Bronsted-Lowry Acid-Base Theory and Chemical Equilibrium Principles

A comprehensive vocabulary flashcard set covering the fundamental definitions, reversible equilibria, amphoterism, conjugate pairs, equilibrium constant expressions, and acid ionization constants based on the lecture notes.

Hydrogen Ion ($H^+$)

Fundamentally a single proton in the context of acid-base chemistry.

Bronsted-Lowry Acid

A chemical species strictly defined as a proton donor.

Bronsted-Lowry Base

A chemical species strictly defined as a proton acceptor.

Equilibrium Arrow ($\rightleftharpoons$)

A bidirectional symbol signifying that a chemical reaction is fully reversible, where reactants collide to form products while products simultaneously react to regenerate reactants, and both forward and reverse rates occur concurrently at all times.

Amphoteric

A term describing a chemical species capable of functioning either as a Bronsted-Lowry acid or as a Bronsted-Lowry base depending on the reaction environment, derived from the prefix 'ampho-' meaning 'both'.

Conjugate Acid-Base Pairs

Two chemical species across a reaction that differ solely by the presence or absence of a single proton ($H^+$).

Dynamic Chemical Equilibrium

A state established in a reversible system when the forward reaction rate equals the reverse reaction rate ($\text{Rate}_{\text{forward}} = \text{Rate}_{\text{reverse}}$), rates are constant, and the concentrations of reactants and products remain invariant over time.

Equilibrium Constant ($K_{\text{eq}}$ or $K$)

An inherent, experimentally measured property unique to each chemical transformation representing the constant ratio of product concentrations to reactant concentrations at equilibrium, expressed in molarity ($M$) at a specified temperature (typically $25\,^{\circ}\text{C}$).

State of Matter Restrictions in Equilibrium Expressions

The rule dictating that pure solids ($(s)$) and pure liquids/solvents ($(l)$) are omitted from equilibrium constant expressions because their concentrations do not vary independently, leaving only dissolved solutes ($(aq)$) and gases ($(g)$).

Acid Ionization Constant ($K_a$)

A specialized equilibrium constant expression denoting the extent of acid dissociation in an aqueous solution, in which the liquid water solvent ($H_2O(l)$) is omitted from the denominator.

Equilibrium Condition: $K > 1$

A thermodynamic condition where product concentration is much greater than reactant concentration ($[\text{Products}] \gg [\text{Reactants}]$), meaning the system heavily favors forward product formation.

Equilibrium Condition: $K < 1$

A thermodynamic condition where reactant concentration is much greater than product concentration ($[\text{Reactants}] \gg [\text{Products}]$), meaning the reaction heavily favors reactants and very little forward conversion occurs.

Strong Acid

An acid characterized by an exceptionally large ionization constant ($K_a \gg 1$, such as hydroiodic acid with $K_a = 10^{10}$) that ionizes completely into hydronium ions and conjugate base ions in aqueous solution, acting as an exceptionally strong proton donor.

Weak Acid

An acid characterized by an ionization constant substantially less than $1$ ($K_a \ll 1$, such as hydrofluoric acid with $K_a = 6.6 \times 10^{-4}$) that donates its proton weakly to water and remains largely un-ionized.

Relative Acid Strength Principle

The chemical principle that the numerical magnitude of the $K_a$ value dictates acid strength across proton donors: a higher $K_a$ value indicates greater product formation and higher ionization into $H_3O^+$, directly defining a stronger acid.