Solubility equilibrium 59529 223932666 2008-07-06T16:07:11Z Broadbot 3410049 robot Adding: [[da:Opløselighedsligevægt]] {{Chemical equilibria}} '''Solubility equilibrium''' is any type [[chemical equilibrium]] between solid and dissolved states of a compound at [[saturation (chemistry)|saturation]]. Solubility equilibria involve application of chemical principles and constants to predict [[solubility]] of substances under specific conditions (because solubility is sensitive to the conditions, while the constants are less so). The substance that is dissolved can be an organic solid such as [[sugar]] or an ionic solid such as [[table salt]]. The main difference is that ionic solids dissociate into constituent ions when they dissolve in water. Most commonly water is the [[solvent]] of interest, although the same basic principles apply with any solvent. In the case of [[environmental science]] studies of water quality, the total concentration of dissolved solids (not necessarily at saturation) is referred to as '''[[total dissolved solids]]'''. == Non-ionic compounds == Dissolution of an '''organic solid''' can be described as an equilibrium between the substance in its solid and dissolved forms: :<math>\mathrm{{C}_{12}{H}_{22}{O}_{11}(s)} \rightleftharpoons \mathrm{{C}_{12}{H}_{22}{O}_{11}(aq)}</math> <!-- reference Atkins, "Physical Chemistry"--> An equilibrium expression for this reaction can be written, as for any chemical reaction (products over reactants): :<math>K = \frac{\left\{\mathrm{{C}_{12}{H}_{22}{O}_{11}}(aq)\right\}}{ \left \{\mathrm{{C}_{12}{H}_{22}{O}_{11}}(s)\right\}}</math> where ''K'' is called the [[equilibrium constant]] (or solubility constant). The curly brackets indicate [[Activity (chemistry)|activity]]. The activity of a pure solid is, by definition, unity. If the activity of the substance in solution is constant (i.e. not affected by any other solutes that may be present) it may be replaced by the concentration. :<math>K_s = \left[\mathrm{{C}_{12}{H}_{22}{O}_{11}}(aq)\right]\,</math> The square brackets mean molar concentration, which is called [[molarity]] with symbol ''M''). This statement says that water '''at equilibrium''' with solid sugar contains a concentration equal to ''K''. For table sugar ([[sucrose]]) at 25 °C, ''K''&nbsp;= 1.971 mol/L. (This solution is very concentrated; sucrose is extremely soluble in water.) This is the maximum amount of sugar that can dissolve at 25 °C; the solution is ''saturated''. If the concentration is below saturation, more sugar dissolves until the solution reaches saturation, or all the solid is consumed. If more sugar is present than is allowed by the solubility expression then the solution is ''supersaturated'' and solid will precipitate until the saturation concentration is reached. This process can be slow; the equilibrium expression describes concentrations when the system reaches equilibrium, not how fast it gets there. ==Ionic compounds== '''Ionic compounds''' normally [[Dissociation (chemistry)|dissociate]] into their constituent ions when they dissolve in water. For example, for [[calcium sulfate]]: :<math>\mathrm{CaSO}_4(s) \rightleftharpoons \mbox{Ca}^{2+}(aq) + \mbox{SO}_4^{2-}(aq)\,</math> As for the previous example, the equilibrium expression is: :<math>K = \frac{\left\{\mbox{Ca} ^{2+}(aq)\right\}\left\{\mbox{SO}_4^{2-}(aq)\right\}}{ \left\{\mbox{CaSO}_4(s)\right\}}</math> <!-- reference Atkins, "Physical Chemistry"--> where ''K'' is called the equilibrium (or solubility) constant and curly brackets indicate activity. The activity of a pure solid is, by definition, equal to one. When the solubility of the salt is very low the activity coefficients of the ions in solution will also be equal to one and this expression reduces to the '''solubility product''' expression: :<math>K_{\mathrm{sp}} = \left[\mbox{Ca}^{2+}(aq)\right]\left[\mbox{SO}_4^{2-}(aq)\right].\,</math> This expression says that an aqueous solution in equilibrium with (''saturated'') solid calcium sulfate has concentrations of these two ions such that their product equals ''K''<sub>sp</sub>; for calcium sulfate ''K''<sub>sp</sub> = 4.93×10<sup>&minus;5</sup>. If the solution contains only calcium sulfate, and the conditions are such that dissolved species are only Ca<sup>2+</sup> and SO<sub>4</sub><sup>2-</sup>, then the concentration of each ion (and the overall [[solubility]] of calcium sulfate) is :<math>\sqrt{ K_{\mathrm{sp}}}=\sqrt{4.93\times10^{-5}}=7.02\times10^{-3}=\left[\mbox{Ca}^{2+}\right]=\left[\mbox{SO}_4^{2-}\right].\,</math> When a solution dissociates into unequal parts as in: :<math>\mathrm{Ca(OH)_2}(s) \rightleftharpoons \mbox{Ca}^{2+}(aq) + \mbox{2OH}^{-}(aq)\,</math>, then determining the solubility from K<sub>sp</sub> is slightly more difficult. Generally, for the dissolution reaction: :<math>\mathrm{A}(s) \rightleftharpoons \mbox{xB}^{p+}(aq) + \mbox{yC}^{q-}(aq)\,</math> the solubility and solubility product are tied with the equation: <math>\sqrt[n]{K_{\mathrm{sp}} \over {x^x \cdot y^y}} = {C \over M_M}</math> :where: :n is the total number of moles on the right hand side, i.e., x+y, dimensionless :x is the number of moles of the cation, dimensionless :y is the number of moles of the anion, dimensionless :K<sub>sp</sub> is the solubility product, (mol/kg)<sup>n</sup> :C is the solubility of A expressed as a mass fraction of the solute A in the solvent (kg of A per kg of solvent) :M<sub>M</sub> is the molecular mass of the compound A, kg/mol. Again, the above equation assumes that the dissolution takes place in pure solvent (no common ion effect), that there is no complexation or hydrolysis (i.e., only ions B<sup>p+</sup> and C<sup>q-</sup> are present in the solution), and that the concentrations are sufficiently low for the activity coefficients to be taken as unity. === Common ion effect === The [[common-ion effect]] refers to the fact that solubility equilibria shift in accordance with [[Le Chatelier's Principle]]. In the above example, addition of [[sulfate]] ions to a saturated solution of calcium sulfate causes CaSO<sub>4</sub> to precipitate until the concentration of the ions in solution are such that they again satisfy the solubility product. (Addition of sulfate ions can, for example, be accomplished by adding a very soluble salt, such as Na<sub>2</sub>SO<sub>4</sub>.) === Salt effect === The [[salt effect]]<ref>{{cite book| title=Vogel's Quantitative Chemical Analysis, 6th edition| editor=J. Mendham, R.C. Denney, J.D. Barnes and M. Thomas| id=ISBN 0-582 22628 7}}</ref> refers to the fact that the presence of another salt, even though there is no common ion, has an effect on the [[ionic strength]] of the solution and hence on the [[activity coefficient]]s of the ions, so that [[solubility]] changes even though K<sub>sp</sub> remains constant (assuming that the activity of the solid remains unity). === Speciation effect === On dissolution, ionic salts typically dissociate into their constituent ions, but the ions may [[Chemical species|speciate]] in the solution. On speciation, the [[solubility]] will always increase although the solubility product does not change. For example, solubility equilibrium for calcium carbonate may be expressed by: :<math>\mathrm{CaCO}_3(s) \rightleftharpoons \mbox{Ca}^{2+}(aq) + \mbox{CO}_3^{2-}(aq)\,</math> :<math>K_{\mathrm{sp}} = \left[\mbox{Ca}^{2+}(aq)\right]\left[\mbox{CO}_3^{2-}(aq)\right].\,</math> Now, if the conditions (e.g., [[pH]]) are such that other carbonate (or calcium) species appear in the solution (for example, [[bicarbonate]] ion HCO<sub>3</sub><sup>-</sup>), then the [[solubility]] of the solid will increase so that the solubility product remains constant. Similarly, if a [[Complex (chemistry)|complexing]] agent, for example [[EDTA]], was present in the solution, [[solubility]] will increase because of the complexation of calcium (a complex has a different chemical identity than uncomplexed Ca<sup>2+</sup> and therefore does not enter the solubility equilibrium). To correctly predict [[solubility]] from a given solubility product, the speciation need to be known (or evaluated, at least approximately). A failure to do so is a common problem and can lead to large errors. === Phase effect === Equilibria are defined for specific crystal [[Phase (matter)|phases]]. Therefore, the solubility product is expected to be different depending on the phase of the solid. For example, [[aragonite]] and [[calcite]] will have different solubility products even though they have both the same chemical identity ([[calcium carbonate]]). Nevertheless, under given conditions, most likely only one phase is thermodynamically stable and therefore this phase enters a true equilibrium. === Particle size effect === The thermodynamic solubility constant is defined for large monocrystals. Solubility will increase with decreasing size of solute particle (or droplet) because of the additional surface energy. This effect is generally small unless particles become very small, typically smaller than 1 μm. The effect of the particle size on solubility constant can be quantified as follows: :<math>\log(^*K_{A}) = \log(^*K_{A \to 0}) + \frac{2 \gamma A_m} {3\ln(10)RT}</math> where <math>^*K_{A}</math> is the solublity constant for the solute particles with the molar surface area A, <math>^*K_{A \to 0}</math> is the solubility constant for substance with molar surface area tending to zero (i.e., when the particles are large), γ is the [[surface tension]] of the solute particle in the solvent, A<sub>m</sub> is the molar surface area of the solute (in m<sup>2</sup>/mol), R is the [[universal gas constant]], and T is the [[absolute temperature]]<ref>Hefter, G.T., Tomkins, R.P.T. (editors), "The Experimental Determination of Solubilities", John Wiley and Sons, Ltd., 2003.</ref>. == Temperature effects == Solubility is sensitive to changes in [[temperature]]. For example, sugar is more soluble in hot water than cool water. It occurs because solubility constants, like other types of equilibrium constant, are functions of temperature. In accordance with [[Le Chatelier's Principle]], when the dissolution process is [[endothermic reaction|endothermic]] (heat is absorbed,) solubility increases with rising temperature, but when the process is [[exothermic]] (heat is released) solubility decreases with rising temperature.<ref name="pauling450">Pauling, Linus: ''General Chemistry'', Dover Publishing, 1970, p 450</ref> <!-- Actually, it is not this simple. Since there is such a favorable entropy change when a solid is dissolved, many solids will become more soluble with increasing temperature, regardless of Le Chatelier's Principle. [[Gibbs Free Energy]] equation expresses the whole picture of this problem. --> == Solubility constants == '''Solubility constants''' have been experimentally determined for a large number of compounds and tables are readily available. For [[Ion|ionic]] compounds the constants are called '''solubility products'''. Concentration units are assumed to be [[molarity|molar]] unless otherwise stated. Solubility is sometimes listed in units of grams dissolved per liter of water. <!--Solubility (and equilibrium) constants themselves are [[dimensionless]] (however, they may have ''units''). The lack of units in the constant may look inconsistent, but it comes about because the use of molar concentration in the solubility expression is only an approximation to ''activity'', a unitless quantity that is approximately equal to molarity at low concentrations. This paragraph is now redundant as the use of concentration products is justified by the treatment of activity coefficients, above --> Some values <ref>{{cite book| title=CRC Handbook of Chemistry and Physics| editor=H.P.R. Frederikse, David R. Lide| id=ISBN 0-8493-0478-4}}</ref> at 25°C: * [[Barium carbonate]]: 2.60{{e|−9}} * [[Copper(I) chloride]]: 1.72{{e|−7}} * [[Lead(II) sulfate]]: 1.81{{e|−8}} * [[Magnesium carbonate]]: 1.15{{e|−5}} * [[Silver chloride]]: 1.70{{e|−10}} * [[Silver bromide]]: 7.7{{e|−13}} * [[Calcium hydroxide]]: 8.0{{e|−6}} See also *[http://srdata.nist.gov/solubility/ IUPAC-NIST solubility database] *[http://www.csudh.edu/oliver/chemdata/data-ksp.htm Solubility products of simple inorganic compounds] ==Table== {| align="left" border="1" cellspacing="0" cellpadding="3" style="margin: 0 0 0 0.5em; background: #FFFFFF; border-collapse: collapse; border-color: #C0C090;" ! {{chembox header}} align="center" colspan="5"| Table of Solubility Products |- | bgcolor="#E0E0E0" align="center"| Compound | bgcolor="#E0E0E0" align="center"| Formula | bgcolor="#E0E0E0" align="center"| Temperature | bgcolor="#E0E0E0" align="center"| ''K''<sub>sp</sub> | bgcolor="#E0E0E0" align="center"|<small> Data Source<br/>(legend below)</small> |- | [[Aluminium hydroxide|Aluminium Hydroxide]] anhydrous | Al(OH)<sub>3</sub> | 20°C | 1.9×10<sup>–33</sub> | L |- | [[Aluminium hydroxide|Aluminium Hydroxide]] anhydrous | Al(OH)<sub>3</sub> | 25°C | 3×10<sup>–34</sub> | w<sub>1</sub> |- | [[Aluminium hydroxide|Aluminium Hydroxide]] trihydrate | Al(OH)<sub>3</sub> | 20°C | 4×10<sup>–13</sub> | C |- | [[Aluminium hydroxide|Aluminium Hydroxide]] trihydrate | Al(OH)<sub>3</sub> | 25°C | 3.7×10<sup>–13</sub> | C |- | [[Aluminium phosphate|Aluminium Phosphate]] | AlPO<sub>4</sub> | 25°C | 9.84×10<sup>–21</sup> | w<sub>1</sub> |- | [[Barium bromate|Barium Bromate]] | Ba(BrO<sub>3</sub>)<sub>2</sub> | 25°C | 2.43×10<sup>–4</sup> | w<sub>1</sub> |- | [[Barium carbonate|Barium Carbonate]] | BaCO<sub>3</sub> | 16°C | 7×10<sup>–9</sup> | C, L |- | [[Barium carbonate|Barium Carbonate]] | BaCO<sub>3</sub> | 25°C | 8.1×10<sup>–9</sup> | C, L |- | [[Barium chromate|Barium Chromate]] | BaCrO<sub>4</sub> | 28°C | 2.4×10<sup>–10</sup> | C, L |- | [[Barium fluoride|Barium Fluoride]] | BaF<sub>2</sub> | 25.8°C | 1.73×10<sup>–6</sup> | C, L |- | [[Barium iodate|Barium Iodate]] dihydrate | Ba(IO<sub>3</sub>)<sub>2</sub> | 25°C | 6.5×10<sup>–10</sup> | C, L |- | [[Barium oxalate|Barium Oxalate]] dihydrate | BaC<sub>2</sub>O<sub>4</sub> | 18°C | 1.2×10<sup>–7</sup> | C, L |- | [[Barium sulfate|Barium Sulfate]] | BaSO<sub>4</sub> | 18°C | 0.87×10<sup>–10</sup> | C, L |- | [[Barium sulfate|Barium Sulfate]] | BaSO<sub>4</sub> | 25°C | 1.08×10<sup>–10</sup> | C, L |- | [[Barium sulfate|Barium Sulfate]] | BaSO<sub>4</sub> | 50°C | 1.98×10<sup>–10</sup> | C, L |- | [[Beryllium hydroxide|Beryllium Hydroxide]] | Be(OH)<sub>2</sub> | 25°C | 6.92×10<sup>–22</sup> | w<sub>1</sub> |- | [[Cadmium carbonate|Cadmium Carbonate]] | CdCO<sub>3</sub> | 25°C | 1.0×10<sup>–12</sup> | w<sub>1</sub> |- | [[Cadmium hydroxide|Cadmium Hydroxide]] | Cd(OH)<sub>2</sub> | 25°C | 7.2×10<sup>–15</sup> | w<sub>1</sub> |- | [[Cadmium oxalate|Cadmium Oxalate]] trihydrate | CdC<sub>2</sub>O<sub>4</sub> | 18°C | 1.53×10<sup>–8</sup> | C, L |- | [[Cadmium phosphate|Cadmium Phosphate]] | Cd<sub>3</sub>(PO<sub>4</sub>)</sub>2</sub> | 25°C | 2.53×10<sup>–33</sup> | w<sub>1</sub> |- | [[Cadmium Sulfide]] | CdS | 18°C | 3.6×10<sup>–29</sup> | C, L |- | [[Calcium Carbonate]] calcite | CaCO<sub>3</sub> | 15°C | 0.99×10<sup>–8</sup> | C, L |- | [[Calcium Carbonate]] calcite | CaCO<sub>3</sub> | 25°C | 0.87×10<sup>–8</sup> | C, L |- | [[Calcium Carbonate]] calcite | CaCO<sub>3</sub> | 18-25°C | 4.8×10<sup>–9</sup> | P |- | [[Calcium chromate|Calcium Chromate]] | CaCrO<sub>4</sub> | 18°C | 2.3×10<sup>–2</sup> | L |- | [[Calcium fluoride|Calcium Fluoride]] | CaF<sub>2</sub> | 18°C | 3.4×10<sup>–11</sup> | C, L |- | [[Calcium fluoride|Calcium Fluoride]] | CaF<sub>2</sub> | 25°C | 3.95×10<sup>–11</sup> | C, L |- | [[Calcium hydroxide|Calcium Hydroxide]] | Ca(OH)<sub>2</sub> | 18°C-25°C | 8×10<sup>–6</sup> | P |- | [[Calcium hydroxide|Calcium Hydroxide]] | Ca(OH)<sub>2</sub> | 25°C | 5.02×10<sup>–6</sup> | w<sub>1</sub> |- | [[Calcium iodate|Calcium Iodate]] hexahydrate | Ca(IO<sub>3</sub>)<sub>2</sub> | 18°C | 6.44×10<sup>–7</sup> | L |- | [[Calcium oxalate|Calcium Oxalate]] monohydrate | CaC<sub>2</sub>O</sub>4</sub> | 18°C | 1.78×10<sup>–9</sup> | C, L |- | [[Calcium oxalate|Calcium Oxalate]] monohydrate | CaC<sub>2</sub>O</sub>4</sub> | 25°C | 2.57×10<sup>–9</sup> | C, L |- | [[Calcium phosphate|Calcium Phosphate]] tribasic | Ca<sub>3</sub>(PO<sub>4</sub>)<sub>2</sub> | 25°C | 2.07×10<sup>–33</sup> | w<sub>1</sub> |- | [[Calcium sulfate|Calcium Sulfate]] | CaSO<sub>4</sub> | 10°C | 6.1×10<sup>–5</sup> | C, L |- | [[Calcium sulfate|Calcium Sulfate]] | CaSO<sub>4</sub> | 25°C | 4.93×10<sup>–5</sup> | w<sub>1</sub> |- | [[Calcium tartrate|Calcium Tartrate]] dihydrate | CaC<sub>4</sub>H<sub>4</sub>O<sub>6</sub> | 18°C | 7.7×10<sup>–7</sup> | C, L |- | [[Chromium hydroxide|Chromium Hydroxide]] II | Cr(OH)<sub>2</sub> | 25°C | 2×10<sup>–16</sup> | w<sub>2</sub> |- | [[Chromium hydroxide|Chromium Hydroxide]] III | Cr(OH)<sub>3</sub> | 25°C | 6.3×10<sup>–31</sup> | w<sub>2</sub> |- | [[Cobalt hydroxide|Cobalt Hydroxide]] II | Co(OH)<sub>2</sub> | 25°C | 1.6×10<sup>–15</sup> | w<sub>2</sub> |- | [[Cobalt sulfide|Cobalt Sulfide]] (less soluble form) | CoS | 18°C | 3×10<sup>–26</sup> | C, L |- | [[Cobalt sulfide|Cobalt Sulfide]] (more soluble form) | CoS | 18°C-25°C | 10<sup>–21</sup> | P |- | [[Copper carbonate|Cupric Carbonate]] | CuCO<sub>3</sub> | 25°C | 1×10<sup>–10</sup> | P |- | [[Copper(II) hydroxide|Cupric Hydroxide]] | Cu(OH)<sub>2</sub> | 18°C-25°C | 6×10<sup>–20</sup> | P |- | [[Copper(II) hydroxide|Cupric Hydroxide]] | Cu(OH)<sub>2</sub> | 25°C | 4.8×10<sup>–20</sup> | w<sub>1</sub> |- | [[Cupric iodate|Cupric Iodate]] | Cu(IO<sub>3</sub>)<sub>2</sub> | 25°C | 1.4×10<sup>–7</sup> | C, L |- | [[Cupric oxalate|Cupric Oxalate]] | CuC<sub>2</sub>O<sub>4</sub> | 25°C | 2.87×10<sup>–8</sup> | C, L |- | [[Copper(II) sulfide|Cupric Sulfide]] | CuS | 18°C | 8.5×10<sup>–45</sup> | C, L |- | [[Copper(I) bromide|Cuprous Bromide]] | CuBr | 18°C-20°C | 4.15×10<sup>–8</sup> | C |- | [[Copper(I) chloride|Cuprous Chloride]] | CuCl | 18°C-20°C | 1.02×10<sup>–6</sup> | C |- | [[Copper(I) oxide|Cuprous Hydroxide]]<br>(in equilib. with Cu<sub>2</sub>O + H<sub>2</sub>O) | Cu(OH) | 25°C | 2×10<sup>–15</sup> | w<sub>1</sub> |- | [[Copper(I) iodide|Cuprous Iodide]] | CuI | 18°C-20°C | 5.06×10<sup>–12</sub> | C |- | [[Copper(I) sulfide|Cuprous Sulfide]] | Cu<sub>2</sub>S | 16°C-18°C | 2×10<sup>–47</sup> | C, L |- | [[Copper(I) thiocyanate|Cuprous Thiocyanate]] | CuSCN | 18°C | 1.64×10<sup>–11</sup> | C, L |- | [[Iron(III) hydroxide|Ferric Hydroxide]] | Fe(OH)<sub>3</sub> | 18°C | 1.1×10<sup>–36</sup> | C, L |- | [[Iron(II) carbonate|Ferrous Carbonate]] | FeCO<sub>3</sub> | 18°C-25°C | 2×10<sup>–11</sup> | P |- | [[Iron(II) hydroxide|Ferrous Hydroxide]] | Fe(OH)<sub>2</sub> | 18°C | 1.64×10<sup>–14</sup> | C, L |- | [[Iron(II) hydroxide|Ferrous Hydroxide]] | Fe(OH)<sub>2</sub> | 25°C | 1×10<sup>–15</sup>; 8.0×10<sup>–16</sup> | P; w<sub>2</sub> |- | [[Ferrous oxalate|Ferrous Oxalate]] | FeC<sub>2</sub>O<sub>4</sub> | 25°C | 2.1×10<sup>–7</sup> | C, L |- | [[Ferrous sulfide|Ferrous Sulfide]] | FeS | 18°C | 3.7×10<sup>–19</sup> | C, L |- | [[Lead bromide|Lead Bromide]] | PbBr<sub>2</sub> | 25°C | 6.3×10<sup>–6</sup>; 6.60×10<sup>–6</sup> | P; w<sub>1</sub> |- | [[Lead carbonate|Lead Carbonate]] | PbCO<sub>3</sub> | 18°C | 3.3×10<sup>–14</sub> | C, L |- | [[Lead chromate|Lead Chromate]] | PbCrO<sub>4</sub> | 18°C | 1.77×10<sup>–14</sup> | C, L |- | [[Lead chloride|Lead Chloride]] | PbCl<sub>2</sub> | 25.2°C | 1.0×10<sup>–4</sup> | L |- | [[Lead chloride|Lead Chloride]] | PbCl<sub>2</sub> | 18°C-25°C | 1.7×10<sup>–5</sup> | P |- | [[Lead fluoride|Lead Fluoride]] | PbF<sub>2</sub> | 18°C | 3.2×10<sup>–8</sup> | C, L |- | [[Lead fluoride|Lead Fluoride]] | PbF<sub>2</sub> | 26.6°C | 3.7×10<sup>–8</sup> | C, L |- | [[Plumbous hydroxide|Lead Hydroxide]] | Pb(OH)<sub>2</sub> | 25°C | 1×10<sup>–16</sup>; 1.43×10<sup>–20</sup> | P; w<sub>1</sub> |- | [[Lead iodate|Lead Iodate]] | Pb(IO<sub>3</sub>)<sub>2</sub> | 18°C | 1.2×10<sup>–13</sup> | C, L |- | [[Lead iodate|Lead Iodate]] | Pb(IO<sub>3</sub>)<sub>2</sub> | 25.8°C | 2.6×10<sup>–13</sup> | C, L |- | [[Lead iodide|Lead Iodide]] | PbI<sub>2</sub> | 15°C | 7.47×10<sup>–9</sup> | C |- | [[Lead iodide|Lead Iodide]] | PbI<sub>2</sub> | 25°C | 1.39×10<sup>–8</sup> | C |- | [[Lead oxalate|Lead Oxalate]] | PbC<sub>2</sub>O<sub>4</sub> | 18°C | 2.74×10<sup>–11</sup> | C, L |- | [[Lead sulfate|Lead Sulfate]] | PbSO<sub>4</sub> | 18°C | 1.06×10<sup>–8</sup> | C, L |- | [[Lead sulfide|Lead Sulfide]] | PbS | 18°C | 3.4×10<sup>–28</sup> | C, L |- | [[Lithium carbonate|Lithium Carbonate]] | Li<sub>2</sub>CO<sub>3</sub> | 25°C | 1.7×10<sup>–3</sup> | C, L |- | [[Lithium fluoride|Lithium Fluoride]] | LiF | 25°C | 1.84×10<sup>–3</sup> | w<sub>1</sub> |- | [[Lithium phosphate|Lithium Phosphate]] tribasic | Li<sub>3</sub>PO<sub>4</sub> | 25° | 2.37×10<sup>–4</sup> | w<sub>1</sub> |- | [[Magnesium ammonium phosphate|Magnesium Ammonium Phosphate]] | MgNH<sub>4</sub>PO<sub>4</sub> | 25°C | 2.5×10<sup>–13</sup> | C, L |- | [[Magnesium carbonate|Magnesium Carbonate]] | MgCO<sub>3</sub> | 12°C | 2.6×10<sup>–5</sup> | C, L |- | [[Magnesium fluoride|Magnesium Fluoride]] | MgF<sub>2</sub> | 18°C | 7.1×10<sup>–9</sup> | C, L |- | [[Magnesium fluoride|Magnesium Fluoride]] | MgF<sub>2</sub> | 25°C | 6.4×10<sup>–9</sup> | C, L |- | [[Magnesium hydroxide|Magnesium Hydroxide]] | Mg(OH)<sub>2</sub> | 18°C | 1.2×10<sup>–11</sup> | C, L |- | [[Magnesium oxalate|Magnesium Oxalate]] | MgC<sub>2</sub>O</sub>4</sub> | 18°C | 8.57×10<sup>–5</sup> | C, L |- | [[Manganese carbonate|Manganese Carbonate]] | MnCO<sub>3</sub> | 18°C-25°C | 9×10<sup>–11</sup> | P |- | [[Manganese hydroxise|Manganese Hydroxide]] | Mn(OH)<sub>2</sub> | 18°C | 4×10<sup>–14</sup> | C, L |- | [[Manganese sulfide|Manganese Sulfide]] (pink) | MnS | 18°C | 1.4×10<sup>–15</sup> | C, L |- | [[Manganese sulfide|Manganese Sulfide]] (green) | MnS | 25°C | 10<sup>–22</sup> | P |- | [[Mercuric bromide|Mercuric Bromide]] | HgBr<sub>2</sub> | 25°C | 8×10<sup>–20</sup> | L |- | [[Mercuric chloride|Mercuric Chloride]] | HgCl<sub>2</sub> | 25°C | 2.6×10<sup>–15</sup> | L |- | [[Mercuric oxide|Mercuric Hydroxide]]<br>(equilib. with HgO + H<sub>2</sub>O) | Hg(OH)<sub>2</sub> | 25°C | 3.6×10<sup>–26</sup> | w<sub>1</sub> |- | [[Mercuric iodide|Mercuric Iodide]] | HgI<sub>2</sub> | 25°C | 3.2×10<sup>–29</sup> | L |- | [[Mercuric sulfide|Mercuric Sulfide]] | HgS | 18°C | 4×10<sup>–53</sup> to 2×10<sup>–49</sup> | C, L |- | [[Mercurous bromide|Mercurous Bromide]] | HgBr | 25°C | 1.3×10<sup>–21</sup> | C, L |- | [[Mercurous chloride|Mercurous Chloride]] | Hg<sub>2</sub>Cl<sub>2</sub> | 25°C | 2×10<sup>–18</sup> | C, L |- | [[Mercurous iodide|Mercurous Iodide]] | HgI | 25°C | 1.2×10<sup>–28</sup> | C, L |- | [[Mercurous sulfate|Mercurous Sulfate]] | Hg<sub>2</sub>SO<sub>4</sub> | 25°C | 6×10<sup>–7</sup>; 6.5×10<sup>–7</sup> | P; w<sub>1</sub> |- | [[Nickel hydroxide|Nickel Hydroxide]] | Ni(OH)<sub>2</sub> | 25°C | 5.48×10<sup>–16</sup> | w<sub>1</sub> |- | [[Nickel sulfide|Nickel Sulfide]] | NiS | 18°C | 1.4×10<sup>–24</sup> | C, L |- | [[Nickel sulfide|Nickel Sulfide]] (less soluble form) | NiS | 18°C-25°C | 10<sup>–27</sup> | P |- | [[Nickel sulfide|Nickel Sulfide]] (more soluble form) | NiS | 18°C-25°C | 10<sup>–21</sup> | P |- | [[Potassium bitartrate|Potassium Acid Tartrate]] | KHC<sub>4</sub>H<sub>4</sub>O<sub>6</sub> | 18°C | 3.8×10<sup>–4</sup> | C, L |- | [[Potassium perchlorate|Potassium Perchlorate]] | KClO<sub>4</sub> | 25°C | 1.05×10<sup>–2</sup> | w<sub>1</sub> |- | [[Potassium periodate|Potassium Periodate]] | KIO<sub>4</sub> | 25° | 3.71×10<sup>–4</sup> | w<sub>1</sub> |- | [[Silver acetate|Silver Acetate]] | AgC<sub>2</sub>H<sub>3</sub>O<sub>2</sub> | 16°C | 1.82×10<sup>–3</sup> | L |- | [[Silver bromate|Silver Bromate]] | AgBrO<sub>3</sub> | 20°C | 3.97×10<sup>–5</sup> | C, L |- | [[Silver bromate|Silver Bromate]] | AgBrO<sub>3</sub> | 25°C | 5.77×10<sup>–5</sup> | C, L |- | [[Silver bromide|Silver Bromide]] | AgBr | 18°C | 4.1×10<sup>–13</sup> | C, L |- | [[Silver bromide|Silver Bromide]] | AgBr | 25°C | 7.7×10<sup>–13</sup> | C, L |- | [[Silver carbonate|Silver Carbonate]] | Ag<sub>2</sub>CO<sub>3</sub> | 25°C | 6.15×10<sup>–12</sup> | C, L |- | [[Silver chloride|Silver Chloride]] | AgCl | 4.7°C | 0.21×10<sup>–10</sup> | C, L |- | [[Silver chloride|Silver Chloride]] | AgCl | 9.7°C | 0.37×10<sup>–10</sup> | L |- | [[Silver chloride|Silver Chloride]] | AgCl | 25°C | 1.56×10<sup>–10</sup> | C, L |- | [[Silver chloride|Silver Chloride]] | AgCl | 50°C | 13.2×10<sup>–10</sup> | C, L |- | [[Silver chloride|Silver Chloride]] | AgCl | 100°C | 21.5×10<sup>–10</sup> | C, L |- | [[Silver chromate|Silver Chromate]] | Ag<sub>2</sub>CrO<sub>4</sub> | 14.8°C | 1.2×10<sup>–12</sup> | C, L |- | [[Silver chromate|Silver Chromate]] | Ag<sub>2</sub>CrO<sub>4</sub> | 25°C | 9×10<sup>–12</sup> | C, L |- | [[Silver cyanide|Silver Cyanide]] | Ag<sub>2</sub>(CN)<sub>2</sub> | 20°C | 2.2×10<sup>–12</sup> | C, L |- | [[Silver dichromate|Silver Dichromate]] | Ag<sub>2</sub>Cr<sub>2</sub>O<sub>7</sub> | 25°C | 2×10<sup>–7</sup> | L |- | [[Silver oxide|Silver Hydroxide]] | AgOH | 20°C | 1.52×10<sup>–8</sup> | C, L |- | [[Silver iodate|Silver Iodate]] | AgIO<sub>3</sub> | 9.4°C | 0.92×10<sup>–8</sup> | C, L |- | [[Silver iodide|Silver Iodide]] | AgI | 13°C | 0.32×10<sup>–16</sup> | C, L |- | [[Silver iodide|Silver Iodide]] | AgI | 25°C | 1.5×10<sup>–16</sup> | C, L |- | [[Silver nitrite|Silver Nitrite]] | AgNO<sub>2</sub> | 25°C | 5.86×10<sup>–4</sup> | L |- | [[Silver oxalate|Silver Oxalate]] | Ag<sub>2</sub>C<sub>2</sub>O<sub>4</sub> | 25°C | 1.3×10<sup>–11</sup> | L |- | [[Silver sulfate|Silver Sulfate]] | Ag<sub>2</sub>SO<sub>4</sub> | 18°C-25°C | 1.2×10<sup>–5</sub> | P |- | [[Silver sulfide|Silver Sulfide]] | Ag<sub>2</sub>S | 18°C | 1.6×10<sup>–49</sup> | C, L |- | [[Silver thiocyanate|Silver Thiocyanate]] | AgSCN | 18°C | 0.49×10<sup>–12</sup> | C, L |- | [[Silver thiocyanate|Silver Thiocyanate]] | AgSCN | 25°C | 1.16×10<sup>–12</sup> | C, L |- | [[Strontium carbonate|Strontium Carbonate]] | SrCO<sub>3</sub> | 25°C | 1.6×10<sup>–9</sup> | C, L |- | [[Strontium chromate|Strontium Chromate]] | SrCrO<sub>4</sub> | 18°C-25°C | 3.6×10<sup>–5</sup> | P |- | [[Strontium fluoride|Strontium Fluoride]] | SrF<sub>2</sub> | 18°C | 2.8×10<sup>–9</sup> | C, L |- | [[Strontium oxalate|Strontium Oxalate]] | SrC<sub>2</sub>O<sub>4</sub> | 18°C | 5.61×10<sup>–8</sup> | C, L |- | [[Strontium sulfate|Strontium Sulfate]] | SrSO<sub>4</sub> | 2.9°C | 2.77×10<sup>–7</sub> | C, L |- | [[Strontium sulfate|Strontium Sulfate]] | SrSO<sub>4</sub> | 17.4°C | 2.81×10<sup>–7</sub> | C, L |- | [[Thallous bromide|Thallous Bromide]] | TlBr | 25°C | 4×10<sup>–6</sup> | L |- | [[Thallous chloride|Thallous Chloride]] | TlCl | 25°C | 2.65×10<sup>–4</sup> | L |- | [[Thallous sulfate|Thallous Sulfate]] | Tl<sub>2</sub>SO<sub>4</sub> | 25°C | 3.6×10<sup>–4</sup> | L |- | [[Thallous thiocyanate|Thallous Thiocyanate]] | TlSCN | 25°C; | 2.25×10<sup>–4</sup> | L |- | [[Stannous hydroxide|Tin Hydroxide]] | Sn(OH)<sub>2</sub> | 18°C-25°C | 1×10<sup>–26</sup> | P |- | [[Stannous hydroxide|Tin Hydroxide]] | Sn(OH)<sub>2</sub> | 25°C | 5.45×10<sup>–27</sup>; 1.4×10<sup>–28</sup> | w<sub>1</sub>; w<sub>2</sub> |- | [[Stannous sulfide|Tin sulfide]] | SnS | 25°C | 10<sup>–28</sup> | P |- | [[Zinc hydroxide|Zinc Hydroxide]] | Zn(OH)<sub>2</sub> | 18°C-20°C | 1.8×10<sup>–14</sup> | C, L |- | [[Zinc oxalate|Zinc Oxalate]] dihydrate | ZnC<sub>2</sub>O<sub>4</sub> | 18°C | 1.35×10<sup>–9</sub> | C, L |- | [[Zinc sulfide|Zinc Sulfide]] | ZnS | 18°C | 1.2×10<sup>–23</sup> | C, L |- ! {{chembox header}} align="center" colspan="5"|<small>'''data source legend:''' L=Lange's 10th ed.; C=CRC 44th ed.; P=General Chemistry by Pauling, 1970 ed.; w<sub>1</sub>=[http://www.ktf-split.hr/periodni/en/abc/kpt.html Web source 1]; w<sub>2</sub>=[http://bilbo.chm.uri.edu/CHM112/tables/KspTable.htm Web source 2]</small> |}<br clear="left"/> ==References== <references /> {{Chemical solutions}} [[Category:Solutions]] [[bs:Proizvod rastvorljivosti]] [[da:Opløselighedsligevægt]] [[de:Löslichkeitsgleichgewicht]] [[fr:Produit de solubilité]] [[it:Prodotto di solubilità]] [[he:מכפלת מסיסות]] [[nl:Oplosbaarheidsproduct]] [[nn:Løysingsevneproduktet]] [[pl:Iloczyn rozpuszczalności]]