Paragraph 1: Chemical Equilibrium

At dynamic equilibrium, the rate of [1] reaction equals the rate of [2] reaction. The equilibrium constant Kc is defined using [3] concentrations only. For gaseous reactions, Kp and Kc are related by Kp = Kc × ([4])^$\Delta n$, where $\Delta n$ represents the difference in moles of [5] species between products and reactants. The reaction quotient Q uses [6] concentrations; if Q < K, the reaction proceeds [7].

Paragraph 2: Le Chatelier's Principle

When a system at equilibrium is subjected to a stress, it shifts to [8] that stress. Adding more reactant shifts equilibrium [9]. Adding a catalyst [10] the equilibrium position and [11] the value of K. Adding an inert gas at [12] volume has no effect, but at constant [13] it shifts toward more gas moles. Only [14] changes the value of K.

Paragraph 3: Ionic Equilibrium

The pH of a solution is defined as −log[[15]]. For a weak acid HA, the hydrogen ion concentration equals [16]. The Henderson-Hasselbalch equation gives buffer pH = [17]. The product Ka × Kb for a conjugate acid-base pair equals [18]. For a sparingly soluble salt, precipitation occurs when the ionic product exceeds [19]. At 25°C, pH + pOH = [20].

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