10 Equilibria in Aqueous Solution — Vocabulary Flashcards

Vocabulary flashcards covering key terms and definitions from the lecture notes on chemical equilibria, Le Châtelier’s principle, solubility, and related concepts.

Chemical equilibrium

The state of a reversible reaction in which the forward and reverse rates are equal and the concentrations of all species remain constant with time.

Dynamic equilibrium

Equilibrium in which the forward and reverse reactions occur at the same rate, so no net change in concentrations.

Reversible reaction

A reaction that can proceed in both forward and reverse directions and form an equilibrium.

Le Châtelier’s Principle

If a system at equilibrium is disturbed by changes in concentration, temperature, or pressure, the system shifts to counteract the change and restore equilibrium.

Equilibrium constant (Kc)

The ratio [C]^c [D]^d / ([A]^a [B]^b) for aA + bB ⇌ cC + dD, based on molar concentrations.

Kp

The equilibrium constant based on partial pressures for gaseous equilibria.

Δn

Change in moles of gas; used in Kp–Kc relation: Δn = (c+d) − (a+b).

Kp = Kc (RT)^Δn

Relation between Kp and Kc for gas-phase equilibria, where R is the gas constant and T is temperature.

Haber Process

Industrial synthesis of ammonia: N2(g) + 3 H2(g) ⇌ 2 NH3(g); optimized at high pressure and moderate temperature.

Homogeneous equilibrium

All reactants and products are in the same phase (gas or aqueous).

Heterogeneous equilibrium

Reactants and products are in different phases; pure solids and liquids are excluded from the expression.

Acid–base equilibrium

Equilibria involving transfer of protons (H+) in water, including conjugate acid–base pairs and Ka/Kb relationships.

Bronsted–Lowry acid

A substance that donates a proton (H+).

Bronsted–Lowry base

A substance that accepts a proton (H+).

Ka

Acid dissociation constant; Ka = [H+][A−]/[HA].

Kb

Base dissociation constant; Kb = [BH+][OH−]/[B].

Kw

Ion-product of water; Kw = [H+][OH−] = 1.0 × 10−14 at 25°C.

pH

The negative logarithm of the hydrogen ion concentration: pH = −log[H+].

pOH

The negative logarithm of the hydroxide ion concentration: pOH = −log[OH−].

pH + pOH = 14

At 25°C, the sum of pH and pOH equals 14.

Ksp (solubility product)

The product of ion concentrations at equilibrium for a sparingly soluble salt, e.g., Ksp = [Ba2+][SO4^2−] for BaSO4.

'Common Ion Effect'

Decrease in solubility of a salt when a common ion is added to the solution.

Solubility from Ksp

Using Ksp to calculate the molar solubility of a salt in water.

Complex ion formation

Formation of a complex ion from a metal ion and ligands, often increasing solubility.

Stability constant Kf

Equilibrium constant for the formation of a complex ion; indicates the stability of the complex.

Complex ion and solubility

Formation of complex ions can increase the solubility of sparingly soluble salts.

Reaction quotient Q

A ratio calculated like K using current concentrations; used to predict the direction to reach equilibrium.

Q vs K interpretation

If Q < K, reaction proceeds toward products; if Q > K, toward reactants; Q = K means equilibrium.

Evaporation equilibrium (water)

For H2O(l) ⇌ H2O(g), the equilibrium constant is based on the partial pressure of water vapor (Kp = P(H2O)).

PbCl2 dissolution

PbCl2(s) ⇌ Pb2+(aq) + 2 Cl−(aq); Ksp = [Pb2+][Cl−]^2.

CaCO3(s) ⇌ CaO(s) + CO2(g)

Solubility equilibrium where solids are excluded; K = PCO2.

K for gas-phase reactions with Δn

K relates to Kp or Kc depending on gas moles; Δn ≠ 0 means Kp ≠ Kc.

Kc vs Kp units

K may be dimensionless or have units depending on Δn; often treated as dimensionless when activities are used.

Pure solids and liquids excluded

In equilibrium expressions, pure solids and liquids do not appear because their activities are constant.

Solubility equilibrium example

BaSO4(s) ⇌ Ba2+(aq) + SO4^2−(aq); Ksp = [Ba2+][SO4^2−].

Partial pressure

The pressure contribution of a single gas in a mixture.

Saturation / Saturated solution

A solution in which no more solute can dissolve at the given temperature.

Selective precipitation

Separating ions by precipitating one ion while others remain in solution due to differing solubilities.

Catalyst effect on equilibrium

A catalyst speeds up both forward and reverse reactions equally, changing rate but not the position of equilibrium.

Solubility product and CO2 (CaCO3 example)

In CaCO3 ⇌ CaO + CO2, solids are excluded, so the equilibrium expression involves only CO2 (K = PCO2).

Partial pressure in equilibrium expressions

For gas phase equilibria, the expression often involves the partial pressures of gaseous species.

Solubility product and precipitation in water treatment

Ksp governs when solids dissolve or precipitate, enabling selective precipitation for purification.

Haber Process optimization (conceptual)

Industrial optimization uses high pressure and moderate temperature to maximize NH3 yield.

Effect of temperature (Le Châtelier)

Exothermic: increasing T shifts to reactants; Endothermic: increasing T shifts to products.

Effect of pressure (Le Châtelier)

Increasing pressure shifts toward fewer moles of gas; decreasing pressure toward more moles.

Effect of catalysts (Le Châtelier)

Catalysts speed up both forward and reverse reactions without changing equilibrium position.

Gaseous equilibria and partial pressures

Kp describes the relative pressures of gases at equilibrium for a reversible reaction.

Solubility product constant (Ksp) example

Ksp provides the basis to calculate solubility and predict precipitation for salts like BaSO4 or PbCl2.