Elektrolity: Reguła van ’t Hoffa

Electrolytes dissociate into multiple ions when dissolved, while non-electrolytes remain as single molecules. A 1 M potassium chloride solution produces approximately 2 moles of ions per liter, compared to 1 mole of dextrose molecules in a 1 M non-electrolyte solution. This doubling of solute particles causes electrolyte solutions to exhibit stronger colligative effects, such as twice the osmotic pressure of equivalent non-electrolyte solutions.

The van't Hoff factor (i) is the ratio of dissolved solute particles to the number of formula units added to solution. It is calculated by dividing the measured value of a colligative property by the value predicted from the formula. For ideal complete dissociation, potassium chloride would have i = 2, but measured values are typically lower due to ion pairing and incomplete dissociation in real solutions.

When electrolytes dissolve, some cations and anions recombine in solution through a phenomenon called ion pairing. This reduces the effective number of dissolved particles below the theoretical maximum. For a 0.100 m potassium chloride solution, the measured freezing point depression is 0.344 °C instead of the predicted 0.372 °C because ion pairing decreases the actual particle count in solution.

Strong electrolytes like iron(III) chloride and magnesium sulfate form strong electrostatic interactions and ion pairs, reducing their van't Hoff factors below the ideal value. Weak electrolytes like ammonium hydroxide undergo incomplete dissociation into ions. Both types yield van't Hoff factors less than expected because fewer independent particles exist in solution than the formula suggests.

As solutions become more dilute, ions separate more widely and residual interionic attractions decrease. In extremely dilute solutions, ion activities approach their actual concentrations, and the van't Hoff factor approaches its ideal value. For example, a 0.05 m sodium chloride solution has i = 1.9, closer to the ideal value of 2 than more concentrated solutions would exhibit.

Solvation of ions and the insulating action of polar water greatly reduce interionic attractions, but do not completely eliminate them. Residual attractions prevent ions from behaving as totally independent particles, allowing positive and negative ions to occasionally touch and form solvated units called ion pairs. This incomplete independence explains why strong electrolytes show lower van't Hoff factors than theoretical predictions.

Osmotic pressure depends on the total number of dissolved particles. A 1 M potassium chloride solution with i ≈ 2 produces twice the osmotic pressure of a 1 M non-electrolyte solution because it contains approximately twice as many solute particles. The van't Hoff factor quantifies this particle multiplication effect, allowing accurate calculation of osmotic pressure and other colligative properties for electrolyte solutions.