CHM 112 Exam 2 Study Outline: Chemical Equilibrium, Acids and Bases, and Solubility

Vocabulary flashcards covering the primary concepts of dynamic equilibrium, Le Chatelier's principle, acid-base theories and strengths, buffers, indicators, polyprotic systems, and solubility equilibria from CHM 112 Exam 2.

Chemical Equilibrium

The dynamic state in which macroscopic observables stop changing, the forward and reverse microscopic reaction rates are equal, and reactants and products are present in fixed relative concentrations independent of initial distribution.

Law of Mass Action (LOMA)

The law stating that for a reversible reaction at equilibrium ($aA + bB \rightleftharpoons cC + dD$), the ratio of the activities of products to reactants raised to their stoichiometric coefficients equals a constant value $K$.

Activity ($a$)

The dimensionless effective concentration or pressure of a chemical species, defined as $\frac{P_A}{P_A^\circ}$ ($P_A^\circ = 1\,\text{atm}$) for gases, $\frac{[A]}{[A]^\circ}$ ($[A]^\circ = 1\,\text{M}$) for aqueous solutions, and exactly $1$ for pure solids and pure liquids.

Equilibrium Constant ($K$)

A unitless number characteristic of a specific reaction whose value depends solely on temperature, indicating whether products are favored ($K$ is large) or reactants are favored ($K$ is small).

Relationship Between $K_p$ and $K_c$

The mathematical formula relating gas equilibrium constants based on pressure and concentration: $K_p = K_c(RT)^{\Delta n}$, where $\Delta n$ is the moles of gaseous products minus the moles of gaseous reactants.

Reaction Quotient ($Q$)

A quantity calculated using the same expression as the equilibrium constant, but evaluated at any conditions to predict reaction direction: if $Q < K$, it proceeds forward; if $Q > K$, it proceeds in reverse; if $Q = K$, it is at equilibrium.

Successive Approximations

A numerical calculation technique used to simplify higher-order equilibrium equations when $K$ is either very small ($\times 10^{-3}$ or smaller) or very large ($\times 10^{3}$ or larger) by assuming the change $x$ is negligible compared to initial values and iteratively refining it.

Le Chatelier's Principle

The principle stating that if a dynamic equilibrium system is perturbed by a change in concentration, pressure, volume, or temperature, the reaction runs in the direction that counteracts and compensates for the change.

Heterogeneous Equilibria

Equilibrium systems containing reactants and products in multiple physical states (solids, liquids, gases, or solutions), wherein pure solids and liquids have an activity of $1$ and do not alter the position of equilibrium upon quantity changes.

Arrhenius Acid and Base

An acid is a substance that produces $H^+$ in aqueous solution, whereas a base is a substance that produces $OH^-$ in aqueous solution.

Brønsted-Lowry Acid and Base

An acid is a species that donates a proton, and a base is a species that accepts a proton.

Lewis Acid and Base

An acid is an electron-pair acceptor, while a base is an electron-pair donor.

Strong Acid

An acid with a very large dissociation constant ($K_a$) that is treated as completely dissociated in water because it is a stronger proton donor than hydronium ($H_3O^+$). Examples include $HCl$, $HBr$, $HI$, $HNO_3$, $HClO_4$, $HClO_3$, and $H_2SO_4$.

Weak Acid

An acid with a small dissociation constant ($K_a$) that is only partially dissociated (rarely more than a few percent) at equilibrium because it is a less effective proton donor than hydronium ($H_3O^+$).

Conjugate See-Saw

The inverse strength relationship between conjugate acid-base pairs: the stronger an acid or base is, the weaker its conjugate base or acid must be.

Halogen Acid ($HX$) Strength Trend

The periodic trend where acid strength increases down Group 17 (halogens) due to longer, weaker $H-X$ bonds resulting from poor orbital overlap between hydrogen and larger halogen atoms.

Oxyacid Strength Trend

The concept that the presence of additional electron-withdrawing oxygen atoms increases the strength of an oxyacid by increasing the polarity of the $H-O$ bond.

Autoionization of Water

The equilibrium reaction $2H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)$, characterized at $25\,^\circ\text{C}$ by the ion-product constant $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$.

pH and pOH

Logarithmic representations of ion concentrations defined by $pH = -\log[H_3O^+]$ and $pOH = -\log[OH^-]$, satisfying $pH + pOH = 14$ in any aqueous solution at $25\,^\circ\text{C}$.

Buffer Solution

A solution typically containing comparable concentrations of a weak acid and its conjugate weak base that maintains an approximately constant $pH$ despite small additions of acid, base, or dilution.

Buffer Capacity

The measure of the amount of acid or base a buffer can absorb for a given change in $pH$, which depends directly on the absolute concentrations of the buffer components rather than their ratio.

Henderson-Hasselbalch Equation

The equation used to determine buffer $pH$: $pH = pK_a + \log\left(\frac{[A^-]_0}{[HA]_0}\right)$, which indicates that an acid is exactly $50\,\%$ dissociated when $pH = pK_a$.

Acid-Base Indicator

A very dilute system consisting of a weak acid and its conjugate base ($HIn + H_2O \rightleftharpoons In^- + H_3O^+$) whose protonated form ($HIn$) dominates at $pH = pK_a - 1$ and deprotonated form ($In^-$) dominates at $pH = pK_a + 1$, each displaying a different color.

Polyprotic Acid

An acid capable of donating more than one proton in a series of successive dissociation equilibria, where successive dissociation constants decrease ($K_{a1} > K_{a2} > K_{a3}$), making each successive step a weaker acid.

Amphiprotic Species

A chemical species that can act as both a proton donor (acid) and a proton acceptor (base), whose salt $pH$ can be approximated by $pH \approx \frac{1}{2}(pK_{a1} + pK_{a2})$.

Solubility

The maximum amount of a solute (measured in grams or moles) that can dissolve in a specified volume of solvent at equilibrium at a particular temperature.

Solubility Product Constant ($K_{sp}$)

The equilibrium constant describing the balance between an undissolved, slightly soluble salt and its dissolved ions in a saturated aqueous solution.

Common Ion Effect

The reduction in the solubility of a slightly soluble salt caused by the presence or addition of an ion already participating in that dissolution equilibrium.

Selective Precipitation

A laboratory separation technique that isolates specific ions from solution based on differences in salt solubilities, requiring their $K_{sp}$ values to differ by several orders of magnitude.