Acids and Bases: Properties, Definitions, and Reactions

Acids and Bases

Introduction to Acids and Bases

Acids and bases are fundamental classes of compounds in chemistry, each with distinct properties and important roles in both laboratory and everyday contexts. Their behavior is central to many chemical reactions, including those in biological and environmental systems.

Properties of Acids

• Sour Taste: Acids typically taste sour. For example, the sour taste of candies like Sour Patch Kids is due to citric and tartaric acids, which release H+ ions that interact with taste receptors.

• Reaction with Metals: Acids can dissolve many metals, producing hydrogen gas and a salt. However, some metals like gold are resistant to acid attack.

• Litmus Test: Acids turn blue litmus paper red.

Examples of Common Acids

• Hydrochloric Acid (HCl): Found in laboratories and the human stomach, used for cleaning metals and processing foods.

• Sulfuric Acid (H2SO4): Widely used in industry, especially in fertilizer and battery production.

• Nitric Acid (HNO3): Used in manufacturing fertilizers and explosives.

• Acetic Acid (HC2H3O2): The main component of vinegar, a carboxylic acid.

• Carboxylic Acids: Organic acids containing the –COOH group, found in many biological substances (e.g., citric acid in lemons, malic acid in apples).

Properties of Bases

• Bitter Taste: Bases taste bitter, which is a natural deterrent against consuming potentially toxic substances (e.g., alkaloids like coniine).

• Slippery Feel: Bases feel slippery because they react with oils on the skin to form soap-like substances.

• Litmus Test: Bases turn red litmus paper blue.

Examples of Common Bases

• Sodium Hydroxide (NaOH): Used in drain cleaners and soap manufacturing.

• Potassium Hydroxide (KOH): Used in similar applications as NaOH.

• Sodium Bicarbonate (NaHCO3): Commonly known as baking soda, used as an antacid.

Definitions of Acids and Bases

Arrhenius Definition

• Acid: Produces H+ ions in aqueous solution.

• Base: Produces OH− ions in aqueous solution.

Brønsted–Lowry Definition

• Acid: Proton donor (gives H+).

• Base: Proton acceptor (receives H+).

This definition is broader and includes substances that do not contain OH− but can accept protons, such as ammonia (NH3).

Amphoteric Substances and Conjugate Acid–Base Pairs

• Amphoteric: Substances like water can act as either an acid or a base.

• Conjugate Acid–Base Pair: Two substances related by the gain or loss of a proton (e.g., NH3/NH4+).

Reactions of Acids and Bases

Neutralization Reactions

When an acid reacts with a base, the H+ from the acid combines with the OH− from the base to form water. The other ions form a salt.

• General Equation: $\mathrm{H^+ (aq) + OH^- (aq) \rightarrow H_2O (l)}$

• With Carbonates/Bicarbonates: Reaction produces water, carbon dioxide gas, and a salt.

Reactions with Metals and Metal Oxides

• Acids + Metals: Produce hydrogen gas and a salt (e.g., $\mathrm{2HCl(aq) + Mg(s) \rightarrow H_2(g) + MgCl_2(aq)}$).

• Acids + Metal Oxides: Produce water and a salt.

Acid–Base Titration

Principle of Titration

Titration is a laboratory technique used to determine the concentration of an unknown acid or base by reacting it with a solution of known concentration. The equivalence point is reached when stoichiometric amounts of acid and base have reacted.

• Indicator: A dye that changes color at (or near) the equivalence point (e.g., phenolphthalein).

• Calculation: Use the volume and concentration of the titrant to find the unknown concentration.

Strength of Acids and Bases

Strong vs. Weak Acids

• Strong Acid: Completely ionizes in solution (e.g., HCl, HNO3).

• Weak Acid: Only partially ionizes in solution (e.g., HF, acetic acid).

Strong vs. Weak Bases

• Strong Base: Completely dissociates in solution (e.g., NaOH, KOH).

• Weak Base: Partially reacts with water to produce OH− (e.g., NH3).

Electrolyte Behavior

• Strong Electrolytes: Solutions of strong acids/bases conduct electricity well due to the presence of many ions.

• Weak Electrolytes: Solutions of weak acids/bases conduct electricity poorly due to fewer ions.

Water and the pH Scale

Self-Ionization of Water

• Water can act as both an acid and a base, undergoing self-ionization: $\mathrm{2H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)}$

• At 25°C, $[\mathrm{H_3O^+}] = [\mathrm{OH^-}] = 1.0 \times 10^{-7}$ M

• The ion product constant for water: $K_w = [\mathrm{H_3O^+}][\mathrm{OH^-}] = 1.0 \times 10^{-14}$

pH and pOH Scales

• pH: $\mathrm{pH = -\log [H_3O^+]}$

• pOH: $\mathrm{pOH = -\log [OH^-]}$

• At 25°C: pH + pOH = 14

• pH < 7: acidic; pH = 7: neutral; pH > 7: basic

Buffers

Buffer Solutions

Buffers are solutions that resist changes in pH when small amounts of acid or base are added. They contain significant amounts of both a weak acid and its conjugate base. Buffers are essential in biological systems, such as human blood, to maintain a stable pH.

Environmental and Health Connections

Acid Rain

Acid rain results from the reaction of sulfur and nitrogen oxides (from fossil fuel combustion) with water in the atmosphere, forming sulfuric and nitric acids. Acid rain can damage buildings, ecosystems, and aquatic environments.

Health Example: Antifreeze Poisoning

Ethylene glycol (antifreeze) is metabolized to glycolic acid in the body, which can overwhelm the blood's buffering capacity, leading to dangerously low blood pH and potentially fatal outcomes if untreated.