6.9. IONIC EQUILIBRIUM IN SOLUTION

Under the effect of change of concentration on the direction of equilibrium, you have incidentally come across with the following equilibrium which involves ions: Fe3+(aq) + SCN–(aq) [Fe(SCN)]2+(aq). There are numerous equilibria that involve ions only.

Ionic equilibrium is the balance between the reactions of ions in a solution. For example, when you mix iron(III) ions (Fe3+) and thiocyanate ions (SCN–), they form a complex ion called [Fe(SCN)]2+. This equilibrium means that there can be changes in the concentrations of these ions, which affects how the reaction proceeds.

Think of ionic equilibrium like a seesaw, where the balance can be tipped by adding more weight on one side. In our example, adding more Fe3+ or SCN– could shift the equilibrium towards creating more of the complex ion.

It is well known that the aqueous solution of sugar does not conduct electricity. However, when common salt (sodium chloride) is added to water it conducts electricity. Also, the conductance of electricity increases with an increase in concentration of common salt. Michael Faraday classified the substances into two categories based on their ability to conduct electricity.

When dissolved in water, substances fall into two categories: electrolytes, which can conduct electricity (like sodium chloride), and non-electrolytes, which cannot (like sugar). Strong electrolytes fully dissociate into ions, while weak electrolytes partially dissociate. For example, when sodium chloride dissolves in water, it separates completely into sodium ions and chloride ions, allowing the solution to conduct electricity.

Imagine a busy street where cars (ions) can move freely. If the street is clear (like in a solution of sodium chloride), the cars (electrical current) can drive easily, but if there are no cars (like sugar in water), no traffic (current) can happen.

Faraday further classified electrolytes into strong and weak electrolytes. Strong electrolytes on dissolution in water are ionized almost completely, while the weak electrolytes are only partially dissociated.

The key difference between strong and weak electrolytes lies in how completely they dissociate into ions. Strong electrolytes like sodium chloride dissociate completely into ions in solution, while weak electrolytes like acetic acid only partially dissociate, meaning that most of the acid remains in its molecular form, with only a few ions present.

You can think of strong electrolytes as a packed concert where everyone is dancing vigorously (complete dissociation). In contrast, at a quiet coffee shop where only a few people are chatting (weak electrolyte), the atmosphere is much calmer and less electric (fewer free ions).

It should be noted that in weak electrolytes, equilibrium is established between ions and the unionized molecules. This type of equilibrium involving ions in aqueous solution is called ionic equilibrium.

Ionic equilibrium in weak electrolytes indicates that even though weak acids don't fully dissociate, there exists a balance between the undissociated acid and its ions. For instance, in a solution of acetic acid, there are both the dissolved molecules of acetic acid and ions of acetate (CH3COO–) and hydronium (H3O+). This balance is crucial for understanding how weak acids behave in solution.

Imagine a seesaw again, but this time with a person on each side (the ionized molecules on one side and the undissociated molecules on the other). They can change places as they interact, representing how equilibrium works — while some acetic acid molecules turn into ions, some ions can also recombine to form acetic acid.

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