8.6
Ostwalds Verdünnungsgesetz
Betrachten wir einen binären Elektrolyten AB mit einer Konzentration 'c' , der reversibel in seine Bestandteile dissoziiert. Der Grad dieser Dissoziat…
Basierend auf Arrhenius' Theorie der elektrolytischen Dissoziation erklärt Ostwalds Verdünnungsgesetz das Gleichgewicht in wässrigen Elektrolytlösungen.
Betrachten wir einen Elektrolyten AB mit einer Konzentration von 'c' mol pro Liter, der reversibel in Ionen dissoziiert.
Wenn α der Dissoziationsgrad ist, hat jedes Ion eine Gleichgewichtskonzentration von cα, während der undissoziierte Anteil (1 − α) eine Konzentration von c(1 − α) besitzt. Diese Begriffe definieren die Gleichgewichtskonstante K, die Dissoziationskonstante genannt wird.
Ostwalds Verdünnungsgesetz gilt nur für schwache Elektrolyte wie Essigsäure und NH₄OH.
Starke Elektrolyte wie HCl und NaF sind fast vollständig dissoziiert, mit α nahe an einem. Unter diesen Bedingungen nähert sich (1 − α) null, wodurch das Gesetz mathematisch ungültig ist.
Das Versagen des Gesetzes für starke Elektrolyte entsteht, weil bei der Auflösung die hohe Dielektrizitätskonstante des Wassers die elektrostatischen Kräfte schwächt, was zu nahezu vollständiger Dissoziation führt. Daher existiert kein Ion-Molekül-Gleichgewicht, und das Ostwaldsche Verdünnungsgesetz wird nicht mehr anwendbar.
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Q1: What is Ostwald's dilution law and how does it relate to electrolyte dissociation?
Ostwald's dilution law explains equilibrium in aqueous electrolyte solutions based on Arrhenius' theory of electrolytic dissociation. For an electrolyte AB at concentration c, the dissociation constant K equals cα²/(1−α), where α is the degree of dissociation. This law describes how weak electrolytes like acetic acid and NH₄OH reversibly dissociate into ions at equilibrium.
Q2: Why does Ostwald's dilution law fail for strong electrolytes?
Strong electrolytes like HCl and NaF are almost completely dissociated, with α approaching one, making (1−α) approach zero. This renders the law mathematically invalid. Water's high dielectric constant weakens electrostatic forces, causing near-complete dissociation and eliminating the ion-molecule equilibrium that the law requires.
Q3: How do you calculate equilibrium concentrations using the degree of dissociation?
For an electrolyte AB with concentration c and degree of dissociation α, each ion has equilibrium concentration cα, while the undissociated fraction has concentration c(1−α). These concentrations define the dissociation constant K, which governs the equilibrium state of weak electrolyte solutions.
Q4: What is the difference between weak and strong electrolytes in terms of dissociation?
Weak electrolytes like acetic acid partially dissociate, with α significantly less than one, allowing Ostwald's dilution law to apply. Strong electrolytes like NaCl and NaF are electrovalent compounds formed by electron transfer; they dissociate almost completely with α near one, making the law inapplicable to them.
Q5: How does the dielectric constant of a solvent affect electrolyte dissociation?
According to Coulomb's law, the dielectric constant of the medium affects electrostatic force strength and conductivity. Water's high dielectric constant weakens electrostatic forces between ions, promoting near-complete dissociation of strong electrolytes and enabling ion mobility for electrical conduction in solutions.
Q6: What is the simplified form of Ostwald's dilution law for very weak electrolytes?
For very weak electrolytes where α is much less than one, the term (1−α) approximates to one. The dissociation constant simplifies to K ≈ cα², making calculations more straightforward for solutions with minimal dissociation and allowing easier prediction of equilibrium behavior.
Q7: What structural difference exists between strong electrolytes and weak electrolytes?
Strong electrolytes like NaCl and NaF are electrovalent compounds composed solely of ions in their crystal structure, not molecules. When dissolved in water, these ions gain mobility and conduct electricity. Weak electrolytes exist as molecules that partially dissociate into ions upon dissolution in solution.