Which of these statements is/are true regarding a 1.0 \(M\) solution of a strong acid HA? (a) \(\left[\mathrm{A}^{-}\right]>\left[\mathrm{H}^{+}\right]\) (b) The \(\mathrm{pH}\) is 0.00 . (c) \(\left[\mathrm{H}^{+}\right]=1.0 \mathrm{M}\) (d) \([\mathrm{HA}]=1.0 \mathrm{M}\).

When discussing the pH of strong acids, one fundamental concept to grasp is that the pH scale is a measure of the acidity or basicity of an aqueous solution. For strong acids, like the hypothetical HA in our exercise, pH calculation is straightforward. The strength of an acid refers to its ability to dissociate in water and strong acids fully dissociate, releasing a high concentration of hydrogen ions \(H^+\).

The pH is calculated by taking the negative logarithm (base 10) of this concentration: \(pH = -\log_{10}\left[\text{H}^+\right]\). A 1.0 M solution of a strong acid, which dissociates completely, results in a \(pH\) of exactly 0 since \(\log_{10}(1) = 0\). This means the solution is highly acidic, as a low pH indicates a high concentration of hydrogen ions.

The process of acid dissociation in water is critical to understanding strong acids like the one given in our exercise. Dissociation is the separation of an acid into its constituent ions. In the case of strong acids, this dissociation is complete, meaning that when a strong acid HA dissolves in water, it breaks apart into its \(H^+\) ions and corresponding anions \(A^-\).

Symbolically, we represent this reaction as: \[HA \rightarrow H^+ + A^-\]. Because strong acids completely dissociate in water, the concentration of the anion \(A^-\) will equal the initial concentration of the acid HA. Therefore, if we start with a 1.0 M solution of HA, we will end up with 1.0 M \(H^+\) and 1.0 M \(A^-\), assuming there are no other interactions in the solution.

Understanding hydrogen ion concentration is fundamental to acid-base chemistry. The concentration of hydrogen ions \(\left[\text{H}^+\right]\) in a solution determines its acidity. Higher \(\left[\text{H}^+\right]\) correlates with lower pH and therefore a more acidic solution. In our exercise, the complete dissociation of a strong acid in water leads to a \(\left[\text{H}^+\right]\) equal to the initial molarity of the acid. In other words, a 1.0 M HA solution results in a 1.0 M concentration of \(H^+\) ions.

The strict one-to-one dissociation ratio is why strong acids are considered strong; they completely give up their hydrogen ions into the solution. This concept is crucial in predicting the behavior of various reactions and calculating the pH in quantitative terms, which has numerous practical applications in chemistry and related fields.