Problem 9 - Chemistry: Principles and Reactions

Question

Using the Brønsted-Lowry model, write equations to show why the following species behave as weak acids in water. (a) \(\mathrm{Ni}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{OH}^{+}\) (b) \(\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{3+}\) (c) \(\mathrm{H}_{2} \mathrm{~S}\) (d) \(\mathrm{HPO}_{4}^{2-}\) (e) \(\mathrm{HClO}_{2}\) (f) \(\mathrm{Cr}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5}(\mathrm{OH})^{+}\)

Step-by-Step Solution

Answer

Answer: The given species act as weak acids in water because they donate a small amount of protons (H+) when they partially dissociate in water, forming their respective conjugate bases.

Step 1: Identifying the Protons to be Donated

For each given species, identify the proton that can be donated according to the Brønsted-Lowry model.

Step 2: Writing the Equations

Write the equations for all species, showing how they react with water and form their respective conjugate bases. (a) \(\mathrm{Ni}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{OH}^{+}\) \(\mathrm{Ni}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{OH}^{+} + \mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{Ni}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{2+} + \mathrm{OH}^{-}\) (b) \(\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{3+}\) \(\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{3+} + \mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{OH}^{2+} + \mathrm{H}_{3} \mathrm{O}^{+}\) (c) \(\mathrm{H}_{2} \mathrm{S}\) \(\mathrm{H}_{2} \mathrm{S} + \mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{HS}^{-} + \mathrm{H}_{3} \mathrm{O}^{+}\) (d) \(\mathrm{HPO}_{4}^{2-}\) \(\mathrm{HPO}_{4}^{2-} + \mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{PO}_{4}^{3-} + \mathrm{H}_{3} \mathrm{O}^{+}\) (e) \(\mathrm{HClO}_{2}\) \(\mathrm{HClO}_{2} + \mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{ClO}_{2}^{-} + \mathrm{H}_{3} \mathrm{O}^{+}\) (f) \(\mathrm{Cr}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5}(\mathrm{OH})^{+}\) \(\mathrm{Cr}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5}(\mathrm{OH})^{+} + \mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{Cr}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{3+} + \mathrm{OH}^{-}\) These equations illustrate how each species behaves as a weak acid in water, by donating a proton (H+) and forming their respective conjugate bases.

Key Concepts

Weak Acids

In the realm of the Brønsted-Lowry model, acids are defined by their ability to donate protons, commonly represented as hydrogen ions ( H^+ ) . Weak acids are a specific category within acids that only partially dissociate in water, meaning they give off some but not all of their protons. This characteristic makes weak acids distinct and affects their behavior in solution.

• They often have a low degree of ionization, indicating that only a small fraction of the acid particles donate protons.
• The dissociation of weak acids results in a dynamic equilibrium, not a complete transformation.

Due to their partial dissociation, weak acids like Ni( H _2O) _5OH^+ and Al( H _2O) _6^{3+} demonstrate reversible reactions with water, shown in the provided equations. This means that in solution, there is always a balance between the acid and its ions, making weak acids less robust than strong acids in their acidifying effect.

Conjugate Base

A conjugate base forms when a Brønsted-Lowry acid donates a proton to the surrounding medium. The concept of conjugate bases is essential for understanding how weak acids behave in water and the equilibrium they achieve.

• When a weak acid gives up a proton, it forms its corresponding conjugate base. This is the remaining part of the acid after losing a H^+ ion.
• The strength of the conjugate base is inversely related to that of the acid—strong acids have weak conjugate bases, while weak acids have relatively stronger ones.

For example, when Al( H _2O) _6^{3+} donates a proton, it forms the conjugate base Al( H _2O) _5OH^{2+} . The presence of a conjugate base indicates that equilibrium is attainable, as it can react with water to regenerate the original acid species.

Proton Donation

The process of proton donation is central to the behavior of acids in water, according to the Brønsted-Lowry theory. This model conceptualizes acids as proton donors and bases as proton acceptors.

• Proton donation involves the transfer of an H^+ ion from the acid to a base, often water in aqueous solutions.
• This transfer helps establish the acid-base equilibrium by altering the concentrations of ions in the solution.

Taking HClO _2 as an example, it donates a proton to water, forming a H _3O^+ ion and leaving behind a ClO _2^- ion, the conjugate base. This showcases how the proton donation sets the stage for equilibrium, determining the acid's reactivity and influence on the pH level.

Acid-Base Equilibrium

Acid-base equilibrium describes the balance achieved when acids donate protons and bases accept them, reaching a state where forward and reverse reactions occur at the same rate. This equilibrium is crucial for the stability and function of chemical systems, especially for weak acids.

• In weak acid solutions, equilibrium involves both the undissociated acid molecules and their dissociated ions.
• The equilibrium position varies, influenced by the strength of the acid, the concentration of the reactants and products, and the temperature.

The reactions involving HPO _4^{2-} and H _2S illustrate how these species establish an equilibrium with water. They produce hydronium ions and conjugate bases, indicating a dynamic balance of dissociation and reassociation. Understanding this equilibrium is essential for predicting the behavior of acids in various conditions and their subsequent chemical reactivity.