Question

Which one of the following statements is not correct? (a) \(\mathrm{Cl}^{-}\) is a lewis acid. (b) The \(\mathrm{pH}\) of \(10^{-8} \mathrm{M} \mathrm{HCl}\) solution is less than \(7 .\) (c) The ionic product of water at \(25^{\circ} \mathrm{C}\) is \(10^{-14} \mathrm{~mol}^{2} \mathrm{~L}^{-2}\) (d) Bronsted-Lowry theory could not explain the acidic character of \(\mathrm{AlCl}_{3}\)

Short answer

Statement (a) is incorrect: \( \text{Cl}^- \) is not a Lewis acid.

Step-by-step solution

Understand Lewis Acid

A Lewis acid is a chemical species that accepts an electron pair. The chloride ion, \( \text{Cl}^- \), already has an extra electron, making it unlikely to accept another pair of electrons. Hence, \( \text{Cl}^- \) is not a Lewis acid, making statement (a) incorrect.

Examine pH of a weak HCl solution

A 10\(^{-8}\) M \( \text{HCl} \) solution seems like it would have a pH of 8, making it more basic than neutral. However, because \( \text{HCl} \) completely ionizes and the concentration is extremely low, the pH is determined by the contribution from water's autoionization, making it slightly acidic (<7). Therefore, statement (b) is correct.

Ionic Product of Water

The ionic product of water at \( 25^{\circ} \mathrm{C} \) is indeed \( 10^{-14} \text{ mol}^2 \text{ L}^{-2} \). This is a well-known fact, affirming that statement (c) is correct.

Bronsted-Lowry and AlCl3

The Bronsted-Lowry theory describes acids as proton donors and bases as proton acceptors. However, \( \text{AlCl}_3 \) acts as an acid by accepting electron pairs (it's a Lewis acid), not by donating protons. Hence, Bronsted-Lowry theory cannot explain its acidic behavior, thus statement (d) is true.

Key concepts

Lewis Acids and Bases

In chemistry, Lewis acids and bases open the door to understanding how various reactions work through the concept of electron pairs. A Lewis acid is any substance that accepts a pair of electrons. This means that the particle doesn't have enough electrons and is eager to gain some. On the contrary, a Lewis base donates a pair of electrons because it has a little extra to spare.
The chloride ion, \( \text{Cl}^- \), serves as a perfect counterexample of a Lewis acid. Already packed with electrons, this ion is not interested in hoarding more. It's like trying to convince your friend to take more candies when they already have their pockets full—it just doesn't work!
By understanding what makes a Lewis acid or base, we can predict how different elements and compounds will interact in various chemical environments.

pH Calculation

Calculating the pH of a solution helps determine its acidity or basicity. A basic pH calculation might consider just the concentration of an acid or base. However, things get interesting when dealing with very dilute solutions, such as a \( 10^{-8} \text{ M} \) \( \text{HCl} \) solution.
On first glance, one might assume such a dilute solution is basic due to its low concentration. But because \( \text{HCl} \) is a strong acid, it completely ionizes in water. In addition, even water itself can influence pH due to its own ionization process.
In extremely dilute solutions, the proton concentration arising from the autoionization of water cannot be ignored, leading to a slightly acidic pH value.
This showcases how pH measurement isn't always straightforward, especially in dilute conditions.

Ionic Product of Water

The ionic product of water reflects a crucial aspect of water's chemical nature, showcasing its ability to partially ionize into hydrogen ions (\( \text{H}^+ \)) and hydroxide ions (\( \text{OH}^- \)) even at room temperature. At \( 25^{\circ} \text{C} \), the ionic product \( K_w \) is precisely \( 10^{-14} \text{ mol}^2 \text{ L}^{-2} \).
This constant result highlights a balance in pure water: the addition of either acid or base will disrupt this equilibrium by altering the concentrations of \( \text{H}^+ \) and \( \text{OH}^- \). If more \( \text{H}^+ \) ions are added, the concentration of \( \text{OH}^- \) must decrease to maintain the product at \( 10^{-14} \).
Understanding the ionic product of water is fundamental when contemplating water’s behavior in different chemical reactions and environments.

Bronsted-Lowry Theory

The Bronsted-Lowry theory reshapes our understanding of acids and bases by defining them as proton donors and acceptors, respectively. This framework provides insight into how substances interact with each other.
For instance, when \( \text{HCl} \) dissolves in water, it donates a proton (\( \text{H}^+ \)) to form \( \text{H}_3\text{O}^+ \), marking \( \text{HCl} \) as a Bronsted-Lowry acid. Water, in this situation, acts as a Bronsted-Lowry base by accepting the proton.
However, not all acids fit neatly into the Bronsted-Lowry model, like \( \text{AlCl}_3 \), which doesn't donate protons but still acts acidic because it accepts electron pairs (being a Lewis acid).
The elegance of the Bronsted-Lowry model lies in its ability to explain a wide range of acid-base reactions, although it sometimes needs support from other theories, like Lewis's, for a comprehensive view.