Acid-Base Equilibria and Solubility Equilibria

These flashcards cover key vocabulary and concepts from Chapter 16 focused on common ion effects, buffer chemistry, titration methods, solubility equilibria, and complex ions.

Common ion effect

The shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance, which suppresses the ionization of a weak acid or a weak base.

Henderson-Hasselbalch equation

The equation used to calculate the $pH$ of a buffer solution: $pH = pKa + \log \frac{[\text{conjugate base}]}{[\text{acid}]}$.

Buffer solution

A solution of a weak acid or a weak base and its salt that has the ability to resist changes in $pH$ upon the addition of small amounts of either acid or base.

Titration

A technique where a solution of accurately known concentration is added gradually to another solution of unknown concentration until the chemical reaction between the two solutions is complete.

Equivalence point

The point in a titration at which the chemical reaction is complete.

Indicator

A substance that changes color at or near the equivalence point of a titration, such as phenolphthalein, methyl red, or thymol blue.

Solubility product constant ($K_{sp}$)

The equilibrium constant for the dissolution of a slightly soluble ionic solid in aqueous solution.

Molar solubility ($s$)

The number of moles of solute dissolved in $1\,L$ of a saturated solution, expressed in units of $mol/L$.

Solubility

The number of grams of solute dissolved in $1\,L$ of a saturated solution, expressed in units of $g/L$.

Reaction Quotient ($Q$) in Solubility

A value used to predict precipitation: if $Q > K_{sp}$, the solution is supersaturated and a precipitate will form; if $Q < K_{sp}$, it is unsaturated; if $Q = K_{sp}$, it is saturated.

Complex ion

An ion containing a central metal cation bonded to one or more molecules or ions.

Formation constant ($K_f$)

The equilibrium constant for complex ion formation, representing the stability of the complex; also known as the stability constant.

Qualitative Analysis

The systematic separation and identification of cations into groups based on their specific precipitation reactions with various reagents like $HCl$, $H_2S$, or $NaOH$.

Carbonic anhydrase

An enzyme that catalyzes the reaction $CO_2 (g) + H_2O (l) \rightleftharpoons H_2CO_3 (aq)$, which is essential in biological processes such as eggshell formation.

$$pKa$$

The negative base-10 logarithm of the acid ionization constant ($Ka$), defined as $pKa = -\log(Ka)$, used in the Henderson-Hasselbalch equation.

Le Chatelier's Principle\n\n

If a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change and re-establish equilibrium.\n\n

Acid-Base Reaction\n\n

A chemical reaction that occurs between an acid and a base, resulting in the formation of water and a salt.\n\n

pH Scale\n\n

A logarithmic scale used to specify the acidity or basicity of an aqueous solution, ranging from 0 (acidic) to 14 (basic), with 7 being neutral.\n\n

Arrhenius Acid\n\n

A substance that increases the concentration of hydrogen ions (H+) in aqueous solution.\n\n

Arrhenius Base\n\n

A substance that increases the concentration of hydroxide ions (OH-) in aqueous solution.\n\n

Bronsted-Lowry Acid\n\n

A proton (H+) donor in a chemical reaction.\n\n

Bronsted-Lowry Base\n\n

A proton (H+) acceptor in a chemical reaction.\n\n

Buffer Capacity\n\n

The ability of a buffer solution to resist changes in pH upon the addition of small amounts of acid or base.\n\n

Hydrolysis\n\n

A chemical reaction involving the breaking down of a compound by reaction with water, often resulting in the formation of an acid and a base.\n\n