Chemical Equilibria Lecture Notes

A collection of vocabulary flashcards covering the fundamental concepts of chemical equilibria, including reaction kinetics, Le Châtelier’s principle, and the Haber process.

Dynamic equilibrium

A reversible reaction in which the rates of the forward and reverse reactions have become equal and there is no net change in the concentrations of the products and reactants.

Equilibrium law

For a reversible reaction at equilibrium, the ratio of the concentration of the products to the concentration of the reactants, each raised to the power of their stoichiometric coefficient, is constant at a given temperature.

Equilibrium constant ($K_c$)

For a generic reversible reaction $aA + bB \rightleftharpoons cC + dD$, it is defined as $K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$ and represents the ratio $\frac{k_{\text{forward}}}{k_{\text{reverse}}}$.

Reaction quotient ($Q_c$)

The quantity of the relative proportion of [products] versus [reactants] at any one time, denotable by the same expression as the equilibrium constant ($K_c$).

Equilibrium position

A description of the relative proportion of [products] versus [reactants] within a system.

ICE table

A table used to track Initial, Change, and Equilibrium concentrations or pressures to solve for unknown quantities in a chemical reaction.

Le Châtelier’s principle

States that the equilibrium position of a reversible reaction will shift in the direction that offsets a disturbance (such as changes in concentration, pressure, or temperature) so as to re-establish dynamic equilibrium.

Exothermic reaction (Equilibrium Shift)

A reaction where the forward process releases energy (e.g., the Haber process); an increase in temperature causes the equilibrium position to shift left, decreasing the value of $K_c$.

Endothermic reaction (Equilibrium Shift)

A reaction where the forward process absorbs energy (e.g., the dissociation of $N_2O_4(g) \rightleftharpoons 2NO_2(g)$, $\Delta H > 0$); an increase in temperature causes the equilibrium position to shift right, increasing the value of $K_c$.

Catalyst effect on equilibrium

The presence of a catalyst (such as iron in the Haber process) provides an alternative reaction mechanism with lower activation energies ($E_a$), increasing both $k_{\text{forward}}$ and $k_{\text{reverse}}$ by the same extent, thus neither $Q_c$ nor $K_c$ changes and the equilibrium position remains unchanged.

Haber process

An industrial process for producing ammonia: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, with $\Delta H = -92\,kJ\,mol^{-1}$, typically conducted at $450$ to $500\,^\circ\text{C}$ and $200$ to $300\,atm$ using a finely-divided $Fe(s)$ catalyst.

Standard Gibbs free energy change relation ($\Delta G^\theta$)

The relationship given by $\Delta G^\theta = -RT \ln K_c$, where $R = 8.31\,J\,K^{-1}\,mol^{-1}$ and $T$ is the temperature in $K$. A more negative $\Delta G^\theta$ results in a value of $K_c$ larger than $1$.

Partial pressure ($p$)

The pressure exerted by an individual gas in a mixture, calculated as a direct proportion of its mole fraction with respect to the total pressure, since $p \propto [ ]$ based on the ideal gas law $pV = nRT$.

Reaction Quotient and Equilibrium direction

If $Q_c < K_c$, there is a net forward reaction until $Q_c = K_c$; if $Q_c > K_c$, there is a net reverse reaction until $Q_c = K_c$.