Comprehensive Study Guide: Acids, Bases, and Chemical Equilibria

Acids are substances that exhibit the following distinct physical and chemical properties: - They possess a characteristically sour taste. - They cause a color change in chemical indicators; specifically, they turn blue litmus paper red. - They react with active metals (such as magnesium or zinc) to produce hydrogen gas ( $H_2$ ). - They have a $pH$ level of less than $7$ .
Bases are substances characterized by the following properties: - They possess a bitter taste. - They have a slippery or soapy feel to the touch. - They cause chemical indicators to change color; specifically, they turn red litmus paper blue. - They have a $pH$ level of greater than $7$ .

Definition of an Electrolyte: An electrolyte is a substance that, when dissolved in water (aqueous solution), dissociates or ionizes into its constituent ions. These mobile ions allow the solution to conduct an electric current.
Conductivity of Acids and Bases: Both acids and bases are classified as electrolytes. This is because: - Acids ionize in aqueous solutions to produce hydrogen ions ( $H^+$ ) or hydronium ions ( $H_3O^+$ ). - Bases dissociate or react with water to produce hydroxide ions ( $OH^-$ ). - The presence of these freely moving charged particles (cations and anions) in the water enables the flow of electricity.

Hydronium Ion: The chemical formula for the hydronium ion is $H_3O^+_{(aq)}$ . It is formed when a free proton ( $H^+$ ) associates with a water molecule ( $H_2O$ ).
Strong Acids and Strong Bases: - Strong Acid: A strong acid is an acid that dissociates or ionizes completely ( $100\%$ ) in aqueous solution. This means every molecule of the acid releases its proton. Example: Hydrochloric acid ( $HCl$ ). - Strong Base: A strong base is a base that dissociates completely ( $100\%$ ) in aqueous solution, providing a high concentration of hydroxide ions. Example: Sodium hydroxide ( $NaOH$ ).
Protic Classification of Acids: - Monoprotic Acid: An acid capable of donating only one proton ( $H^+$ ) per molecule. Example: Hydrochloric acid ( $HCl$ ). - Diprotic Acid: An acid capable of donating two protons ( $H^+$ ) per molecule in a step-wise manner. Example: Sulfuric acid ( $H_2SO_4$ ). - Triprotic Acid: An acid capable of donating three protons ( $H^+$ ) per molecule. Example: Phosphoric acid ( $H_3PO_4$ ).

Arrhenius Definition: Defines an acid as a substance that increases the concentration of $H^+$ ions when dissolved in water. - a) Hydrochloric Acid: $HCl_{(aq)} \rightarrow H^+{(aq)} + Cl^-{(aq)}$ - b) Nitric Acid ( $HNO_3$ ): $HNO_3(aq) \rightarrow H^+{(aq)} + NO_3^-{(aq)}$ - c) Hydrofluoric Acid: $HF_{(aq)} \rightarrow H^+{(aq)} + F^-{(aq)}$ - d) Acetic Acid: $CH_3COOH_{(aq)} \rightarrow H^+{(aq)} + CH_3COO^-{(aq)}$ - e) Bicarbonate Ion: HCO_3^-{(aq)} \rightarrow H^+{(aq)} + CO_3^{2-}_{(aq)}
Brønsted-Lowry Definition: Defines an acid as a proton ( $H^+$ ) donor and involves the transfer of a proton to a water molecule to form hydronium. - a) Hydrochloric Acid: $HCl_{(aq)} + H_2O_{(l)} \rightarrow H_3O^+{(aq)} + Cl^-{(aq)}$ - b) Nitric Acid: $HNO_3(aq) + H_2O_{(l)} \rightarrow H_3O^+{(aq)} + NO_3^-{(aq)}$ - c) Hydrofluoric Acid: $HF_{(aq)} + H_2O_{(l)} \rightleftharpoons H_3O^+{(aq)} + F^-{(aq)}$ - d) Acetic Acid: $CH_3COOH_{(aq)} + H_2O_{(l)} \rightleftharpoons H_3O^+{(aq)} + CH_3COO^-{(aq)}$ - e) Bicarbonate Ion: $HCO_3^-{(aq)} + H_2O{(l)} \rightleftharpoons H_3O^+{(aq)} + CO_3^{2-}{(aq)}$

Base Dissociation Equations: - a) Sodium Hydroxide: $NaOH_{(aq)} \rightarrow Na^+{(aq)} + OH^-{(aq)}$ - b) Barium Hydroxide: $Ba(OH)2(aq) \rightarrow Ba^{2+}{(aq)} + 2OH^-{(aq)}$ - c) Bicarbonate Ion (acting as a base): $HCO_3^-{(aq)} + H_2O_{(l)} \rightleftharpoons H_2CO_3(aq) + OH^-_{(aq)}$
Amphoterism of Hydrogen Sulfate ( $HSO_4^-$ ): A substance is amphoteric if it can act as both an acid and a base. Based on Brønsted-Lowry theory: - Acting as an Acid: It donates a proton to water: $HSO_4^-{(aq)} + H_2O{(l)} \rightarrow SO_4^{2-}{(aq)} + H_3O^+{(aq)}$ - Acting as a Base: It accepts a proton from water: $HSO_4^-{(aq)} + H_2O{(l)} \rightarrow H_2SO_4(aq) + OH^-_{(aq)}$
Autoionization of Water: Water is the most common amphoteric substance. The dissociation equation is: - $H_2O_{(l)} + H_2O_{(l)} \rightleftharpoons H_3O^+{(aq)} + OH^-{(aq)}$ - It is considered both an acid and a base because one water molecule donates a proton (acting as a Brønsted-Lowry acid) while the other accepts a proton (acting as a Brønsted-Lowry base).

Chemical Formula: The formula for ammonia is $NH_3$ .
Definition as a Base: To classify ammonia as a base, the Brønsted-Lowry definition is most appropriate, as it defines a base as a proton ( $H^+$ ) acceptor ( $NH_3$ lacks a hydroxide group to fit the strict Arrhenius definition).
Dissociation in Water: When ammonia reacts with water, the following equilibrium occurs: - $NH_3(aq) + H_2O_{(l)} \rightleftharpoons NH_4^+{(aq)} + OH^-{(aq)}$
Conjugate Acid-Base Pairs: - Pair 1: $NH_3$ (Base) and $NH_4^+$ (Conjugate Acid). - Pair 2: $H_2O$ (Acid) and $OH^-$ (Conjugate Base).
Mechanism of Basicity: Ammonia is considered a base because its dissociation in water results in the release of hydroxide ( $OH^-$ ) ions. In this reaction, water acts as the proton donor (acid). If water ( $H_2O$ ) were named using binary acid-naming rules ( $\text{hydro-} + \text{root} + \text{-ic acid}$ ), it would be called hydrohydroxic acid.