Maximum solubility of salts
Salt solubility is the equilibrium amount of a specified solid that can dissolve in water at a stated temperature. At 20 °C, the values below show how widely common salts differ: calcium chloride dissolves hundreds of thousands of times more by mass than barium sulfate.
Basis and limitations: The table gives approximate grams of salt per 100 mL of initially pure water at 20 °C. It is a reference comparison, not a process-design value. Published tables may report grams per 100 g of water or per 100 mL of saturated solution instead; these bases are not identical. Hydrate form, pH, dissolved carbon dioxide, other ions, and measurement convention can also change reported values. Barium chloride is shown as the dihydrate because that form is used in the cited reference.
Approximate solubility at 20 °C
| Salt | Formula | Approximate solubility (g/100 mL water) | General description |
|---|---|---|---|
| Calcium chloride | CaCl2 | 74.5 | Very soluble |
| Ammonium sulfate | (NH4)2SO4 | 75.4 | Very soluble |
| Magnesium chloride hexahydrate | MgCl2·6H2O | 54.6 | Very soluble |
| Ammonium chloride | NH4Cl | 37.2 | Soluble |
| Sodium chloride | NaCl | 35.9 | Soluble |
| Barium chloride dihydrate | BaCl2·2H2O | 35.7 | Soluble |
| Potassium chloride | KCl | 34.2 | Soluble |
| Magnesium sulfate | MgSO4 | 30.0 | Soluble; commonly encountered as a hydrate |
| Sodium carbonate | Na2CO3 | 21.5 | Soluble |
| Sodium sulfate | Na2SO4 | 19.5 | Soluble |
| Potassium sulfate | K2SO4 | 11.1 | Soluble |
| Potassium carbonate | K2CO3 | 112 | Very soluble |
| Ammonium carbonate | (NH4)2CO3 | No single stable value | Reported values are method- and composition-dependent; the material can release ammonia and carbon dioxide. |
| Calcium sulfate (gypsum) | CaSO4 | 0.21 | Sparingly soluble |
| Calcium carbonate (calcite) | CaCO3 | 0.0014 | Very sparingly soluble |
| Barium carbonate | BaCO3 | 0.0014 | Very sparingly soluble |
| Barium sulfate (barite) | BaSO4 | 0.00024 | Very sparingly soluble |
How to interpret salt combinations
Many sodium, potassium, and ammonium salts dissolve readily, but “always soluble” is too broad: the anion, temperature, and solution composition matter. In contrast, calcium and barium carbonates and sulfates have low solubilities in pure water. Their presence does not by itself mean that a solid will immediately form. Precipitation occurs when dissolved-ion activities exceed the relevant mineral's saturation condition; pH, alkalinity, ionic strength, complexing agents, and temperature affect that condition.
These values are useful for comparing pure-water reference solubilities. Natural water and process streams contain mixtures of ions, so their behavior can differ substantially from the single-salt values shown here. For a practical discussion of how solubility and mixed-ion chemistry shape high-recovery reverse osmosis and brine treatment, see Overcoming the RO Wall: Salt Solubility and Advanced Brine Treatment.
Salts commonly concentrated in RO brine
Reverse osmosis removes water while retaining most dissolved ions in the concentrate. Seawater and many brackish waters contain substantial sodium and chloride, as well as magnesium, sulfate, calcium, potassium and bicarbonate alkalinity. The proportions vary by source; bicarbonate is an aqueous alkalinity species, not usually a stable solid bicarbonate mineral in ordinary water.
| Common dissolved component | Representative formula or ions | Why it matters during concentration |
|---|---|---|
| Sodium chloride | NaCl | Usually the dominant dissolved salt in seawater; its concentration rises as water is recovered. |
| Magnesium chloride | MgCl2 | Represents magnesium and chloride in saline waters; magnesium can also influence carbonate and silica chemistry. |
| Magnesium sulfate | MgSO4 | Represents a common combination of major seawater ions and becomes more concentrated in the brine. |
| Potassium chloride and sodium sulfate | KCl; Na2SO4 | Concentrate with the other highly soluble salts; at advanced evaporation, potassium- and magnesium-bearing minerals may crystallise. |
| Calcium and bicarbonate alkalinity | Ca2+; HCO3− | These dissolved species can lead to calcium carbonate supersaturation as concentration, pH and carbon dioxide conditions change. |
Minerals that may precipitate
Concentration does not guarantee that a solid will form. Precipitation becomes possible when the water reaches supersaturation for a mineral; pH, temperature, ion activities, alkalinity and other dissolved constituents affect the result. Common potential deposits in RO systems include:
| Potential solid | Formula | Relevance |
|---|---|---|
| Calcite or aragonite | CaCO3 | Carbonate scale where calcium and carbonate-system chemistry permit supersaturation. |
| Gypsum or anhydrite | CaSO4·2H2O; CaSO4 | Calcium sulfate scale; the solid phase depends on temperature and solution conditions. |
| Barite | BaSO4 | Low-solubility sulfate scale when barium and sulfate are both present. |
| Celestite | SrSO4 | Potential strontium sulfate scale in waters containing strontium and sulfate. |
| Calcium phosphate phases | Several possible formulas | May form in waters with calcium and phosphate; the solid phase depends on pH and water chemistry. |
| Fluorite | CaF2 | Potential scale where calcium and fluoride concentrations are sufficient. |
| Silica deposits | SiO2 and metal silicates | Silica is not a salt, but can limit recovery by forming deposits in some concentrates. |
During strong evaporation of seawater-like brine, carbonate minerals and calcium sulfate can precipitate before halite; still more concentrated residual brine can form potassium- and magnesium-bearing salts. This is an illustrative sequence for seawater-like composition, not a universal order for every RO concentrate or industrial water.
Use the right comparison: The first table compares approximate solubility in pure water at 20 °C. RO concentrates are mixed, high-ionic-strength solutions, so those single-salt values cannot predict their saturation limits directly. Evaluate an actual stream with its full analysis, temperature, pH and recovery conditions.
Discuss your brine treatment challenge
De-risking a brine project: from concept to pilot
- Characterise the stream: Obtain a representative sample and analyse major ions, alkalinity, silica, metals, organics and relevant operating conditions.
- Model the chemistry: Evaluate osmotic pressure and mineral saturation using a model appropriate for the brine’s ionic strength and temperature range.
- Test the process: Run laboratory trials on representative samples to check separation performance, precipitation, fouling and residuals.
- Validate at pilot scale: Where warranted, use pilot operation to assess variability, cleaning needs, control strategy and practical performance before scale-up.
Lenntech’s engineering team can assess separation and concentration options for a specific water or brine composition.
To start a discussion:
email info@lenntech.com.
Sources and further reading
- Royal Society of Chemistry: sulfate and carbonate solubility of Group 1 and 2 salts (20 °C reference values for calcium and barium sulfate and carbonate).
- Merck: solubility table for compounds in water versus temperature (values per 100 g water; hydrate forms identified).
- PubChem: calcium chloride and PubChem: potassium carbonate.
- OECD SIDS: ammonium carbonate (reported solubility and possible decomposition).
- Health Canada: barium in drinking water (physicochemical properties of selected barium compounds).
- USDA: magnesium chloride properties (identifies the hexahydrate and reports solubility at 20 °C).
- WHO/ILO International Chemical Safety Card: magnesium sulfate (water solubility at 20 °C and hydrate forms).
- USGS: major dissolved constituents and brines.
- Sandia National Laboratories: seawater composition and mineral precipitation during evaporation.
- Review: reverse-osmosis fouling and mineral scaling.