Electrolytic dissociation

The electrolytic dissociation means the reversible decomposition of a compound in anions and cations in a solvent . Such solutions are called electrolytes . Such a dissociation takes place, for example, when salts are dissolved in water. In this polar solvent, what is dissolved is in the form of freely moving ions . These freely moving anions and cations lead to an electrical conductivity of the water.

history

Already in 1795, after the discovery of electrochemical electricity by Luigi Galvani, Alessandro Volta was considering a charge separation of particles by different metals . In 1805, Theodor Grotthuss developed a theory for the charge separation of water molecules. Berzelius and Daniell thought about salts in water, they assumed that the salts separated into positive and negative charge carriers. Faraday coined the term ion, but he did not yet have the current idea of ​​a localized charge on an atom or part of a molecule in solution. Rudolf Clausius believed that the charge carriers are not firmly connected, but rather vibrate. Even then he recognized the possibility of electrolytes dissociating. Since the first electricity was only obtained from electrochemical batteries and the basic laws of electricity and electronics were developed with these batteries, the question of electrolytes and the conductivity of the solutions was of great importance.

Fundamental work on the actual dissociation, however, was only carried out by Friedrich Wilhelm Georg Kohlrausch (production of a conductivity meter with alternating current for liquids, formulation of a law for strong electrolytes, kohlrausches square root law ), Jacobus Henricus van 't Hoff ( osmotic pressure depending on the number of particles), Svante Arrhenius (work on conductivity with Kohlrausch's conductometer and postulate of the dissociation of salts into positive and negative charge carriers, founder of the dissociation theory) and Wilhelm Ostwald (dissemination of Arrhenius ideas, application of the law of mass action to dissociation and establishment of the important law for weak electrolytes, Ostwald's Law of Dilution ).

A mathematical theory of dissociation was later developed by Petrus Debye and Erich Hückel. With this Debye-Hückel theory , the degrees of dissociation of the electrolytes can be mathematically calculated from known limit conductivities. However, the model is only suitable for low electrolyte concentrations (up to 0.01 mol / l).

Max von Laue was able to show through X-ray structure analysis that ions are also present in solid salts.

Chemical definition

Electrolytic dissociation is the reversible breakdown of a compound into anions and cations in a solvent . The proportion of dissociated ions in relation to the total content of undissociated and dissociated ions of the same type is called the degree of dissociation (see also activity ). It depends on the concentration in the solvent. At a very high dilution, the acetic acid is also completely dissociated into acetate and hydronium ions, at a high concentration there is still a large proportion of undissociated acetic acid. The dissociation, or more precisely the degree of dissociation of salts or organic molecules, can be determined by electrical conductivity measurements ( conductometry ) and by pH measurements of aqueous solutions.

Sodium chloride (NaCl), which is dissolved in tap or distilled water, dissolves in the form of positive sodium cations and negative chlorine anions. These ions are dissociated, i.e. i.e., spatially separated from one another and with an electrical charge. Such solutions are called electrolytes . NaCl is a strong electrolyte because the ions are almost completely separated in the solution. Vinegar essence , which consists of the ingredients acetic acid and water, is also partially in the form of ions, namely as oxonium ions and acetate anions . Acetic acid is not completely dissociated into ions, only approx. 0.3% of the acetic acid molecules (at 1 mol / l) are in dissociated form. Acetic acid is a weak electrolyte. Depending on the type of dissolved particles, there are all transitions between strong and weak electrolytes.

In the case of electrolytes, the equilibrium constant from dissociated products to undissociated starting materials can be determined from the degree of dissociation and the law of mass action (MWG). Wilhelm Ostwald formulated that the product of the active mass concentrations of the dissociated particles (in the case of acetic acid, the acetate concentration multiplied by the hydronium concentration) by the concentration of the undissociated mass particles (undissociated acetic acid) is always a constant. This constant is called the dissociation constant (outdated: affinity constant) and it is used, for example, to determine the pKa values ​​of acids. With the equilibrium constant K _a (a abbreviation for acid ), or as a negative decadic logarithm (−logK _a = pK _a ), the proportion of dissociated ions can be precisely determined for each concentration.

${\displaystyle \mathrm {AB\ \rightleftharpoons \ A^{+}\ +\ B^{-}} }$.

${\displaystyle \mathrm {NaCl+Wasser\rightarrow Na^{+}+Cl^{-}+Wasser\quad |k_{1}=Geschwindigkeitskonstante1} }$

${\displaystyle \mathrm {Na^{+}+Cl^{-}+Wasser\rightarrow NaCl+Wasser\quad |k_{2}=Geschwindigkeitskonstante2} }$

${\displaystyle {\tfrac {k_{1}}{k_{2}}}=K_{a}}$ .

K _a > 1 applies to strong electrolytes (table salt in water).

In the case of acetic acid in water, the equilibrium is on the left. We have K _a <1.

${\displaystyle \mathrm {H_{3}C{-}COOH+H_{2}O\ \rightleftharpoons \ H_{3}C{-}COO^{-}\ +\ H_{3}O^{+}} }$

${\displaystyle \mathrm {H_{3}CCO_{2}H+H_{2}O\rightarrow H_{3}CCO_{2}^{-}+H_{3}O^{+}\quad |k1=Geschwindigkeitskonstante1} }$

${\displaystyle \mathrm {H_{3}CCO_{2}^{-}+H_{3}O^{+}\rightarrow H_{3}CCO_{2}H+H_{2}O\quad |k2=Geschwindigkeitskonstante2} }$

If the gas hydrogen chloride (HCl) is introduced into water, an electrolytic solution is formed, which is called hydrochloric acid :

${\displaystyle \mathrm {H_{2}O\ +\ HCl\ \rightleftharpoons \ H_{3}O^{+}\ +\ Cl^{-}} }$.

If the gas ammonia (NH ₃ ) is introduced into water, NH ₄ ⁺ is formed as a cation and OH ^- as an anion

${\displaystyle \mathrm {NH_{3}+H_{2}O\rightleftharpoons \ NH_{4}^{+}+OH^{-}} }$

The equilibrium reactions of these last three examples are also called protolysis ; it is described in more detail in the article of the same name. This behavior turns acetic and hydrochloric acid into acids . The behavior of ammonia makes ammonia a base . The electrical conductivity of these solutions is the experimental proof of the formation of freely moving anions and cations.

A very important special case of dissociation is the equilibrium of pure water, distilled in the absence of air. Water dissociates to a very small extent into hydroxide and oxonium ions. Kohlrausch and Heydweiller determined the conductivity of distilled water in 1894 to be 0.06 · 10 ⁻⁶  Ω ⁻¹ . The equilibrium constant, the ionic product of water , can be calculated from the conductivity and the knowledge of the limit conductivity (see equivalent conductivity ) for hydronium and hydroxide ions . It has a value of K _W = 10 ⁻¹⁴  mol ² / L ² or pK _W = 14. Knowing the concentration of electrolytes and the corresponding pK _a values ​​from tables can easily determine the pH value and degree of dissociation for each electrolyte .

When dissolving exactly 100.0 g of 98% sulfuric acid in exactly 962.7 g of distilled water, the concentration of the sulfuric acid is exactly 1 mol / L. If you take 106.38 g of the sulfuric acid solution and dissolve it in 900 mL of distilled water, this solution has a concentration of 0.1 mol / L, if you take 10.64 g of the first solution and dissolve it in 990 mL of distilled water the concentration 0.01 mol / L, a 0.001 mol / L aqueous sulfuric acid solution can be prepared analogously. The sulfuric acid can be dissociated in water, i.e. This means that the hydrogen atoms split off from the sulfuric acid as positively charged oxonium ions and hydrogen sulfate ions or sulfate ions are formed as counterions. The strength of the dissociation is the pK _s determined value. Acids having a negative pK _s value are always completely dissociated before. The lower the pK _s value, the stronger the acid is dissociated.

Let us assume that the 0.1 molar sulfuric acid would be completely dissociated at this concentration, then both hydrogen atoms would have to be split off from the sulfuric acid and the oxonium concentration would have to be 0.2 mol / L. The equivalent concentration (or normality) of sulfuric acid would be: 2 · 0.1 mol / L = 0.2 mol / L. In relation to 1000 g of water molecules, 3.6 g of these water molecules could be present as oxonium ions. The pH would therefore be

${\displaystyle pH=-\log(H+)=-\log(0{,}2)=0{,}7}$.

Similarly, for oxonium concentrations of 0.02 mol / L: pH = 1.7 and 0.002 mol / L: pH = 2.7. However, the following pH values ​​were measured: 0.2 mol / L: pH = 1.23, 0.02 mol / L: pH = 1.94, 0.002 mol / L: pH = 2.78. In fact, the measured pH value is only halfway correct for the 0.002 mol / L solution. Only there is the sulfuric acid completely dissociated. In the more concentrated solutions, the approximate formula:

${\displaystyle pH=(pK_{\text{s}}-\log C)/2}$

(with C for the concentration in mol / L * equivalent number) can be used according to Ostwald's law of dilution to determine the pH dependency. For sulfuric acid, 0.2 mol / L: pH = 1.30 (measured: 1.23), 0.02 mol / L: pH = 1.80 (measured: 1.94) is obtained. At a higher concentration, a high proportion of sulfuric acid is therefore present as hydrogen sulfate (over 40% in a 0.1 molar sulfuric acid) and not as a doubly negative sulfate, which explains the deviations in the pH measurements. For hydrochloric acid and nitric acid, the logarithmic concentration dependency applies when determining the pH in accordance with the first relationship, for sodium hydroxide solution in an analogous manner for the hydroxide concentration. When titrating with sodium hydroxide solution or when further diluting the solution, the sulfuric acid particles are dissociated and the correct total concentration is obtained.

Weak acids and bases can act as buffers . The pH value of a solution remains fairly constant and corresponds to the pKa value of the corresponding acid if a weak acid and its anion are in almost the same concentration in the solution. Important biological buffers are the carbonic acid-bicarbonate system and the dihydrogen phosphate buffer, they ensure a constant pH value in biological systems. With acid-base titrations , a suitable pH indicator or pH electrode must be used in such solutions in order to correctly determine the concentration of the acid and base sought.

Salts can also dissociate to a greater or lesser extent. Salt solutions then show different properties at high concentrations than would be expected if the physical measured value were extrapolated from dilute solutions to solutions with high concentrations. The physical properties can be, for example, the freezing point, the boiling point or the electrical conductivity. To describe a solution that has a linear dependency on a hypothetical concentration ( activity ), the concentration of a solution must be multiplied by a factor, the activity coefficient. To describe the concentration-dependent equivalent conductivity, the conductivity coefficient according to Kohlrausch's law is used; this differs significantly from the activity coefficient.

With the so-called real or permanent electrolytes , the ions are already present in the solid body ( ion lattice ). In the case of solid table salt, Na ⁺ and Cl ^- ions are already present in the lattice . When the salt is dissolved in water, freely moving ions are now formed in the water. With the dissociation of salts into ions, the very high lattice energy of the crystal is applied by hydration energy during the dissolution process.

In the so-called potential electrolytes, there are no ionic bonds in the pure substances. As a pure substance, they are non-conductors . When these pure substances (AB) are introduced into a solvent, ions are formed through a chemical reaction between the dissolved matter and the solvent:

${\displaystyle \mathrm {AB\ \rightleftharpoons \ A^{+}\ +\ B^{-}} }$.

A prerequisite for such a reaction is a polar bond between parts A and B of the compound (AB) and a polar solvent. If, for example, pure acetic acid is added to water, the anions H ₃ C – COO ^- are formed as cations H ₃ O ⁺

Ions are not only dissociated in water. An almost equally high degree of dissociation is also observed in polar organic solvents such as formamide, acetonitrile or nitromethane. The dielectric constant is decisive for the dissociation in organic solvents, as Walter Nernst discovered.

Individual evidence

1. ↑ Rudolf Clausius: About the conduction of electricity in electrolytes . In: Pogg. Ann . tape 101 , 1857, pp. 338-360 ( fakesimile in UrMEL of ThULB).
2. ↑ Svante Arrhenius: About the dissociation of substances dissolved in water . In: Journal of physical chemistry . 1st vol., No. 11-12, 1887, pp. 631-648.
3. ↑ ^{a b} Friedrich Wilhelm Ostwald: About the dissociation theory of electrolytes . In: Z. f. physics. Chemistry . tape 2 , 1888, p. 270-283 .
4. ↑ Journal for phys. Chemistry . tape 69 , 1909, pp. 1 ff .
5. ↑ Hans Bouma, Walter Jansen (Ed.): Handbuch der experimental Chemie. Secondary level II. Volume 6: Elektrochemie, Aulis Verlag Deubner & Co. KG, Cologne 1994, ISBN 3-7614-1630-X , p. 24.
6. ↑ Journal f. physics. Chemistry . tape 14 , 1894, pp. 317 .
7. ^ W. Nernst: Dielectric constant and chemical equilibrium. In: Z. phys. Chem. 13, No. 3, 1894, pp. 531-536.