase. When a deglaciation period is initiated, the progressive warming of the oceans releases CO2 in the atmosphere because of its lower solubility in warmer sea water. On its turn, higher levels of CO2 in the atmosphere increase the greenhouse effect and carbon dioxide acts as an amplifier of the general warming.
Polarity A popular aphorism used for predicting solubility is "like dissolves like" also expressed in the Latin language as "Similia similibus solventur". This statement indicates that a solute will dissolve best in a solvent that has a similar chemical structure to itself. This view is simplistic, but it is a useful rule of thumb. The overall solvation capacity of a solvent depends primarily on its polarity.[nb 1] For example, a very polar (hydrophilic) solute such as urea is very soluble in highly polar water, less soluble in fairly polar methanol, and practically insoluble in non-polar solvents such as benzene. In contrast, a non-polar or lipophilic solute such as naphthalene is insoluble in water, fairly soluble in methanol, and highly soluble in non-polar benzene.
Dissolution of sodium chloride in water. In even more simple terms a simple ionic compound (with positive and negative ions) such as sodium chloride (common salt) is easily soluble in a highly polar solvent (with some separation of positive (δ+) and negative (δ-) charges in the covalent molecule) such as water, as thus the sea is salty as it accumulates dissolved salts since early geological ages. The solubility is favored by entropy of mixing (ΔS) and depends on enthalpy of dissolution (ΔH) and the hydrophobic effect. The free energy of dissolution (Gibbs energy) depends on temperature and is given by the relationship: ΔG = ΔH – TΔS. Chemists often exploit differences in solubilities to separate and purify compounds from reaction mixtures, using the technique of liquid-liquid extraction. This applies in vast areas of chemistry from drug synthesis to spent nuclear fuel reprocessing.
Rate of dissolution Dissolution is not an instantaneous process. The rate of solubilization (in kg/s) is related to the solubility product and the surface area of the material. The speed at which a solid dissolves may depend on its crystallinity or lack thereof in the case of amorphous solids and the surface area (crystallite size) and the presence of polymorphism. Many practical systems illustrate this effect, for example in designing methods for controlled drug delivery. In some cases, solubility equilibria can take a long time to establish (hours, days, months, or many years; depending on the nature of the solute and other factors). The rate of dissolution can be often expressed by the Noyes–Whitney equation or the Nernst and Brunner equation of the form:
displaystyle frac mathrm d m mathrm d t =A frac D d (C_ mathrm s -C_ mathrm b )
m = mass of dissolved material t = time A = surface area of the interface between the dissolving substance and the solvent D = diffusion coefficient d = thickness of the boundary layer of the solvent at the surface of the dissolving substance Cs = mass concentration of the substance on the surface Cb = mass concentration of the substance in the bulk of the solvent For dissolution limited by diffusion (or mass transfer if mixing is present), Cs is equal to the solubility of the substance. When the dissolution rate of a pure substance is normalized to the surface area of the solid (which usually changes with time during the dissolution process), then it is expressed in kg/m2s and referred to as "intrinsic dissolution rate". The intrinsic dissolution rate is defined by the United States Pharmacopeia. Dissolution rates vary by orders of magnitude between different systems. Typically, very low dissolution rates parallel low solubilities, and substances with high solubilities exhibit high dissolution rates, as suggested by the Noyes-Whitney equation.
Quantification of solubility
Differential solubility In flowing systems, differences in solubility often determine the dissolution-precipitation driven transport of species. This happens when different parts of the system experience different conditions. Even slightly different conditions can result in significant effects, given sufficient time. For example, relatively low solubility compounds are found to be soluble in more extreme environments, resulting in geochemical and geological effects of the activity of hydrothermal fluids in the Earth's crust. These are often the source of high quality economic mineral deposits and precious or semi-precious gems. In the same way, compounds with low solubility will dissolve over extended time (geological time), resulting in significant effects such as extensive cave systems or Karstic land surfaces.
AgCl(s) ⇌ Ag+(aq) + Cl−(aq) However, there is a limit to how much salt can be dissolved in a given volume of water. This amount is given by the solubility product, Ksp. This value depends on the type of salt (AgCl vs. NaCl, for example), temperature, and the common ion effect. One can calculate the amount of AgCl that will dissolve in 1 liter of water, some algebra is required.
Ksp = [Ag+] × [Cl−] (definition of solubility product) Ksp = 1.8 × 10−10 (from a table of solubility products) [Ag+] = [Cl−], in the absence of other silver or chloride salts,
[Ag+]2 = 1.8 × 10−10 [Ag+] = 1.34 × 10−5 The result: 1 liter of water can dissolve 1.34 × 10−5 moles of AgCl(s) at room temperature. Compared with other types of salts, AgCl is poorly soluble in water. In contrast, table salt (NaCl) has a higher Ksp and is, therefore, more soluble.
Group I (expect lithium phosphate) and NH4+ compounds Carbonates (Except Group I, NH4+ and uranyl compounds)
Nitrates Sulfites (Except Group I and NH4+ compounds)
Acetates (ethanoates) (Except Ag+ compounds) Phosphates (Except Group I (except for Li+) and NH4+ compounds)
Chlorides (chlorates and perchlorates), bromides and iodides (Except Ag+, Pb2+, Cu+ and Hg22+) Hydroxides and oxides (Except Group I, NH4+, Ba2+, Sr2+ and Tl+)
Sulfates (Except Ag+, Pb2+, Ba2+, Sr2+ and Ca2+) Sulfides (Except Group I, Group II and NH4+ compounds)
Incongruent dissolution Many substances dissolve congruently; i.e., the composition of the solid and the dissolved solute stoichiometrically match. However, some substances may dissolve incongruently, whereby the composition of the solute in solution does not match that of the solid. This solubilization is accompanied by alteration of the "primary solid" and possibly formation of a secondary solid phase. However, in general, some primary solid also remains and a complex solubility equilibrium establishes. For example, dissolution of albite may result in formation of gibbsite.
NaAlSi3O8(s) + H+ + 7H2O ⇌ Na+ + Al(OH)3(s) + 3H4SiO4. In this case, the solubility of albite is expected to depend on the solid-to-solvent ratio. This kind of solubility is of great importance in geology, where it results in formation of metamorphic rocks.
Thermodynamic cycle for calculating solvation via sublimation Thermodynamic cycle for calculating solvation via fusion These cycles have been used for attempts at first principles predictions (solving using the fundamental physical equations) using physically motivated solvent models, to create parametric equations and QSPR models  and combinations of the two. The use of these cycles enables the calculation of the solvation free energy indirectly via either gas (in the sublimation cycle) or a melt (fusion cycle). This is helpful as calculating the free energy of solvation directly is extremely difficult. The free energy of solvation can be converted to a solubility value using various formulae, the most general case being shown below, where the numerator is the free energy of solvation, R is the gas constant and T is the temperature in kelvins.
log S (
− 2.303 R T
displaystyle log S(V_ m )= frac Delta G_ text solvation -2.303RT
Well known fitted equations for solubility prediction are the general solubility equations. These equations stem from the work of Yalkowsky et al. The original formula is given first followed by a revised formula which takes a different assumption of complete miscibility in octanol. These equations are founded on the principles of the fusion cycle.
( S ) = 0.8 −
( P ) − 0.01 (
− 25 )
displaystyle log _ 10 (S)=0.8-log _ 10 (P)-0.01( text melting point -25)
( S ) = 0.5 −
( P ) − 0.01 (
− 25 )
displaystyle log _ 10 (S)=0.5-log _ 10 (P)-0.01( text melting point -25)
Biopharmaceutics Classification System
Flexible SPC water model
Hot water extraction
Rate of solution
^ The solvent polarity is defined as its solvation power according to Reichardt
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