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14.3 Relative Strengths of Acids and Bases

3 min read•june 25, 2024

Acid-base strength and equilibrium are crucial concepts in chemistry. They help us understand how substances interact in solution, influencing and chemical reactions. Knowing how to compare acid and base strengths allows us to predict and control these interactions.

Molecular structure plays a key role in determining acid-base strength. Factors like , resonance, and affect how easily substances donate or accept protons. This knowledge is essential for understanding chemical behavior and solving equilibrium problems.

Acid-Base Strength and Equilibrium

Acid vs base strength comparisons

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Top images from around the web for Acid vs base strength comparisons

  • Strength of Acids | Boundless Chemistry View original

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  • Relative Strengths of Acids and Bases | General Chemistry View original

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  • Strength of Acids | Boundless Chemistry View original

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  • (KaK_a) quantifies acid strength higher KaK_a indicates stronger acid (HCl vs CH3COOH)
  • (KbK_b) quantifies base strength higher KbK_b indicates stronger base (NaOH vs NH3)
  • pKapK_a and pKbpK_b scales are negative logarithms of KaK_a and KbK_b lower pKapK_a indicates stronger acid, lower pKbpK_b indicates stronger base
  • related by Ka×Kb=KwK_a \times K_b = K_w, where KwK_w is ionization constant of water (1.0×10−141.0 \times 10^{-14} at 25°C)
  • is the process by which an acid releases a proton in solution, influencing its strength

Molecular structure in acid-base strength

  • Inductive effects influence strength (halogens, nitro) stabilize conjugate base making acid stronger, (alkyl) destabilize conjugate base making acid weaker
  • influence strength stabilization of conjugate base makes acid stronger (carboxylic acids, phenols) due to charge delocalization
  • of atom bonded to acidic hydrogen affects strength sp2sp^2 or spsp hybridized atoms form stronger acids than sp3sp^3 due to greater s-character in bonding orbital
  • of atom bonded to acidic hydrogen affects strength more electronegative atoms (oxygen, nitrogen) form stronger acids than less electronegative (carbon)
  • occurs when a solvent limits the strength of an acid or base, making very strong acids or bases appear equally strong in that solvent

Weak acid-base equilibrium problems

  1. Write acid-base equilibrium expression and corresponding KaK_a or KbK_b expression
    • : HA+H2O⇌H3O++A−HA + H_2O \rightleftharpoons H_3O^+ + A^-, Ka=[H3O+][A−][HA]K_a = \frac{[H_3O^+][A^-]}{[HA]}
    • : B+H2O⇌BH++OH−B + H_2O \rightleftharpoons BH^+ + OH^-, Kb=[BH+][OH−][B]K_b = \frac{[BH^+][OH^-]}{[B]}
  2. Use initial, change, equilibrium (ICE) table to set up equilibrium concentrations
    • Initial concentrations based on given problem
    • Change in concentrations determined by reaction stoichiometry
    • Equilibrium concentrations are sum of initial and change values
  3. Substitute equilibrium concentrations into KaK_a or KbK_b expression and solve for unknown variable
    • If necessary, use approximation x<<[HA]0x << [HA]_0 or x<<[B]0x << [B]_0 to simplify calculation, where xx is change in concentration and [HA]0[HA]_0 or [B]0[B]_0 is initial concentration of acid or base
  4. Calculate pH of solution using equilibrium concentration of H3O+H_3O^+ or OH−OH^-
    • pH=−log⁡[H3O+]pH = -\log{[H_3O^+]} and [pOH](https://www.fiveableKeyTerm:pOH)=−log⁡[OH−][pOH](https://www.fiveableKeyTerm:pOH) = -\log{[OH^-]}
    • pH+pOH=14pH + pOH = 14 at 25°C

Advanced Acid-Base Concepts

  • Buffer solutions resist changes in pH when small amounts of acid or base are added
  • The relates pH to the concentrations of a weak acid and its conjugate base in a
  • occurs when water molecules react with each other to produce hydronium and hydroxide ions, maintaining a constant ion product in pure water

Key Terms to Review (37)

A-: The term 'A-' refers to the relative strength of an acid in the context of acid-base chemistry. It is a measure of the extent to which an acid dissociates or ionizes in water, with a higher 'A-' value indicating a stronger acid.
Acid Dissociation: Acid dissociation is the process by which an acid releases hydrogen ions (H+) when dissolved in water, forming a solution of hydrogen ions and the conjugate base of the acid. This process is crucial in understanding the relative strengths of acids and bases and their behavior in aqueous solutions.
Acid Ionization Constant: The acid ionization constant, denoted as $K_a$, 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 hydrogen ions, providing a numerical value that indicates the extent to which an acid will ionize in water.
Alpha (α) decay: Alpha (α) decay is a type of radioactive decay where an unstable nucleus emits an alpha particle, consisting of 2 protons and 2 neutrons. This process reduces the atomic number by 2 and the mass number by 4.
Autoionization of Water: Autoionization of water is the spontaneous dissociation of water molecules into hydrogen ions (H+) and hydroxide ions (OH-) within the water itself, without the presence of any other substances. This process is crucial in understanding the relative strengths of acids and bases.
B: B is a fundamental concept that spans multiple topics in chemistry, including atomic structure, acid-base chemistry, and the relative strengths of acids and bases. It is a versatile and essential term that underpins our understanding of various chemical phenomena.
Base Ionization Constant: The base ionization constant, denoted as $K_b$, is a measure of the strength of a base in aqueous solutions. It quantifies the extent to which a base dissociates and releases hydroxide ions (OH^-) when dissolved in water, which is a key factor in determining the relative strengths of acids and bases.
Base-ionization constant (Kb): The base-ionization constant (Kb) quantifies the strength of a base in a solution. It is the equilibrium constant for the dissociation of a base into its conjugate acid and hydroxide ion.
BH+: BH+ is a positively charged species that represents a protonated base, formed when a base (B) accepts a proton (H+) in an acid-base reaction. This term is particularly relevant in the context of understanding the relative strengths of acids and bases, as it provides insight into the equilibrium dynamics between conjugate acid-base pairs.
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 is a mixture of a weak acid and its conjugate base, or a weak base and its conjugate acid, which helps maintain a relatively stable pH in a chemical system.
Conjugate Acid-Base Pairs: Conjugate acid-base pairs are related chemical species that differ by the presence or absence of a single proton (H+). When an acid donates a proton, it becomes a conjugate base, and when a base accepts a proton, it becomes a conjugate acid. These pairs are fundamental to understanding the Brønsted-Lowry theory of acids and bases, as well as the concepts of pH, relative acid-base strengths, hydrolysis, polyprotic acids, and acid-base titrations.
Electron-Donating Groups: Electron-donating groups are substituents or functional groups in organic molecules that have the ability to donate or contribute electrons, increasing the electron density around a particular atom or region of the molecule. These groups play a crucial role in determining the relative strengths of acids and bases.
Electron-Withdrawing Groups: Electron-withdrawing groups are functional groups or atoms within a molecule that have a strong tendency to attract and withdraw electrons from the surrounding atoms. This property can significantly influence the reactivity and stability of the molecule.
Electronegativity: Electronegativity is a measure of an atom's ability to attract and hold onto electrons within a chemical bond. It is a dimensionless quantity usually assigned values on the Pauling scale.
Electronegativity: Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. It is a fundamental property that influences the nature and strength of chemical bonds, as well as the physical and chemical properties of substances.
H3O+: H3O+ is the hydronium ion, a positively charged species formed when a proton (H+) interacts with a water molecule (H2O). It is a crucial concept in understanding the behavior of acids, bases, and the pH of aqueous solutions.
HA: HA, or Brønsted-Lowry acid, is a chemical species that can donate a proton (H+) to another substance, thereby acting as an acid in a chemical reaction. This term is particularly relevant in the context of understanding Brønsted-Lowry acid-base theory and the relative strengths of acids and bases.
Henderson-Hasselbalch equation: The Henderson-Hasselbalch equation is a mathematical relationship that describes the pH of a solution containing a weak acid or a weak base. It is a fundamental concept in understanding the relative strengths of acids and bases, as well as the behavior of polyprotic acids and buffers.
Hybridization: Hybridization is the concept in chemistry where atomic orbitals combine to form new hybrid orbitals that are suitable for the pairing of electrons to form chemical bonds. This idea helps explain molecular geometry and bonding properties, linking the arrangement of atoms in a molecule to their electron configurations and the types of bonds formed.
ICE Table: An ICE table, also known as an Initial, Change, and Equilibrium table, is a tool used in chemistry to organize and analyze the concentrations of reactants and products in a chemical equilibrium system. It provides a structured way to visualize and calculate the changes in concentrations as a reaction approaches equilibrium.
Inductive Effects: Inductive effects refer to the ability of atoms or functional groups to influence the electron distribution within a molecule through the transmission of electronic effects along the carbon chain or molecular framework. This phenomenon plays a crucial role in understanding the relative strengths of acids and bases.
Ion-product constant for water, Kw: The ion-product constant for water, $K_w$, is the equilibrium constant for the self-ionization of water. It is defined as the product of the molar concentrations of hydrogen ions and hydroxide ions in water at a given temperature.
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 constituent ions, providing insight into the extent to which an acid ionizes in water.
Kb: Kb, or the base dissociation constant, is a measure of the strength of a base in aqueous solution. It quantifies the extent to which a base dissociates or ionizes in water, providing insight into the relative strengths of different bases.
Kw: Kw, or the equilibrium constant for water, is a fundamental concept in chemistry that describes the self-ionization of water and its relationship to the acidity or basicity of a solution. This term is crucial in understanding Brønsted-Lowry acid-base theory, pH and pOH calculations, as well as the relative strengths of acids and bases.
Leveling Effect: The leveling effect refers to the phenomenon where acids and bases of vastly different strengths are able to ionize to the same extent in aqueous solutions. This occurs due to the high concentration of water molecules, which can act as both an acid and a base, effectively 'leveling' the strengths of the solutes.
OH-: OH- is the hydroxide ion, a negatively charged particle composed of one oxygen atom and one hydrogen atom. This ion is a key component in understanding acid-base chemistry, as it plays a central role in the concepts of pH, pOH, relative strengths of acids and bases, hydrolysis of salts, buffers, acid-base titrations, Lewis acids and bases, and coupled equilibria.
Oxyacids: Oxyacids are acids that contain hydrogen, oxygen, and another element (the central atom). They dissociate in water to produce H+ ions and the corresponding oxyanion.
Percent ionization: Percent ionization is the percentage of acid or base molecules that dissociate into ions in a solution. It indicates the strength of an acid or base in a given concentration.
PH: pH, or potential of hydrogen, is a measure of the acidity or basicity of a solution. It is a scale that ranges from 0 to 14, with 7 being neutral, values less than 7 being acidic, and values greater than 7 being basic or alkaline. The pH of a solution is directly related to the concentration of hydrogen ions (H+) present, and it is a critical factor in many chemical and biological processes.
PKa: pKa is a measure of the strength of an acid, representing the pH at which a weak acid is 50% dissociated. It is a critical parameter that helps determine the relative strengths of acids and bases, the behavior of polyprotic acids, the effectiveness of buffers, and the progress of acid-base titrations.
PKb: pKb is the negative logarithm of the equilibrium constant (Kb) for a base in an aqueous solution. It is a measure of the strength of a base, providing information about its ability to accept protons and form hydroxide ions in water.
POH: pOH, or the negative logarithm of the hydroxide ion concentration, is a measure of the acidity or basicity of a solution. It is a crucial concept in understanding the relationship between pH and the relative strengths of acids and bases, as well as its application in acid-base titrations.
Resonance Effects: Resonance effects refer to the stabilization or destabilization of molecules or ions due to the delocalization of electrons across multiple atoms or functional groups. This phenomenon plays a crucial role in understanding the relative strengths of acids and bases, as it can influence the stability and reactivity of these species.
Weak Acid: A weak acid is a type of acid that only partially dissociates into its constituent ions when dissolved in water. This means that at equilibrium, a significant portion of the acid molecules remain undissociated, unlike strong acids which completely dissociate.
Weak Base: A weak base is a type of base that only partially dissociates in an aqueous solution, producing a relatively low concentration of hydroxide ions. This is in contrast to strong bases, which completely dissociate in water, producing a high concentration of hydroxide ions.
Weak bases: Weak bases are substances that partially ionize in water, producing a relatively small number of hydroxide ions ($OH^-$). They have a higher $pH$ compared to strong acids but do not completely dissociate in solution.
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Glossary
14.4 Hydrolysis of Salts