Problem 22 - A Complete Resource Book in Chemistry for JEE Main

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 \(\mathrm{AlCl}_{3}\).

Step-by-Step Solution

Answer

Statement (a) is not correct; Cl⁻ is not a Lewis acid.

Step 1: Understand the Concepts

We must understand that a Lewis acid is a species that can accept an electron pair. A Bronsted-Lowry acid is a species capable of donating a proton (H⁺). AlCl₃ is known to be a Lewis acid but not a Bronsted-Lowry acid. The pH scale is used to determine acidity or alkalinity of substances, where a pH less than 7 indicates acidity. The ionic product of water is the product of the concentrations of hydrogen ions and hydroxide ions in water, which is 10⁻¹⁴ mol² L⁻² at 25°C.

Step 2: Analyze Statement (a)

The statement (a) claims that Cl⁻ is a Lewis acid. Recall that a Lewis acid accepts an electron pair. Cl⁻, being negatively charged and already having a full electron shell, does not accept more electrons. Therefore, Cl⁻ is not a Lewis acid.

Step 3: Analyze Statement (b)

Statement (b) indicates that the pH of a 10⁻⁸ M HCl solution is less than 7. When calculating the pH of such a dilute solution, we must consider the H⁺ contribution from water itself. Normally, the pH of a neutral solution at 25°C is 7. However, with additional H⁺ from HCl, the pH will indeed be less than 7, as a minor increase in H⁺ concentration from HCl does lower the pH slightly below 7.

Step 4: Analyze Statement (c)

Statement (c) mentions the ionic product of water at 25°C being 10⁻¹⁴ mol² L⁻². This is correct by definition and represents the product [H⁺][OH⁻] in pure water at 25°C.

Step 5: Analyze Statement (d)

Statement (d) asserts that the Bronsted-Lowry theory could not explain the acidic character of AlCl₃. Since Bronsted-Lowry acids are described as proton donors, and AlCl₃ does not donate a proton, it is commonly known that Bronsted-Lowry theory doesn't explain AlCl₃'s acidity, rather Lewis theory does, as it acts as an electron-pair acceptor.

Step 6: Conclusion

After analyzing all statements, it is clear that statement (a) is not correct because Cl⁻ is not a Lewis acid.

Key Concepts

Lewis Acid

In the realm of chemistry, a Lewis acid is a fascinating concept. It's a substance that can accept a pair of electrons. Unlike the traditional view that focuses solely on protons, this is about electrons. Consider it like a vacuum for electrons.
This concept, named after the chemist Gilbert Lewis, has broadened our understanding of chemical reactions. For instance, when AlCl₃ acts as a Lewis acid, it accepts electrons, forming new bonds.

• Lewis acids like AlCl₃ do not have to involve a proton (H⁺).
• They focus on electron pair acceptance.
• These acids can be positive ions or neutral molecules with empty orbitals.

When you encounter substances in chemical equations, remembering that a Lewis acid is about electron acceptance rather than proton donation can simplify complex reactions. It highlights the diverse nature of acids in chemistry.

Bronsted-Lowry Acid

The Bronsted-Lowry theory refines our comprehension of acids and bases. Here, the focus is on protons. A Bronsted-Lowry acid is defined as a substance capable of donating a proton (H⁺ ion). In essence, these acids are like buckets shedding their extra protons.
In aqueous solutions, common examples include substances like HCl, which release H⁺ ions in water, a direct pathway to acidity. However, not all acidic substances fit into this framework. AlCl₃'s acidity, for instance, is not due to proton donation, showing the limitations of Bronsted-Lowry's exclusively proton-focusing definition.

• Proton donation is the key identifier in Bronsted-Lowry acids.
• It contrasts with Lewis acids that involve electron pairs.
• This makes the concept more relatable with everyday acids like many organic acids.

Understanding Bronsted-Lowry acids helps clarify why certain reactions occur based on proton exchange, providing a more intuitive insight into reaction mechanics and pH shifts in solutions.

pH of Solutions

pH is a crucial concept when discussing the acidity or basicity of a solution. It measures the concentration of hydrogen ions, with the formula \[ \text{pH} = -\log_{10} [\text{H}^+] \]. A low pH (less than 7) indicates acidity, while a pH greater than 7 suggests basicity.
A fascinating scenario arises with very dilute HCl solutions, such as a solution with a concentration of \(10^{-8} \) M. Normally, pure water has a pH of 7, due to equal concentrations of H⁺ and OH⁻ ions. However, introducing any acid like HCl reduces the pH slightly, making it less than 7, due to the increased H⁺ concentration.

• pH less than 7 means the solution is acidic.
• Neutral water contributes to overall hydrogen ion concentration in solutions.
• Even very dilute acid solutions can exhibit acidic pH.

Recognizing how pH reflects the acidity or basicity allows us to evaluate the chemical environment of a solution, predicting how various reactions might proceed.

Ionic Product of Water

At room temperature, the ionic product of water is an essential concept in understanding chemical equilibriums. It is described by the product \[ [\text{H}^+][\text{OH}^-] = 10^{-14} \text{ mol}^2 \text{ L}^{-2} \] at 25°C. This equilibrium information is critical because it indicates the balance between hydrogen ions and hydroxide ions in pure water.
By knowing this product, we grasp that in a neutral solution, [H⁺] and [OH⁻] are both equal to \(10^{-7}\). Any deviation from this balance due to added acids or bases shifts the pH accordingly.

• This ionic product reveals the self-ionization of water.
• pH and pOH are directly derived from these ion concentrations.
• Changes in temperature influence this equilibrium constant.

The ionic product of water is pivotal in understanding why the pH of neutral water is 7 and aids in predicting outcomes when acids or bases are introduced into a system. It underscores the delicate balance in aqueous chemistry.