Salt effect

As salting- or salting is known in the chemistry to increase the solubility of a substance in water , when salts are added. A saline solution can e.g. B. usually dissolve another salt better than pure water.

The term is often used in connection with proteins . By adding small amounts of salt, protein aggregation (connection via amino acid residues ) is suppressed, and the resulting enlarged capillary allows more water to be absorbed.

There is also a salting-out effect that can occur with dissolved non-electrolytes after adding salt. In the case of proteins, the salt ions compete with the protein for free water due to a high amount of salt , which reduces the protein's water binding.

The effect of salts with regard to the salting-in and salting-out effect on biomolecules is characterized by the Hofmeister series .

Clear explanation

In the saturated solution of a salt, some of the undissolved sediment still goes into solution, while an equally large part precipitates again . In order to be able to precipitate, a cation and an anion of the associated ion pair must “meet” in the solution. If another salt is dissolved, the likelihood of this meeting will decrease. Precipitation is hindered and more salt can go into solution.

Physico-chemical derivation

The measure of the solubility of a salt is the solubility product of its ions. It is often given as the product of the concentrations : ${\displaystyle L_{\mathrm {AB} }}$ ${\displaystyle c}$

${\displaystyle L_{\mathrm {AB} }=c_{A}^{+}\cdot c_{B}^{-}}$

Strictly speaking, the concentrations can only be used for infinitely diluted ( ideal ) solutions. In real solutions, instead of concentration, there is activity : ${\displaystyle a}$

${\displaystyle L_{\mathrm {AB} }=a_{A}^{+}\cdot a_{B}^{-}}$

The activity of an ion is the product of its activity coefficient and the concentration : ${\displaystyle f_{i}}$ ${\displaystyle c_{i}}$

${\displaystyle a_{i}=f_{i}\cdot c_{i}}$

The activity coefficient for already dissolved ions is reduced according to the Debye-Hückel law, since the ionic strength increases when other salts are added: ${\displaystyle I}$

${\displaystyle \lg f_{i}=-A\cdot z_{i}^{2}\cdot {\sqrt {I}}}$

With

• temperature-dependent constant , e.g. B. 0.509 at 25 ° C ${\displaystyle A}$
• Charge number of the respective type of ion ${\displaystyle z_{i}}$

The Debye-Hückel law is used to calculate the mean activity coefficient of monovalent electrolytes in solutions with an ionic strength of less than 0.01 mol / l.

The activity coefficient for ions of all charge numbers can be determined even in solutions of several electrolytes with an ionic strength of up to 0.5 mol / l using the Davies approximation:

${\displaystyle \lg f_{i}=-A\cdot z_{i}^{2}\cdot \left({\frac {\sqrt {I}}{1+{\sqrt {I}}}}-0,3\cdot I\right)}$