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Chapter 16: Problem 92

(a) How does the strength of an acid vary with the polarity and strength of the \(\mathrm{H}-\mathrm{X}\) bond? (b) How does the acidity of the binary acid of an element vary as a function of the electronegativity of the element? How does this relate to the position of the element in the periodic table?

Short Answer

Expert verified
(a) The strength of an acid increases with increasing bond polarity and decreasing bond strength of the \(\mathrm{H}-\mathrm{X}\) bond. (b) The acidity of a binary acid typically increases with increasing electronegativity of the element X. In the periodic table, electronegativity increases from left to right across a period and from bottom to top within a group, which corresponds to the increasing acidity of binary acids in these directions. For example, in Group 16, the acidity increases as H2O < H2S < H2Se < H2Te, and in Group 17, the acidity increases as HF < HCl < HBr < HI.

Step by step solution

01

(a) Strength of an Acid and \(\mathrm{H}-\mathrm{X}\) Bond Polarity and Strength

To determine the strength of an acid (typically represented as \(\mathrm{HX}\)), we will consider the polarity and strength of the \(\mathrm{H}-\mathrm{X}\) bond. The strength of an acid is determined by its ability to donate a proton (\(\mathrm{H}^+\)), or in other words, the extent to which it ionizes in a given solution: \[\mathrm{HX} \rightleftharpoons \mathrm{H}^+ + \mathrm{X}^-\] The more polar the \(\mathrm{H}-\mathrm{X}\) bond, the greater the difference in electronegativity between the two bonded atoms. This makes it easier for the bond to break and form ions. Additionally, a weak \(\mathrm{H}-\mathrm{X}\) bond will break more readily, which will also increase the acidity of the binary acid. As a result, we can conclude that the strength of an acid increases with increasing bond polarity and decreasing bond strength.
02

(b) Acidity and Electronegativity of Binary Acids

The acidity of a binary acid typically increases with increasing electronegativity of the element X. This is because, as the electronegativity of X increases, the polarity of the \(\mathrm{H}-\mathrm{X}\) bond increases, making it easier for the bond to break and form ions.
03

(b) Acidity, Electronegativity, and the Periodic Table

In the periodic table, electronegativity increases from left to right across a period and from bottom to top within a group. Therefore, the acidity of binary acids tends to increase from left to right across a period and from bottom to top within a group. For example, consider the binary acids of elements in groups 16 and 17: Group 16: H2O < H2S < H2Se < H2Te (least to most acidic) Group 17: HF < HCl < HBr < HI (least to most acidic) In both groups, the acidity of the binary acid increases with the increase in electronegativity, representing the general trend seen in the periodic table.

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Most popular questions from this chapter

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

The amino acid glycine \(\left(\mathrm{H}_{2} \mathrm{~N}-\mathrm{CH}_{2}-\mathrm{COOH}\right)\) can participate in the following equilibria in water: \(\mathrm{H}_{2} \mathrm{~N}-\mathrm{CH}_{2}-\mathrm{COOH}+\mathrm{H}_{2} \mathrm{O}=\) \(\mathrm{H}_{2} \mathrm{~N}-\mathrm{CH}_{2}-\mathrm{COO}^{-}+\mathrm{H}_{3} \mathrm{O}^{+} \quad K_{a}=4.3 \times 10^{-3}\) \(\mathrm{H}_{2} \mathrm{~N}-\mathrm{CH}_{2}-\mathrm{COOH}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons\) \({ }^{+} \mathrm{H}_{3} \mathrm{~N}-\mathrm{CH}_{2}-\mathrm{COOH}+\mathrm{OH}^{-} \quad K_{b}=6.0 \times 10^{-5}\) (a) Use the values of \(K_{a}\) and \(K_{b}\) to estimate the equilibrium constant for the intramolecular proton transfer to form a zwitterion: \(\mathrm{H}_{2} \mathrm{~N}-\mathrm{CH}_{2}-\mathrm{COOH} \rightleftharpoons{ }^{+} \mathrm{H}_{3} \mathrm{~N}-\mathrm{CH}_{2}-\mathrm{COO}^{-}\) What assumptions did you need to make? (b) What is the pH of a \(0.050 \mathrm{M}\) aqueous solution of glycine? (c) What would be the predominant form of glycine in a solution with pH 13? With pH 1?

\begin{aligned} &\text { In your own words, define or explain (a) } K_{w} \text { , (b) } K_{a r}\\\ &\text { (d) } \mathrm{pK}_{\mathrm{b}}\\\ &\text { (c) } \mathrm{pOH} \text { , } \end{aligned}

(a) Write an equation for the reaction in which \(\mathrm{H}_{2} \mathrm{C}_{6} \mathrm{H}_{7} \mathrm{O}_{5}^{-}(a q)\) acts as a base in \(\mathrm{H}_{2} \mathrm{O}(l) .\) (b) Write an equation for the reaction in which \(\mathrm{H}_{2} \mathrm{C}_{6} \mathrm{H}_{7} \mathrm{O}_{5}^{-}(\mathrm{aq})\) acts as an acid in \(\mathrm{H}_{2} \mathrm{O}(l)\). (c) What is the conjugate acid of \(\mathrm{H}_{2} \mathrm{C}_{6} \mathrm{H}_{7} \mathrm{O}_{5}^{7} ?\) What is its conjugate base?

(a) The hydrogen oxalate ion \(\left(\mathrm{HC}_{2} \mathrm{O}_{4}{ }^{-}\right)\) is amphiprotic. Write a balanced chemical equation showing how it acts as an acid toward water and another equation showing how it acts as a base toward water. (b) What is the conjugate acid of \(\mathrm{HC}_{2} \mathrm{O}_{4}^{-} ?\) What is its conjugate base?
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