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Chapter 16: Problem 24

(a) Which of the following is the stronger Br酶nsted-Lowry acid, \(\mathrm{HClO}_{3}\) or \(\mathrm{HClO}_{2}\) ? (b) Which is the stronger BronstedLowry base, \(\mathrm{HS}^{-}\)or \(\mathrm{HSO}_{4}^{-}\)?

Short Answer

Expert verified
(a) The stronger Br酶nsted-Lowry acid is \(\mathrm{HClO}_{3}\), as its conjugate base \(\mathrm{ClO}_{3}^{-}\) is more stable due to a higher number of oxygen atoms. (b) The stronger Br酶nsted-Lowry base is \(\mathrm{HSO}_{4}^{-}\), as its conjugate acid \(\mathrm{H}_{2}\mathrm{SO}_{4}\) is more stable due to more electronegative oxygen atoms.

Step by step solution

01

Compare the acidic strengths of \(\mathrm{HClO}_{3}\) and \(\mathrm{HClO}_{2}\)

In order to compare their acidic strengths, let's analyze the conjugate bases of both acids after losing a proton: - Conjugate base of \(\mathrm{HClO}_{3}\): \(\mathrm{ClO}_{3}^{-}\) - Conjugate base of \(\mathrm{HClO}_{2}\): \(\mathrm{ClO}_{2}^{-}\) The stability of the conjugate base is determined by its ability to distribute the negative charge among the more electronegative atoms. Electronegativity of oxygen is high, so the higher the number of oxygen atoms attached to the central chlorine atom, the more stable the conjugate base will be due to electron delocalization. In this case, \(\mathrm{ClO}_{3}^{-}\) has more oxygen atoms than \(\mathrm{ClO}_{2}^{-}\). Therefore, the conjugate base of \(\mathrm{HClO}_{3}\) is more stable than the conjugate base of \(\mathrm{HClO}_{2}\). Since a more stable conjugate base corresponds to a stronger acid, we can conclude that \(\mathrm{HClO}_{3}\) is the stronger Br酶nsted-Lowry acid.
02

Compare the basic strengths of \(\mathrm{HS}^{-}\) and \(\mathrm{HSO}_{4}^{-}\)

In order to compare their basic strengths, let's analyze the conjugate acids of both bases after gaining a proton: - Conjugate acid of \(\mathrm{HS}^{-}\): \(\mathrm{H}_{2}\mathrm{S}\) - Conjugate acid of \(\mathrm{HSO}_{4}^{-}\): \(\mathrm{H}_{2}\mathrm{SO}_{4}\) The stability of the conjugate acid is determined by its ability to hold the extra proton while distributing the positive charge. In this case, \(\mathrm{H}_{2}\mathrm{S}\) has only one oxygen atom, while \(\mathrm{H}_{2}\mathrm{SO}_{4}\) has four oxygen atoms. Since oxygen is more electronegative than sulfur, the positive charge is more effectively distributed among the highly electronegative oxygen atoms in \(\mathrm{H}_{2}\mathrm{SO}_{4}\), making it a more stable conjugate acid. Since a more stable conjugate acid corresponds to a stronger base, we can conclude that \(\mathrm{HSO}_{4}^{-}\) is the stronger Br酶nsted-Lowry base. So, the final answers are: (a) The stronger Br酶nsted-Lowry acid is \(\mathrm{HClO}_{3}\). (b) The stronger Br酶nsted-Lowry base is \(\mathrm{HSO}_{4}^{-}\).

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Understanding Acidic Strength
The acidic strength of a compound is an essential aspect of chemistry, especially in understanding how substances behave in various reactions. Acidic strength refers to an acid's ability to donate a proton (\(\text{H}^+\)), and the stronger the acid, the more readily it will lose this proton.
One way to determine acidic strength is to consider the stability of its conjugate base after the acid has donated a proton. If the conjugate base is stable, the acid tends to be strong. This is because stable conjugate bases are better at handling the negative charge resulting from the loss of a proton.
In our exercise, we see that \(\text{HClO}_3\) is stronger compared to \(\text{HClO}_2\), primarily due to the increased stability of its conjugate base.
The Role of Conjugate Bases
Conjugate bases are formed when an acid donates a proton. They play a crucial role in determining the acid's strength because a stable conjugate base indicates a strong acid. The stability of these bases can be attributed to several factors including:
  • Electron delocalization: More electronegative atoms can disperse the negative charge.
  • Structure: The presence of electronegative atoms, like oxygen, supports the spread of negative charges, making the conjugate base more stable.
The conjugate base of \(\text{HClO}_3\), \(\text{ClO}_3^-\), benefits from having more oxygen atoms compared to \(\text{ClO}_2^-\), the conjugate base of \(\text{HClO}_2\). This results in better charge distribution and increased stability of \(\text{ClO}_3^-\), thereby making \(\text{HClO}_3\) a stronger acid.
Electronegativity and Its Impact
Electronegativity is the tendency of an atom to attract shared electrons in a bond. This property significantly influences both acidic strength and the stability of conjugate bases and acids.
In acids and their conjugate bases, highly electronegative atoms like oxygen can stabilize negative charge through electron delocalization. A greater number of such atoms can lead to increased stability, as seen in the case of \(\text{HClO}_3\). More electronegative elements result in a more stable conjugate base, rendering the corresponding acid stronger.
  • Electronegativity is crucial for distributing charges effectively.
  • The greater the electronegativity difference in a molecule, the more stable the resulting ion will be.
Understanding this concept is key to comprehending why certain acids behave differently in chemical reactions, as observed in the comparison between \(\text{HClO}_3\) and \(\text{HClO}_2\).
Conjugate Acid and Base Pairing
In any Br酶nsted-Lowry reaction, the acid and base each have a corresponding conjugate base and conjugate acid. This pairing is vital for understanding reactions. When a base gains a proton, it forms a conjugate acid, and similarly, when an acid loses a proton, it forms a conjugate base.
For example, in our exercise, \(\text{HS}^-\) becomes \(\text{H}_2\text{S}\) (its conjugate acid), and \(\text{HSO}_4^-\) transforms into \(\text{H}_2\text{SO}_4\), a much more stable acid due to oxygen's high electronegativity. The stability of the conjugate acid, \(\text{H}_2\text{SO}_4\), shows why \(\text{HSO}_4^-\) is a stronger base.
This interrelation is critical for predicting behavior in acid-base reactions, harnessing the concept of conjugate acid-base pairs as a framework for understanding chemical interactions.

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Most popular questions from this chapter

In many reactions the addition of \(\mathrm{AlCl}_{3}\) produces the same effect as the addition of \(\mathrm{H}^{+}\). (a) Draw a Lewis structure for \(\mathrm{AlCl}_{3}\) in which no atoms carry formal charges, and determine its structure using the VSEPR method. (b) What characteristic is notable about the structure in part (a) that helps us understand the acidic character of \(\mathrm{AlCl}_{3}\) ? (c) Predict the result of the reaction between \(\mathrm{AlCl}_{3}\) and \(\mathrm{NH}_{3}\) in a solvent that does not participate as a reactant. (d) Which acid-base theory is most suitable for discussing the similarities between \(\mathrm{AlCl}_{3}\) and \(\mathrm{H}^{+}\)?

An unknown salt is either \(\mathrm{NaF}, \mathrm{NaCl}\), or \(\mathrm{NaOCl}\). When \(0.050 \mathrm{~mol}\) of the salt is dissolved in water to form \(0.500 \mathrm{~L}\) of solution, the pH of the solution is \(8.08\). What is the identity of the salt? \(0.100 \mathrm{M}\) solution of the salt is neutral, what is the identity of the salt?

Indicate whether each of the following statements is true or false. For each statement that is false, correct the statement to make it true. (a) Acid strength in a series of \(\mathrm{H}\) - A molecules increases with increasing size of \(A\). (b) For acids of the same general structure but differing electronegativities of the central atoms, acid strength decreases with increasing electronegativity of the central atom. (c) The strongest acid known is \(\mathrm{HF}\) because fluorine is the most electronegative element. Lewis Acids and Bases (Section 16.11)

Predict whether aqueous solutions of the following compounds are acidic, basic, or neutral: (a) \(\mathrm{NH}_{4} \mathrm{Br}\), (b) \(\mathrm{FeCl}_{3 \text {, }}\) (c) \(\mathrm{Na}_{2} \mathrm{CO}_{3}\), (d) \(\mathrm{KClO}_{4}\), (e) \(\mathrm{NaHC}_{2} \mathrm{O}_{4}\).

Write the chemical equation and the \(K_{a}\) expression for the ionization of each of the following acids in aqueous solution. First show the reaction with \(\mathrm{H}^{+}(a q)\) as a product and then with the hydronium ion: (a) \(\mathrm{HBrO}_{2}\) (b) \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{COOH}\).
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