8.6: Molecular Structures of Acids and Bases
Introduction
This section explores the relationship between molecular structure and the acidic or basic properties of molecules. Understanding the structure of acids and bases helps to predict their behavior in chemical reactions. Acid Strength and Conjugate Base Strength and Base Strength and Conjugate Acid Strength are crucial factors.
Acids
Binary Acids
Binary acids are composed of hydrogen and one other element. The acidity of binary acids increases down a group and across a period.
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Across a period: Acidity increases with increasing electronegativity of the nonmetal. This is because the bond becomes more polarized, making it easier to release $ H^+ $ . Example: $ HF < HCl < HBr < HI $ .
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Down a group: Acidity increases with increasing size of the nonmetal. The bond strength decreases, making it easier to release $ H^+ $ .
Oxyacids
Oxyacids contain hydrogen, oxygen, and another element (usually a nonmetal). The acidity of oxyacids depends on:
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Electronegativity of the central atom: Acidity increases with increasing electronegativity of the central atom. This is because the central atom pulls electron density away from the $ O-H $ bond, making it easier to release $ H^+ $ .
$$ HClO < HBrO < HIO $$
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Number of oxygen atoms: Acidity increases with the number of oxygen atoms bonded to the central atom. Each additional oxygen atom pulls more electron density away from the $ O-H $ bond, making it easier to release $ H^+ $ .
$$ HClO < HClO_2 < HClO_3 < HClO_4 $$
Carboxylic Acids
Carboxylic acids contain the $ -COOH $ group. The acidity of carboxylic acids is influenced by:
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Inductive effect: Electron-withdrawing groups near the carboxyl group increase acidity by stabilizing the conjugate base (carboxylate ion) through dispersal of the negative charge.
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Resonance stabilization: The carboxylate ion ( $ RCOO^− $ ) is stabilized by resonance, delocalizing the negative charge over two oxygen atoms. Resonance Structures
Bases
Molecular Structure and Basicity
The basicity of a molecule depends on its ability to accept a proton ( $ H^+ $ ). Key factors are:
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Availability of lone pair electrons: A molecule with a lone pair of electrons is more likely to act as a base. Lewis structures
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Electronegativity: Less electronegative atoms are better bases because they hold onto their electrons less tightly, making them more available for bonding to a proton.
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Steric hindrance: Bulky groups around the basic site can hinder protonation, decreasing basicity.
Amines
Amines are derivatives of ammonia ( $ NH_3 $ ) where one or more hydrogen atoms are replaced by organic groups (R).
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Aliphatic amines: Amines with alkyl groups (R) are generally stronger bases than ammonia because alkyl groups are electron-donating and increase the electron density on the nitrogen atom.
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Aromatic amines: Aromatic amines (e.g., aniline) are generally weaker bases than ammonia because the lone pair of electrons on the nitrogen atom is delocalized into the aromatic ring, reducing its availability for protonation.
Relationship Between Structure and Acid/Base Strength
The strength of an acid is inversely related to the strength of its conjugate base, and vice versa. A strong acid has a weak conjugate base, and a weak acid has a strong conjugate base. The reverse is also true: Strong bases will have weak conjugate acids, while weak bases will have strong conjugate acids. Conjugate Acid-Base Pairs
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Stabilizing the conjugate base: Any structural feature that stabilizes the conjugate base of an acid will increase the acidity of the acid. For example, resonance, inductive effects, and the size of the anion all stabilize the conjugate base and increase acidity.
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Stabilizing the conjugate acid: Any structural feature that stabilizes the conjugate acid of a base will increase the basicity of the base. For example, inductive effects can stabilize the positive charge on the conjugate acid.
Polyprotic Acids
Polyprotic acids have more than one ionizable proton. The acid dissociation constants ( $ K_a $ ) for each proton are different, and they decrease with each successive ionization. Polyprotic Acids
$$ H_2A \rightleftharpoons HA^- + H^+ \quad K_{a1} $$ $$ HA^- \rightleftharpoons A^{2-} + H^+ \quad K_{a2} $$
Generally, $ K_{a1} > K_{a2} $ , because it is more difficult to remove a proton from a negatively charged ion.
Amphoteric Substances
Amphoteric substances can act as both acids and bases, depending on the reaction conditions. Amphoteric Substances
Example: Comparing Acid Strengths
Consider the following acids: $ HClO_4 $ , $ HClO_3 $ , $ HClO_2 $ , and $ HClO $ . Arrange them in order of increasing acid strength.
Solution: The acidity increases with the number of oxygen atoms. Therefore, the order of increasing acid strength is:
$$ HClO < HClO_2 < HClO_3 < HClO_4 $$