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 industrial processes.

Properties of Acids

• Sour Taste: Acids typically have a sour taste, as experienced with citric acid in lemons or tartaric acid in candies.

• Reaction with Metals: Acids can dissolve many metals, producing hydrogen gas and a salt.

• Effect on Indicators: Acids turn blue litmus paper red.

Examples of Acids:

• Hydrochloric Acid (HCl): Found in stomach acid and used in industry for cleaning metals and processing foods.

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

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

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

• Carboxylic Acids: Organic acids containing the –COOH group, found in many biological substances.

Properties of Bases

• Bitter Taste: Bases have a bitter taste, which is less common in foods due to evolutionary aversion to potentially toxic alkaloids.

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

• Effect on Indicators: Bases turn red litmus paper blue.

Examples of 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.

Limitations: The Arrhenius definition is restricted to aqueous solutions and cannot explain basicity in substances that do not contain OH− ions.

Brønsted–Lowry Definition

• Acid: Proton donor (gives H+).

• Base: Proton acceptor (receives H+).

This definition applies to a broader range of reactions, including those in nonaqueous solutions.

Conjugate Acid–Base Pairs

Any two substances related by the gain or loss of a proton are called a conjugate acid–base pair. For example, NH3 and NH4+ or H2O and OH−.

Reactions of Acids and Bases

Neutralization Reactions

• When an acid reacts with a base, H+ from the acid combines with OH− from the base to form water.

• The reaction also produces a salt, which is an ionic compound composed of the cation from the base and the anion from the acid.

Net Ionic Equation:

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

Reactions with carbonates or bicarbonates produce water, carbon dioxide gas, and a salt.

Reactions with Metals

• Acids react with many metals to produce hydrogen gas and a dissolved salt containing the metal ion.

• Some metals, such as gold, do not react with most acids.

Reactions with Metal Oxides

• Acids react with metal oxides to produce water and a dissolved salt.

Reactions of Bases

• Bases neutralize acids in solution.

• Some bases, such as NaOH, can react with certain metals like aluminum, dissolving them.

Acid–Base Titration

Principle of Titration

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

Strong and Weak Acids and Bases

Strong Acids

• Completely ionize in solution, producing a high concentration of H3O+ ions.

• Examples: HCl, HBr, HI, HNO3, HClO4, H2SO4 (first proton only).

Weak Acids

• Only partially ionize in solution; most molecules remain intact.

• Examples: HF, acetic acid, formic acid, carbonic acid, phosphoric acid.

Strong Bases

• Completely dissociate in solution to produce OH− ions.

• Examples: NaOH, KOH, Ba(OH)2, Sr(OH)2.

Weak Bases

• Partially react with water to produce OH− ions; most molecules remain unreacted.

• Examples: Ammonia (NH3), organic amines.

Water: Acid and Base in One (Amphoterism)

Water is amphoteric, meaning it can act as either an acid or a base. Water undergoes self-ionization to produce equal concentrations of H3O+ and OH− ions.

$ \mathrm{2H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)} $

At 25°C, $[\mathrm{H_3O^+}] = [\mathrm{OH^-}] = 1.0 \times 10^{-7} \ \mathrm{M}$ and $K_w = 1.0 \times 10^{-14}$.

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: basic; pH = 7: neutral

Buffers

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

Chemistry and Health: Acid Rain and Antifreeze Poisoning

• Acid Rain: Caused by sulfur and nitrogen oxides reacting with water to form acids, which can damage buildings and ecosystems.

• Antifreeze Poisoning: Ethylene glycol is metabolized to glycolic acid, which can overwhelm the body's buffer system and lower blood pH to dangerous levels.