Precipitations
Reaction types
What is precipitations?
A precipitation is a solid phase coming out of a liquid. This can be done more or less elegantly and controlled.
For educational purposes, the classical reaction between silver and chloride ions is what the students see, and it is also shown here, but after a few calculation on solubility and solubility products, the relevance of the subject usually stops there. Solubility and solubility products are definitely important and of great industrial relevance, it is just not the only thing for which you use precipitations.
For educational purposes, the classical reaction between silver and chloride ions is what the students see, and it is also shown here, but after a few calculation on solubility and solubility products, the relevance of the subject usually stops there. Solubility and solubility products are definitely important and of great industrial relevance, it is just not the only thing for which you use precipitations.
Mixing of ions
The classical way of looking at precipitations is two aqueous solutions of salts. When the two solutions are mixed, a solid is formed. The example favored by the educational system for chemistry classes is NaCl and AgNO3, that are both soluble in water:
NaCl(s)
Na+(aq) + Cl−(aq)
AgNO3(s) Ag+(aq) + NO3−(aq)
If the two solutions are mixed:
Na+(aq) + Cl−(aq) + Ag+(aq) + NO3−(aq) AgCl(s) + Na+(aq) + NO3−(aq)
AgCl is insoluble and phases out as a white precipitate. As seen the sodium and nitrate ions don not participate in the reaction, and thus do not need to be included in the reaction equation. It is sufficient to write
Ag+(aq) + Cl−(aq) AgCl(s)
This is a model, and it isn't entirely correct. There is still a small amount of both Ag+ and Cl− in the solution, and also a small amount of AgCl(aq) will be present in the solution. In reality this is also an equilibrium, not an irreversible process, that looks like this:
Ag+(aq) + Cl−(aq)
AgCl(aq) AgCl(s)
The reason for simplifying the reaction, as it is done, is that the amount of dissolved AgCl, Ag+ and Cl− is negligible and we focus on the precipitation. As soon as you start doing calculations on solubility products and multiple reactions influencing each other, you do use the reaction equation with the equilibria.
Precipitations comes from some units, in this case ions, having a high affinity for each other, and upon contact they are bound to each other, forming a compound that is insoluble in that particular solvent. In the classic chemistry for educational purposes, precipitations are generally associated with ions in aqueous solution, but it is important to keep in mind, that precipitations is far more than this.
The table below shows some combinations of ions that are soluble and insoluble. BUT, precipitations is not an either/or phenomenon, so, at low concentrations the precipitate may not appear, and at sufficiently high concentrations, precipitates may appear, even for salts that are considered soluble.
An important aspect about precipitations, that is easily overlooked, is that precipitations may only occur in a limited concentration interval, e.g. precipitation of Zn2+ with OH−:
Start:
Zn2+(aq) + 2 OH−(aq) Zn(OH)2(s)
At excess OH−:
Zn(OH)2(s) + 2 OH−(aq) [Zn(OH)4]2−(aq)
Therefore, you cannot just blindly mix two sets of ions, and from that conclude whether you have a precipitation or something else (in this case a complex formation).
NaCl(s)
Na+(aq) + Cl−(aq)AgNO3(s)
Ag+(aq) + NO3−(aq)If the two solutions are mixed:
Na+(aq) + Cl−(aq) + Ag+(aq) + NO3−(aq)
AgCl(s) + Na+(aq) + NO3−(aq)AgCl is insoluble and phases out as a white precipitate. As seen the sodium and nitrate ions don not participate in the reaction, and thus do not need to be included in the reaction equation. It is sufficient to write
Ag+(aq) + Cl−(aq)
AgCl(s)This is a model, and it isn't entirely correct. There is still a small amount of both Ag+ and Cl− in the solution, and also a small amount of AgCl(aq) will be present in the solution. In reality this is also an equilibrium, not an irreversible process, that looks like this:
Ag+(aq) + Cl−(aq)
AgCl(aq) AgCl(s)The reason for simplifying the reaction, as it is done, is that the amount of dissolved AgCl, Ag+ and Cl− is negligible and we focus on the precipitation. As soon as you start doing calculations on solubility products and multiple reactions influencing each other, you do use the reaction equation with the equilibria.
Precipitations comes from some units, in this case ions, having a high affinity for each other, and upon contact they are bound to each other, forming a compound that is insoluble in that particular solvent. In the classic chemistry for educational purposes, precipitations are generally associated with ions in aqueous solution, but it is important to keep in mind, that precipitations is far more than this.
The table below shows some combinations of ions that are soluble and insoluble. BUT, precipitations is not an either/or phenomenon, so, at low concentrations the precipitate may not appear, and at sufficiently high concentrations, precipitates may appear, even for salts that are considered soluble.
| Negative ion | Positive ion | Solubility |
| All | Li+, Na+, K+, NH4+ | Soluble |
| NO3− (nitrate) CH3COO− (acetate) | All | Soluble |
| Cl− (chloride) Br− (bromide) I− (iodide) | Ag+, Pb2+, Hg2+, Cu+ | Insoluble or almost insoluble |
| All the other | Soluble | |
| SO42− (sulfate) | Ag+, Ca2+, Ba2+, Pb2+ | Insoluble or almost insoluble |
| All the other | Soluble | |
| S2− (sulfide) | Li+, Na+, K+, NH4+, Mg2+, Ca2+, Sr2+, Ba2+ | Soluble |
| All the other | Insoluble or almost insoluble | |
| OH− (hydroxide) | Li+, Na+, K+, NH4+, Sr2+, Ba2+ | Soluble |
| All the other | Insoluble or almost insoluble | |
| CO32− (carbonate) PO43− (phosphate) | Li+, Na+, K+, NH4+ | Soluble |
| All the other | Insoluble or almost insoluble |
An important aspect about precipitations, that is easily overlooked, is that precipitations may only occur in a limited concentration interval, e.g. precipitation of Zn2+ with OH−:
Start:
Zn2+(aq) + 2 OH−(aq)
Zn(OH)2(s)At excess OH−:
Zn(OH)2(s) + 2 OH−(aq)
[Zn(OH)4]2−(aq)Therefore, you cannot just blindly mix two sets of ions, and from that conclude whether you have a precipitation or something else (in this case a complex formation).
Removal of the solvent
Another way of getting precipitation, is by removing the solvent. Removing solvent is the precipitation method people don't know that they know. The classic experiment in schools is the solution of copper(II)sulfate (the pentahydrate CuSO4 · 5 H2O). At exposure to the atmosphere, the water evaporates, and as the water evaoporates, big, blue crystals of copper(II)sulfate starts precipitating, when you reach the limit for how much copper(II)sulfate you can dissolve in the water.
This method is used both low tech, e.g. saltworks, where sea salt is produced by evaporating sea water, and high tech, for crystallization of macromolecules for x-ray crystallography (a type of analysis for seeing the 3D structure of large molecules).
This method is used both low tech, e.g. saltworks, where sea salt is produced by evaporating sea water, and high tech, for crystallization of macromolecules for x-ray crystallography (a type of analysis for seeing the 3D structure of large molecules).
Changing the solvent's properties
The alternative to removing the solvent is changing the solvent's properties. It can be something as simple as changing the temperature. For most chemicals, the solubility increases with temperature, and vice versa, so by lowering the temperature, you can get some chemicals to precipitate
Changes to pH or ionic strength are also fairly common methods for precipitations. The polymer poly(HEMA) is soluble in water but insoluble in salt water, provided the concentration is sufficiently high. Precipitations by adding salts are er commonly known as salting out.
Changing the solvent's hydrophilic/hydrophobic properties is yet another option. Zinc sulfate can be synthesized by adding a weak solution of sulfuric acid to elemental zinc. The result is an aqueous solution of zinc sulfate. Zinc sulfate is insoluble in ethanol, so by adding more and more ethanol to the solution, the solvent becomes more and more like ethanol, and the zinc sulfate precipitates.
Changes to pH or ionic strength are also fairly common methods for precipitations. The polymer poly(HEMA) is soluble in water but insoluble in salt water, provided the concentration is sufficiently high. Precipitations by adding salts are er commonly known as salting out.
Changing the solvent's hydrophilic/hydrophobic properties is yet another option. Zinc sulfate can be synthesized by adding a weak solution of sulfuric acid to elemental zinc. The result is an aqueous solution of zinc sulfate. Zinc sulfate is insoluble in ethanol, so by adding more and more ethanol to the solution, the solvent becomes more and more like ethanol, and the zinc sulfate precipitates.