Solubility equilibrium
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Solubility equilibrium is any chemical equilibrium between solid and dissolved states of a compound at saturation.
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.
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[edit] Non-ionic compounds
Dissolution of an organic solid can be described as an equilibrium between the substance in its solid and dissolved forms:
An equilibrium expression for this reaction can be written, as for any chemical reaction (products over reactants):
}{ \left \{\mathrm{{C}_{12}{H}_{22}{O}_{11}}\right\}(s)}](../../../math/6/e/f/6ef0f1b005b104940f149c1bf46db5f9.png)
where K is called the equilibrium constant (or solubility constant) and the square brackets mean molar concentration in mol/L (sometimes called molarity with symbol M). Because a notion of concentration for a solid doesn't make sense, curly brackets are used, which mean activity, around the solid. Luckily, the activity of a solid is almost always equal to one. So, a very simple expression suffices:
\,](../../../math/8/9/3/893d0618c8891f00e41d5f5aa3cdeaa8.png)
This statement says that water at equilibrium with solid sugar contains a concentration equal to K. For table sugar (sucrose) at 25 °C, K = 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.
[edit] Ionic compounds
Ionic compounds normally dissociate into their constituent ions when they dissolve in water. For example, for calcium sulfate:
As for the previous example, the equilibrium expression is:
![K_c = \frac{\left[\mbox{Ca} ^{2+}(aq)\right]\left[\mbox{SO}_4^{2-}(aq)\right]}{ \left\{\mbox{CaSO}_4(s)\right\}}](../../../math/e/d/7/ed7f908ecc647323ff9596ef701bf96c.png)
where K is called the equilibrium (or solubility) constant, the square brackets mean molar concentration (M, or mol/L), and curly brackets mean activity. Since the activity of a pure solid is equal to one, this expression reduces to the solubility product expression:
![K_{sp} = \left[\mbox{Ca}^{2+}(aq)\right]\left[\mbox{SO}_4^{2-}(aq)\right].\,](../../../math/d/6/1/d617991e79c7b1524afa91c66216582a.png)
This expression says that an aqueous solution in equilibrium with (saturated with) solid calcium sulfate has concentrations of these two ions such that their product equals Ksp; for calcium sulfate Ksp = 4.93×10−5. If the solution contains only calcium sulfate the concentration of each ion (and the overall solubility of calcium sulfate) is
![\sqrt{ K_{sp}}=\sqrt{4.93\times10^{-5}}=7.02\times10^{-3}=\left[\mbox{Ca}^{2+}\right]=\left[\mbox{SO}_4^{2-}\right].\,](../../../math/d/2/1/d2153ebe653840e73372f073e0a27ab1.png)
[edit] Solubility constants
Solubility constants have been experimentally determined for a large number of compounds and tables are readily available. For ionic compounds the constants are called solubility products. Concentration units are assumed to be molar (moles per liter) unless otherwise stated. Solubility is sometimes listed in mass units such as 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.
The common-ion effect refers to the fact that solubility equilibria shift in response to Le Chatelier's Principle. In the above example, addition of sulfate ions to a saturated solution of calcium sulfate causes CaSO4 to precipitate until the ions in solution again satisfy the solubility expression. (Addition of sulfate ions could be accomplished by adding a very soluble salt, such as Na2SO4.)
Solubility is sensitive to 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. A thermodynamic approach is required to predict how much and in what direction a particular constant changes.
Some values [1] at 25°C:
- Barium carbonate: 2.60×10-9
- Copper(I) chloride: 1.72×10-7
- Lead(II) sulfate: 1.81×10-8
- Magnesium carbonate: 1.15×10-5
- Silver chloride: 1.70×10-10
- Silver bromide: 7.7×10-13
- Calcium hydroxide: 8.0×10-6
[edit] Table
| Table of Solubility Products | ||||
|---|---|---|---|---|
| Compound | Formula | Temperature | Ksp | Data Source (legend below) |
| Aluminum Hydroxide anhydrous | Al(OH)3 | 20°C | 1.9×10-33 | L |
| Aluminum Hydroxide anhydrous | Al(OH)3 | 25°C | 3×10-34 | w1 |
| Aluminum Hydroxide trihydrate | Al(OH)3 | 20°C | 4×10-13 | C |
| Aluminum Hydroxide trihydrate | Al(OH)3 | 25°C | 3.7×10-13 | C |
| Aluminum Phosphate | AlPO4 | 25°C | 9.84×10-21 | w1 |
| Barium Bromate | Ba(BrO3)2 | 25°C | 2.43×10-4 | w1 |
| Barium Carbonate | BaCO3 | 16°C | 7×10-9 | C, L |
| Barium Carbonate | BaCO3 | 25°C | 8.1×10-9 | C, L |
| Barium Chromate | BaCrO4 | 28°C | 2.4×10-10 | C, L |
| Barium Fluoride | BaF2 | 25.8°C | 1.73×10-6 | C, L |
| Barium Iodate dihydrate | Ba(IO3)2 | 25°C | 6.5×10-10 | C, L |
| Barium Oxalate dihydrate | BaC2O4 | 18°C | 1.2×10-7 | C, L |
| Barium Sulfate | BaSO4 | 18°C | 0.87×10-10 | C, L |
| Barium Sulfate | BaSO4 | 25°C | 1.08×10-10 | C, L |
| Barium Sulfate | BaSO4 | 50°C | 1.98×10-10 | C, L |
| Beryllium Hydroxide | Be(OH)2 | 25°C | 6.92×10-22 | w1 |
| Cadmium Carbonate | CdCO3 | 25°C | 1.0×10-12 | w1 |
| Cadmium Hydroxide | Cd(OH)2 | 25°C | 7.2×10-15 | w1 |
| Cadmium Oxalate trihydrate | CdC4O4 | 18°C | 1.53×10-8 | C, L |
| Cadmium Phosphate | Cd3(PO4)2 | 25°C | 2.53×10-33 | w1 |
| Cadmium Sulfide | CdS | 18°C | 3.6×10-29 | C, L |
| Calcium Carbonate calcite | CaCO3 | 15°C | 0.99×10-8 | C, L |
| Calcium Carbonate calcite | CaCO3 | 25°C | 0.87×10-8 | C, L |
| Calcium Carbonate calcite | CaCO3 | 18-25°C | 4.8×10-9 | P |
| Calcium Chromate | CaCrO4 | 18°C | 2.3×10-2 | L |
| Calcium Fluoride | CaF2 | 18°C | 3.4×10-11 | C, L |
| Calcium Fluoride | CaF2 | 25°C | 3.95×10-11 | C, L |
| Calcium Hydroxide | Ca(OH)2 | 18°C-25°C | 8×10-6 | P |
| Calcium Iodate hexahydrate | Ca(IO3)2 | 18°C | 6.44×10-7 | L |
| Calcium Oxalate monohydrate | CaC2O4 | 18°C | 1.78×10-9 | C, L |
| Calcium Oxalate monohydrate | CaC2O4 | 25°C | 2.57×10-9 | C, L |
| Calcium Phosphate tribasic | Ca3(PO4)2 | 25°C | 2.07×10-33 | w1 |
| Calcium Sulfate | CaSO4 | 10°C | 6.1×10-5 | C, L |
| Calcium Sulfate | CaSO4 | 25°C | 4.93×10-5 | w1 |
| Calcium Tartrate dihydrate | CaC4H4O6 | 18°C | 7.7×10-7 | C, L |
| Chromium Hydroxide II | Cr(OH)2 | 25°C | 2×10–16 | w2 |
| Chromium Hydroxide III | Cr(OH)3 | 25°C | 6.3×10–31 | w2 |
| Cobalt Hydroxide II | Co(OH)2 | 25°C | 1.6×10–15 | w2 |
| Cobalt Sulfide (less soluble form) | CoS | 18°C | 3×10-26 | C, L |
| Cobalt Sulfide (more soluble form) | CoS | 18°C-25°C | 10-21 | P |
| Cupric Iodate | Cu(IO3)2 | 25°C | 1.4×10-7 | C, L |
| Cupric Oxalate | CuC2O4 | 25°C | 2.87×10-8 | C, L |
| Cupric Sulfide | CuS | 18°C | 8.5×10-45 | C, L |
| Cuprous Bromide | CuBr | 18°C-20°C | 4.15×10-8 | C |
| Cuprous Chloride | CuCl | 18°C-20°C | 1.02×10-6 | C |
| Cuprous Iodide | CuI | 18°C-20°C | 5.06×10-12 | C |
| Cuprous Sulfide | Cu2S | 16°C-18°C | 2×10-47 | C, L |
| Cuprous Thiocyanate | CuSCN | 18°C | 1.64×10-11 | C, L |
| Ferric Hydroxide | Fe(OH)3 | 18°C | 1.1×10-36 | C, L |
| Ferrous Carbonate | FeCO3 | 18°C-25°C | 2×10-11 | P |
| Ferrous Hydroxide | Fe(OH)2 | 18°C | 1.64×10-14 | C, L |
| Ferrous Hydroxide | Fe(OH)2 | 25°C | 1×10-15; 8.0×10-16 | P; w2 |
| Ferrous Oxalate | FeC2O4 | 25°C | 2.1×10-7 | C, L |
| Ferrous Sulfide | FeS | 18°C | 3.7×10-19 | C, L |
| Lead Bromide | PbBr2 | 25°C | 6.3×10-6; 6.60×10-6 | P; w1 |
| Lead Carbonate | PbCO3 | 18°C | 3.3×10-14 | C, L |
| Lead Chromate | PbCrO4 | 18°C | 1.77×10-14 | C, L |
| Lead Chloride | PbCl2 | 25.2°C | 1.0×10-4 | L |
| Lead Chloride | PbCl2 | 18°C-25°C | 1.7×10-5 | P |
| Lead Fluoride | PbF2 | 18°C | 3.2×10-8 | C, L |
| Lead Fluoride | PbF2 | 26.6°C | 3.7×10-8 | C, L |
| Lead Hydroxide | Pb(OH)2 | 25°C | 1×10-16; 1.43×10-20 | P; w1 |
| Lead Iodate | Pb(IO3)2 | 18°C | 1.2×10-13 | C, L |
| Lead Iodate | Pb(IO3)2 | 25.8°C | 2.6×10-13 | C, L |
| Lead Iodide | PbI2 | 15°C | 7.47×10-9 | C |
| Lead Iodide | PbI2 | 25°C | 1.39×10-8 | C |
| Lead Oxalate | PbC2O4 | 18°C | 2.74×10-11 | C, L |
| Lead Sulfate | PbSO4 | 18°C | 1.06×10-8 | C, L |
| Lead Sulfide | PbS | 18°C | 3.4×10-28 | C, L |
| Lithium Carbonate | Li2CO3 | 25°C | 1.7×10-3 | C, L |
| Lithium Fluoride | LiF | 25°C | 1.84×10-3 | w1 |
| Lithium Phosphate tribasic | Li3PO4 | 25° | 2.37×10-4 | w1 |
| Magnesium Ammonium Phosphate | MgNH4PO4 | 25°C | 2.5×10-13 | C, L |
| Magnesium Carbonate | MgCO3 | 12°C | 2.6×10-5 | C, L |
| Magnesium Fluoride | MgF2 | 18°C | 7.1×10-9 | C, L |
| Magnesium Fluoride | MgF2 | 25°C | 6.4×10-9 | C, L |
| Magnesium Hydroxide | Mg(OH)2 | 18°C | 1.2×10-11 | C, L |
| Magnesium Oxalate | MgC2O4 | 18°C | 8.57×10-5 | C, L |
| Manganese Carbonate | MnCO3 | 18°C-25°C | 9×10-11 | P |
| Manganese Hydroxide | Mn(OH)2 | 18°C | 4×10-14 | C, L |
| Manganese Sulfide (pink) | MnS | 18°C | 1.4×10-15 | C, L |
| Manganese Sulfide (green) | MnS | 25°C | 10-22 | P |
| Mercuric Bromide | HgBr2 | 25°C | 8×10-20 | L |
| Mercuric Chloride | HgCl2 | 25°C | 2.6×10-15 | L |
| Mercuric Iodide | HgI2 | 25°C | 3.2×10-29 | L |
| Mercuric Sulfide | HgS | 18°C | 4×10-53 to 2×10-49 | C, L |
| Mercurous Bromide | HgBr | 25°C | 1.3×10-21 | C, L |
| Mercurous Chloride | HgCl | 25°C | 2×10-18 | C, L |
| Mercurous Iodide | HgI | 25°C | 1.2×10-28 | C, L |
| Mercurous Sulfate | Hg2SO4 | 25°C | 6.5×10-7 | w1 |
| Nickel Hydroxide | Ni(OH)2 | 25°C | 5.48×10-16 | w1 |
| Nickel Sulfide | NiS | 18°C | 1.4×10-24 | C, L |
| Nickel Sulfide (less soluble form) | NiS | 18°C-25°C | 10-27 | P |
| Nickel Sulfide (more soluble form) | NiS | 18°C-25°C | 10-21 | P |
| Potassium Acid Tartrate | KHC4H4O6 | 18°C | 3.8×10-4 | C, L |
| Potassium Perchlorate | KClO4 | 25°C | 1.05×10-2 | w1 |
| Potassium Periodate | KIO4 | 25° | 3.71×10-4 | w1 |
| Silver Acetate | AgC2H3O2 | 16°C | 1.82×10-3 | L |
| Silver Bromate | AgBrO3 | 20°C | 3.97×10-5 | C, L |
| Silver Bromate | AgBrO3 | 25°C | 5.77×10-5 | C, L |
| Silver Bromide | AgBr | 18°C | 4.1×10-13 | C, L |
| Silver Bromide | AgBr | 25°C | 7.7×10-13 | C, L |
| Silver Carbonate | Ag2CO3 | 25°C | 6.15×10-12 | C, L |
| Silver Chloride | AgCl | 4.7°C | 0.21×10-10 | C, L |
| Silver Chloride | AgCl | 9.7°C | 0.37×10-10 | L |
| Silver Chloride | AgCl | 25°C | 1.56×10-10 | C, L |
| Silver Chloride | AgCl | 50°C | 13.2×10-10 | C, L |
| Silver Chloride | AgCl | 100°C | 21.5×10-10 | C, L |
| Silver Chromate | Ag2CrO4 | 14.8°C | 1.2×10-12 | C, L |
| Silver Chromate | Ag2CrO4 | 25°C | 9×10-12 | C, L |
| Silver Cyanide | Ag2(CN)2 | 20°C | 2.2×10-12 | C, L |
| Silver Dichromate | Ag2Cr2O7 | 25°C | 2×10-7 | L |
| Silver Hydroxide | AgOH | 20°C | 1.52×10-8 | C, L |
| Silver Iodate | AgIO3 | 9.4°C | 0.92×10-8 | C, L |
| Silver Iodide | AgI | 13°C | 0.32×10-16 | C, L |
| Silver Iodide | AgI | 25°C | 1.5×10-16 | C, L |
| Silver Nitrite | AgNO2 | 25°C | 5.86×10-4 | L |
| Silver Oxalate | Ag2C2O4 | 25°C | 1.3×10-11 | L |
| Silver Sulfate | Ag2SO4 | 18°C-25°C | 1.2×10-5 | P |
| Silver Sulfide | Ag2S | 18°C | 1.6×10-49 | C, L |
| Silver Thiocyanate | AgSCN | 18°C | 0.49×10-12 | C, L |
| Silver Thiocyanate | AgSCN | 25°C | 1.16×10-12 | C, L |
| Strontium Carbonate | SrCO3 | 25°C | 1.6×10-9 | C, L |
| Strontium Chromate | SrCrO4 | 18°C-25°C | 3.6×10-5 | P |
| Strontium Fluoride | SrF2 | 18°C | 2.8×10-9 | C, L |
| Strontium Oxalate | SrC2O4 | 18°C | 5.61×10-8 | C, L |
| Strontium Sulfate | SrSO4 | 2.9°C | 2.77×10-7 | C, L |
| Strontium Sulfate | SrSO4 | 17.4°C | 2.81×10-7 | C, L |
| Thallous Bromide | TlBr | 25°C | 4×10-6 | L |
| Thallous Chloride | TlCl | 25°C | 2.65×10-4 | L |
| Thallous Sulfate | Tl2SO4 | 25°C | 3.6×10-4 | L |
| Thallous Thiocyanate | TlSCN | 25°C; | 2.25×10-4 | L |
| Tin Hydroxide | Sn(OH)2 | 25°C | 5.45×10-27; 1.4×10–28 | w1; w2 |
| Tin sulfide | SnS | 25°C | 10-28 | P |
| Zinc Hydroxide | Zn(OH)2 | 18°C-20°C | 1.8×10-14 | C, L |
| Zinc Oxalate dihydrate | ZnC2O4 | 18°C | 1.35×10-9 | C, L |
| Zinc Sulfide | ZnS | 18°C | 1.2×10-23 | C, L |
| data source legend: L=Lange's 10th ed.; C=CRC 44th ed.; P=General Chemistry by Pauling, 1970 ed.; w1=Web source 1; w2=Web source 2 | ||||
[edit] References
- ^ H.P.R. Frederikse, David R. Lide: CRC Handbook of Chemistry and Physics. ISBN 0-8493-0478-4.
