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2.8 Acid and Base Strength

2 min read•may 7, 2024

Acid and are crucial concepts in organic chemistry. They're determined by factors like , atom size, and . Understanding these principles helps predict reactivity and explain chemical behavior.

The scale quantifies , with lower values indicating stronger acids. There's an inverse relationship between acid and strength. These concepts are fundamental for understanding reactions and equilibria in organic systems.

Acid and Base Strength

Acid strength comparison using pKa

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  • pKa negative logarithm of acid dissociation constant () defined by equation pKa=−log(Ka)pKa = -log(Ka)
  • Lower pKa values indicate stronger acids
    • Lower pKa means higher Ka greater extent of dissociation
    • HCl (pKa = -7) stronger acid than acetic acid (pKa = 4.76)
  • Factors affecting acid strength:
    • Electronegativity of atom bonded to acidic hydrogen
      • More electronegative atoms (fluorine) stabilize increasing acid strength
    • Size of atom bonded to acidic hydrogen
      • Larger atoms (iodine) better stabilize negative charge of conjugate base
    • Resonance stabilization of conjugate base
      • Delocalization of negative charge (benzoate anion) increases acid strength
    • Inductive effects from neighboring atoms
      • Electron-withdrawing groups (nitro -NO2) increase acid strength by stabilizing conjugate base

Inverse relationship of acid-base strength

  • -base pairs related by loss or gain of proton (H+) in reaction: Acid + H2O ⇌ Conjugate base + H3O+ ()
  • Stronger acids have weaker conjugate bases
    • Strong acid (HCl) readily donates proton leaving weak conjugate base (Cl-)
  • Weaker acids have stronger conjugate bases
    • Weak acid (acetic acid CH3COOH) holds onto proton more tightly resulting in stronger conjugate base (acetate CH3COO-)
  • pKa of acid and pKb of conjugate base related by equation: pKa+pKb=14pKa + pKb = 14 (in aqueous solution at 25℃)

Calculation of water's pKa

  • of water () equilibrium constant for autoionization of water: H2O + H2O ⇌ H3O+ + OH-
    • Kw=[H3O+][OH−]=1.0×10−14Kw = [H3O+][OH-] = 1.0 × 10^-14 (at 25℃)
  • In pure water, concentrations of H3O+ and OH- are equal
    • [H3O+]=[OH−]=Kw=1.0×10−7[H3O+] = [OH-] = \sqrt{Kw} = 1.0 × 10^-7 M
  • pKa of water calculated using concentration of H3O+ in pure water
    • pKa=−log(Ka)=−log([H3O+])=−log(1.0×10−7)=7pKa = -log(Ka) = -log([H3O+]) = -log(1.0 × 10^-7) = 7
  • Therefore, pKa of water is 7 at 25℃

Acid-Base Theories and Concepts

  • Brønsted-Lowry theory: Defines acids as proton donors and bases as proton acceptors
  • : Describes acids as electron pair acceptors and bases as electron pair donors
  • : In a solvent, acids stronger than the solvent's will be leveled to the same strength
  • : A mixture of a weak acid and its conjugate base (or a weak base and its conjugate acid) that resists changes in pH when small amounts of acid or base are added

Key Terms to Review (22)

Acid Strength: Acid strength refers to the ability of an acid to donate protons (H+) in an aqueous solution. It is a measure of the extent to which an acid dissociates and releases hydrogen ions, which determines the acidity of the solution. Acid strength is a crucial factor in understanding acid-base reactions and predicting their outcomes.
Acid-Base Pair: An acid-base pair refers to a conjugate acid-base relationship, where an acid and its corresponding base form a reversible pair that can exchange protons (H+ ions) in chemical reactions. This concept is central to the Brønsted-Lowry definition of acids and bases and understanding the relative strengths of acids and bases.
Acidity constant (Ka): The acidity constant (Ka) measures the strength of an acid in solution, indicating how completely an acid dissociates into its ions. A higher Ka value signifies a stronger acid, capable of donating protons more readily.
Base Strength: Base strength refers to the ability of a base to accept protons (H+ ions) and the extent to which a base can be deprotonated. It is a measure of the base's capacity to donate electron density and its propensity to form covalent bonds with protons, which is crucial in understanding acid-base reactions and the formation of enolate ions.
Branched-chain alkane: A branched-chain alkane is an alkane that has one or more alkyl groups attached to its continuous chain of carbon atoms, creating a non-linear structure. These compounds are a type of hydrocarbon where the carbon atoms are connected by single bonds in a branching pattern, differing from straight-chain alkanes.
Buffer Solution: A buffer solution is an aqueous solution that resists changes in pH upon the addition of small amounts of an acid or base. It maintains a relatively stable pH, which is crucial in various chemical and biological processes.
Conjugate acid: A conjugate acid is formed when a base gains a proton (H+ ion) during a chemical reaction. It is the species that remains after a base has accepted a proton in the context of the Brønsted–Lowry acid-base theory.
Conjugate Acid: A conjugate acid is the species formed when a base accepts a proton (H+) in a Brønsted-Lowry acid-base reaction. It is the acid that results when a base is protonated, and it is a weaker acid than the original acid. Conjugate acids play a crucial role in understanding acid-base chemistry, the strength of acids and bases, and their behavior in various reactions, including SN1 reactions, amine basicity, and the Henderson-Hasselbalch equation for biological amines and amino acids.
Conjugate base: A conjugate base is the species that remains after an acid has donated a proton (H+ ion) during a chemical reaction. It is capable of gaining a proton in the reverse reaction, forming the original acid.
Conjugate Base: A conjugate base is the species formed when an acid loses a proton (H+) in an acid-base reaction. It is the base that is left behind when an acid donates a proton to another substance, becoming the conjugate acid-base pair. This term is central to understanding acid-base chemistry, as well as its applications in organic reactions and biological systems.
Electronegativity: Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. It is a fundamental concept in understanding the nature and strength of chemical bonds, as well as predicting the polarity and reactivity of molecules.
Electronegativity (EN): Electronegativity is a measure of an atom's ability to attract and hold onto electrons when it is part of a compound. The higher the electronegativity value, the more strongly an atom can pull electrons towards itself.
Hydronium Ion: The hydronium ion, represented as H3O+, is a cation formed when a proton (H+) is added to a water molecule. It is a key player in the understanding of acid-base chemistry and the strength of acids and bases.
Inductive effect: The inductive effect is a phenomenon observed in organic chemistry where the polarization of chemical bonds occurs due to the shifting of electrons along a chain of atoms within a molecule, caused by differences in electronegativity. This effect influences the distribution of electron density across the molecule, affecting its reactivity and properties.
Inductive Effect: The inductive effect is an electronic effect in which the unequal sharing of electrons between atoms in a molecule results in a partial charge being transmitted through the bonds of the molecule. This effect can influence the reactivity and stability of various functional groups and intermediates in organic chemistry.
Ion-Product Constant: The ion-product constant, also known as the ionic product, is a fundamental concept in acid-base chemistry that describes the equilibrium between the concentration of hydrogen ions (H+) and hydroxide ions (OH-) in a solution. It is a measure of the strength of an acid or base and is a critical factor in determining the pH of a solution.
Ka: Ka, or the acid dissociation constant, is a quantitative measure of the strength of an acid in a solution. It represents the equilibrium constant for the dissociation of an acid into its conjugate base and a hydrogen ion. The value of Ka is used to determine the pH of an acid solution and to predict the extent of acid-base reactions.
Kw: Kw, also known as the ion product constant or water autoionization constant, is a fundamental parameter in the study of acid-base chemistry. It represents the equilibrium constant for the self-ionization of water, which is a crucial concept in understanding the strength and behavior of acids and bases.
Leveling Effect: The leveling effect refers to the phenomenon where strong acids and bases in aqueous solutions are effectively neutralized and behave as if they were weak acids and bases. This concept is crucial in understanding acid-base strength and predicting acid-base reactions.
Lewis Theory: The Lewis theory, developed by Gilbert N. Lewis, provides a conceptual framework for understanding acid-base reactions and the formation of chemical bonds. It focuses on the sharing and transfer of electron pairs, rather than the traditional Arrhenius definition of acids and bases based on the presence of hydrogen ions (H+) or hydroxide ions (OH-) in aqueous solutions.
PKa: pKa, or the acid dissociation constant, is a measure of the strength of an acid in a solution. It represents the pH at which a particular acid is 50% dissociated into its conjugate base. This value is crucial in understanding the behavior and properties of acids, bases, and their reactions in organic chemistry.
Resonance Stabilization: Resonance stabilization is a phenomenon where the delocalization of electrons in a molecule or ion leads to a more stable configuration compared to a single Lewis structure. This concept is crucial in understanding the behavior and properties of various organic compounds, including their acidity, basicity, reactivity, and stability.
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2.9 Predicting Acid–Base Reactions from pKa Values