7.2
Associação Jônica
A associação iônica é a associação de íons com carga oposta em uma solução eletrolítica para formar pares de íons. Bjerrum definiu pares de íons como…
Eletrólitos fortes são frequentemente considerados inteiramente como íons livres em solução aquosa. Na realidade, íons com carga oposta podem se associar parcialmente para formar pares de íons, exceto em muitos eletrólitos 1:1, como NaCl ou NaOH.
Bjerrum definiu um par de íons como dois íons com cargas opostas suficientemente próximos para que sua atração eletrostática exceda a energia térmica, quantificada como 2kT, onde k é a constante de Boltzmann e T é a temperatura absoluta. Quando essa condição é atendida, a associação de íons torna-se favorável.
O emparelhamento de íons aumenta com cargas iônicas mais altas, como em eletrólitos 2:1 ou 2:2, levando a uma fração significativa de pares iônicos mesmo em concentrações baixas. Essa previsão é comprovada experimentalmente. Quando a porcentagem de cátions em pares de íons é representada em função da molalidade, os resultados correspondem de perto à teoria de Bjerrum.
Também é fortemente influenciado pelo solvente: a alta constante dielétrica da água enfraquece a atração eletrostática e limita o pareamento de íons, enquanto solventes com constantes dielétricas mais baixas aumentam a atração íon–íon, tornando a formação de pares de íons significativa mesmo para eletrólitos 1:1.
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Q1: What is an ion pair according to Bjerrum's definition?
An ion pair consists of two oppositely charged ions positioned close enough that their electrostatic attraction exceeds the thermal energy, quantified as 2kT, where k is Boltzmann's constant and T is absolute temperature. When this condition is met, the ions remain associated rather than freely dissolved in solution.
Q2: How does ionic charge affect ion pair formation?
Ion pairing increases significantly with higher ionic charges, such as in 2:1 or 2:2 electrolytes, leading to substantial ion-pair fractions even at low concentrations. In contrast, 1:1 electrolytes like NaCl show minimal ion pairing in aqueous solutions due to weaker electrostatic interactions between singly charged ions.
Q3: Why does water limit ion pair formation compared to other solvents?
Water's high dielectric constant weakens electrostatic attraction between oppositely charged ions, reducing ion-pair formation. Solvents with lower dielectric constants are less effective at stabilizing separated charges, resulting in stronger electrostatic attraction and increased ion pairing even for 1:1 electrolytes.
Q4: How does ionic association affect electrical conductivity?
Ionic association reduces electrical conductivity because ions form associated species like CaSO₄ and MgF₂, decreasing the number of free charge carriers in solution. The extent of association can be estimated from conductivity measurements, with association becoming significant in concentrated solutions and negligible at infinite dilution.
Q5: What is the difference between ion pairs and complex ions?
Ion pairs are held together by electrostatic forces and often retain part of their solvent shells, whereas complex ions involve bonds with significant covalent character, commonly forming in transition-metal salt solutions. Absorption spectroscopy can distinguish between these species, and solutions may contain both types simultaneously.
Q6: How do temperature and molality influence ionic association?
Both temperature and molality affect the extent of ionic association in solution. At infinite dilution, the degree of association approaches zero, while association becomes significant in concentrated solutions. Higher temperatures increase thermal energy, reducing favorable ion-pair formation conditions.
Q7: How does Bjerrum's theory compare to experimental observations of ion pairing?
Bjerrum's theory predictions closely match experimental data when the percentage of cations in ion pairs is plotted against molality. This agreement validates the theoretical framework that electrostatic attraction must exceed thermal energy for ion-pair formation, supporting the Debye–Hückel Theory of Electrolyte Solutions as a foundation for understanding nonideal behavior.