Unit 6: Controlling Chemical Change - Acid-Base Equilibria and Chemical Stability

Comprehensive vocabulary flashcards covering interaction characterization, equilibrium shifts, mathematical pH models, buffer systems, and structural determinants of chemical stability across Unit 6.

Brønsted-Lowry Acid

A substance that acts as a proton ($H^+$) donor in an aqueous reaction, thereby increasing the concentration of hydronium ions ($[H_3O^+]$).

Brønsted-Lowry Base

A substance that acts as a proton ($H^+$) acceptor in an aqueous reaction, thereby increasing the concentration of hydroxide ions ($[OH^-]$).

Mechanistic Curved Arrows

Symbolic arrows in chemical structures representing the movement of electron pairs, typically starting at lone pairs or covalent bonds and terminating at an atom.

Autoionization of Water

The endothermic proton transfer process between two water molecules ($2H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)$) with $\Delta H^\circ_{\text{rxn}} = 55.9\,\text{kJ}$ and $\Delta S^\circ_{\text{rxn}} = -80.5\,\text{J}\cdot\text{K}^{-1}$.

Ion-Product Constant of Water ($K_w$)

The thermodynamic equilibrium constant for the self-ionization of water, equal to $[H_3O^+][OH^-] = 1.0 \times 10^{-14}$ at $25^\circ\text{C}$.

pH

A logarithmic measure of hydronium ion concentration defined as $-\log([H_3O^+])$, where values below $7$ represent acidic conditions and values above $7$ represent basic conditions at $25^\circ\text{C}$.

pOH

A logarithmic measure of hydroxide ion concentration defined as $-\log([OH^-])$, which satisfies $pH + pOH = 14$ in aqueous solutions at $25^\circ\text{C}$.

Acid Dissociation Constant ($K_a$)

The equilibrium constant for the reaction of an acid with water ($HA(aq) + H_2O(l) \rightleftharpoons A^-(aq) + H_3O^+(aq)$), expressed as $K_a = \frac{[A^-][H_3O^+]}{[HA]}$.

Base Ionization Constant ($K_b$)

The equilibrium constant for the reaction of a base with water ($B(aq) + H_2O(l) \rightleftharpoons HB^+(aq) + OH^-(aq)$), expressed as $K_b = \frac{[HB^+][OH^-]}{[B]}$.

Conjugate Acid-Base Pair

Two molecular species that differ only by the presence or absence of a single proton ($H^+$), linked quantitatively by the relationship $K_a \times K_b = K_w$.

$$pK_a$$

The negative base-10 logarithm of the acid dissociation constant ($-\log(K_a)$), where lower values signify stronger acids.

5% Rule for Weak Acids

A mathematical criterion validating the assumption that $[HA]_0 - x \approx [HA]_0$, which holds true when $x < 0.05[HA]_0$ or equivalently when $\frac{[HA]_0}{K_a} > 400$.

Le Chatelier's Principle

The rule stating that when an external perturbation alters a variable of a system at equilibrium, the system will shift its composition in the direction that minimizes that change.

Reaction Quotient ($Q$)

A mathematical expression identical in form to the equilibrium constant expression using non-equilibrium concentrations, where $Q < K$ drives a shift toward products and $Q > K$ drives a shift toward reactants.

Kinetic Argument for Equilibrium Shifts

An explanation of equilibrium shifts based on disturbances causing an imbalance between forward and reverse rates, shifting to products when $\text{Rate}_f > \text{Rate}_b$ and to reactants when $\text{Rate}_b > \text{Rate}_f$.

Henderson-Hasselbalch Equation

The logarithmic buffer relationship expressed as $pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right)$, relating solution pH to the relative concentrations of conjugate species.

Acid-Base Buffer

An aqueous solution containing significant, comparable concentrations of a weak acid and its conjugate base that resists substantial pH changes upon the addition of acids or bases.

Buffer Capacity Range

The effective operational pH span of a buffer, typically maintained within $pH = pK_a \pm 1$, functioning at maximum capacity when $[HA] = [A^-]$ at $pH = pK_a$.

Charge Density

The ratio of charge to ionic volume; smaller ions have higher charge density, generating stronger electrostatic attractions but imposing greater entropic ordering on surrounding water molecules.

Electronegativity Effect on Acidity

The horizontal periodic trend where higher electronegativity of the central atom energetically stabilizes the negative charge of the conjugate base, increasing parent acid strength ($CH_4 < NH_3 < H_2O < HF$).

Atomic Size Effect on Acidity

The vertical periodic trend where a larger atomic radius disperses negative charge over a larger volume, reducing charge density and solvent ordering to yield a more stable conjugate base and stronger acid ($HF < HCl < HBr < HI$).

Resonance Stabilization of Conjugate Base

The delocalization of negative charge across multiple atoms through conjugated pi systems, which lowers charge density, stabilizes the anion, and significantly increases acid strength.

Inductive Effect

The transmission of electronic charge through sigma bonds caused by electronegative electron-withdrawing groups (which stabilize conjugate bases) or electron-donating alkyl groups (which destabilize conjugate bases).

Steric Hindrance in Bases

The spatial crowding of bulky alkyl groups around a basic nitrogen atom that obstructs proton access to the lone pair, thereby decreasing base strength and increasing $pK_b$.

Lewis Acid

A chemical species that acts as an electron-pair acceptor in the formation of a coordinate covalent bond.

Lewis Base

A chemical species that acts as an electron-pair donor in the formation of a coordinate covalent bond.

Directionality of Acid-Base Reactions

The thermodynamic tendency of proton-transfer processes to proceed spontaneously in the direction that forms the weaker acid and weaker base.