Principles of Chemical Equilibrium, Acids, Bases, and Complexes

Flashcards covering chemical equilibrium principles, Le Châtelier's principle, acid-base theories (Arrhenius, Brønsted-Lowry, Lewis), pH calculations, and complex ion formation.

Chemical Equilibrium

A state in which the rates of the forward and reverse reactions are equal and the amounts of reactants and products stop changing.

Dynamic Process

The nature of chemical equilibrium where the system is not static; reactions continue to occur in both directions at equal rates.

Reaction Quotient ($Q$)

A mathematical expression that allows us to express the amounts of reactants and products present at any point in a reversible reaction.

Law of Mass Action

The rule stating that when a reaction has achieved equilibrium at a given temperature, the reaction quotient always has the same value ($K$).

Equilibrium Constant ($K$)

A unitless value representing the ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of their coefficients.

Relationship between $K_p$ and $K_c$

Expressed by the formula $K_p = K_c(RT)^{\triangle n_g}$, where $R = 0.08206\,L\cdot atm/mol\cdot K$, $T$ is temperature in Kelvin, and $\triangle n_g$ is the change in moles of gas.

Heterogeneous Equilibrium

A system in which reactants and products are in two or more different phases; pure solids and pure liquids are excluded from the equilibrium expression.

Le Châtelier’s Principle

States that when a chemical system at equilibrium is disturbed, it returns to equilibrium by counteracting the disturbance.

Exothermic Reaction (Equilibrium Shift)

A reaction where heat is a product; increasing the temperature will shift the equilibrium away from the products (toward the reactants).

Endothermic Reaction (Equilibrium Shift)

A reaction where heat is a reactant; increasing the temperature will shift the equilibrium away from the reactants (toward the products).

Catalyst

A substance that speeds up both forward and reverse reaction rates, causing the system to reach equilibrium more quickly without affecting the equilibrium concentrations or the value of $K$.

Arrhenius Acid

A substance that produces $H^+$ ions in aqueous solution.

Arrhenius Base

A substance that produces $OH^-$ ions in aqueous solution.

Hydronium Ion ($H_3O^+$)

A complex ion formed when an $H^+$ ion (proton) reacts with a water molecule.

Brønsted-Lowry Acid

A compound that donates a proton ($H^+$) to another compound.

Brønsted-Lowry Base

A compound that accepts a proton ($H^+$) from another compound.

Conjugate Pair

Two species related by the loss or gain of a single proton ($H^+$); for every acid there is a conjugate base, and for every base there is a conjugate acid.

Amphoteric (Amphiprotic) Substances

Materials that can act as either an acid or a base because they possess both a transferable $H$ and an atom with lone pair electrons.

Ion Product Constant for Water ($K_w$)

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

pH

A logarithmic measure of acidity defined as $pH = -\log[H_3O^+]$; values less than 7 are acidic, and values greater than 7 are basic.

pOH

A logarithmic measure of basicity defined as $pOH = -\log[OH^-]$; at $25\,^\circ\text{C}$, the sum of $pH$ and $pOH$ equals $14.0$.

Strong Acid

An acid that is a strong electrolyte and ionizes practically $100\%$ in water, such as $HCl$, $HBr$, $HI$, $HClO_4$, $HNO_3$, and $H_2SO_4$.

Acid Ionization Constant ($K_a$)

The equilibrium constant for the reaction of a weak acid with water; a larger $K_a$ indicates a stronger acid.

Percent Ionization

The ratio of the concentration of ionized acid at equilibrium to the initial concentration of the acid, multiplied by $100\%$. This value decreases as the initial concentration of the acid increases.

Binary Acid Strength

Trends where acidity increases to the right across a period (due to bond polarization) and increases down a column (due to decreasing bond strength).

Oxyacid Strength

Trends where acidity increases with the electronegativity of the central atom and with the number of oxygen atoms attached to the central atom.

Lewis Acid

A species that accepts a pair of electrons to form a coordinate covalent bond.

Lewis Base

A species that donates a pair of electrons to form a coordinate covalent bond; it must possess a lone pair of electrons.

Coordinate Covalent Bond

A type of bond formed when one atom provides both of the bonding electrons; also known as a dative bond.

Complex Ion

A polyatomic ion consisting of a central metal atom surrounded by ions or molecules called ligands which are bonded via coordinate covalent bonds.

Ligands

Ions or neutral molecules that act as Lewis bases by donating electron pairs to a central metal atom in a complex ion.

Formation Constant ($K_f$)

The equilibrium constant for the formation of a complex ion directly from its components in solution.

Dissociation Constant ($K_d$)

The equilibrium constant for the decomposition of a complex ion into its components; it is the inverse of the formation constant ($K_d = \frac{1}{K_f}$).

Amphoteric Metal Hydroxides

Insoluble metal hydroxides that become more soluble in both acidic and basic solutions, such as those containing $Al^{3+}$, $Cr^{3+}$, $Zn^{2+}$, $Pb^{2+}$, and $Sb^{2+}$.