\(\mathrm{HNO}_{3}\) is a strong acid, whereas \(\mathrm{HNO}_{2}\) is a weak acid. Likewise, \(\mathrm{H}_{2} \mathrm{SO}_{4}\) is a strong acid, whereas \(\mathrm{H}_{2} \mathrm{SO}_{3}\) is a weak acid. Likewise, \(\mathrm{HClO}_{3}\) is a strong acid, whereas \(\mathrm{HClO}_{2}\) and \(\mathrm{HClO}\) are weak acids. (a) After water, what would be the predominant species present in a solution of \(\mathrm{HNO}_{3} ?\) (b) After water, what would be the predominate species present in a solution of \(\mathrm{HNO}_{2} ?\) (c) For acids of the general formula \(\mathrm{HXO}_{n}\) where \(\mathrm{X}\) is a nonmetal atom (these are called oxy-acids), how does the acid strength depend on the value of \(n\) ?

Understanding the difference between strong and weak acids is crucial to grasping the concepts of acidity and pH. A strong acid is defined by its ability to completely dissociate into its ions in a solution. This means that when a strong acid is added to water, it will release all of its hydrogen ions, leaving no molecules of the original acid intact. An example of a strong acid is hydrochloric acid (\( \mathrm{HCl} \)), which dissociates into hydrogen (\( \mathrm{H}^+ \)) and chloride ions (\( \mathrm{Cl}^- \)).

On the other hand, a weak acid only partially dissociates in water, which means that it releases some hydrogen ions into the solution, but also leaves a significant number of undissociated acid molecules. Acetic acid (\( \mathrm{CH}_3\mathrm{COOH} \)), commonly found in vinegar, is a classic example; it does not release all of its hydrogen ions when dissolved in water. This partial dissociation is what makes weak acids less corrosive and less conductive compared to strong acids.

The process by which an acid releases hydrogen ions into a solution is known as acid dissociation. This process is best represented by a chemical equation showing the acid (\( \mathrm{HA} \)) breaking apart into its constituent ions, hydrogen (\( \mathrm{H}^+ \)) and the conjugate base (\( \mathrm{A}^- \)). For strong acids, this reaction goes to completion, signified by a single arrow:\[ \mathrm{HA} \rightarrow \mathrm{H}^+ + \mathrm{A}^- \].

For weak acids, the dissociation is represented as an equilibrium, shown by a double arrow, indicating that the reaction can proceed in both forward and reverse directions:\[ \mathrm{HA} \leftrightarrows \mathrm{H}^+ + \mathrm{A}^- \].

The equilibrium position of this reaction is crucial, as it determines the concentration of hydrogen ions in the solution, which in turn affects the pH. The strength of an acid is often expressed by its acid dissociation constant (\( K_a \)), which is a numerical value that indicates how well an acid dissociates in solution. A large \( K_a \) denotes a strong acid, while a smaller \( K_a \) indicates a weaker acid.

Oxy-acids are a class of acids that contain hydrogen, oxygen, and another element, typically a nonmetal, forming the central atom. The general formula for an oxy-acid is \( \mathrm{HXO}_n \), where \( X \) is the central atom and \( n \) represents the number of oxygen atoms. These acids are noteworthy because their acid strength is heavily influenced by the number of oxygen atoms present.

The presence of more oxygen atoms increases the acid's strength by drawing electron density away from the hydrogen-oxygen bond, effectively making it weaker and more likely to release the hydrogen ion. This phenomenon is due to the electronegativity of oxygen, which pulls shared electrons towards itself, enhancing the acid’s ability to ionize. Some common examples of oxy-acids are sulfuric acid (\( \mathrm{H}_2\mathrm{SO}_4 \)), which is a strong acid, and sulfurous acid (\( \mathrm{H}_2\mathrm{SO}_3 \)), which is a weaker acid comparatively.

In chemical reactions, chemical equilibrium is the state where the rate of the forward reaction equals the rate of the reverse reaction. In the context of acids, this equilibrium is paramount for understanding weak acids and their dissociation in water. The reversible nature of weak acid dissociation sets up an equilibrium condition in the solution, characterized by specific concentrations of the reactants and products that remain constant at a given temperature — as long as the system is closed and no additional substances are added.

The equilibrium constant (\( K \)) for a reaction is the ratio of the products' concentrations to the reactants' concentrations, each raised to the power of their stoichiometric coefficients. For the dissociation of a weak acid \( HA \), the equilibrium constant (\( K_a \)) is given by:\[ K_a = \frac{\left[ \mathrm{H}^+ \right]\left[ \mathrm{A}^- \right]}{\left[ \mathrm{HA} \right]} \].

This value is especially important, as it helps in predicting how much of an acid will dissociate, determining both the pH of the solution and the concentrations of the constituents at equilibrium. By applying these principles, one can predict not just the behavior of acids and bases in solution, but equilibrium positions of a wide array of chemical reactions.