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 behaviors. Their study is essential for understanding chemical reactions, biological processes, and industrial applications.

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 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 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 skin oils 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 industrial processes.

• Sodium Bicarbonate (NaHCO3): 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 does not account for bases that do not contain OH− or for reactions in nonaqueous solvents.

Brønsted–Lowry Definition

• Acid: Proton (H+) donor.

• Base: Proton (H+) acceptor.

This definition is broader and includes more substances, such as ammonia (NH3), which acts as a base by accepting a proton from water.

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 (base) and NH4+ (conjugate acid).

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

• Gas Evolution: Acids react with carbonates or bicarbonates to produce water, carbon dioxide gas, and a salt.

Acids Reacting with Metals

Acids react with many metals to produce hydrogen gas and a dissolved salt containing the metal ion. Not all metals react; for example, gold does not dissolve in most acids.

Acids Reacting with Metal Oxides

Acids react with metal oxides to produce water and a dissolved salt. For example, hydrochloric acid reacts with potassium oxide to form water and potassium chloride.

Bases Reacting with Metals

Some metals, such as aluminum, can dissolve in strong bases like sodium hydroxide, producing hydrogen gas and a soluble aluminate ion.

Acid–Base Titration

Quantifying Acid or Base Concentration

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.

Strong and Weak Acids and Bases

Strong Acids

Strong acids completely ionize in solution, producing a high concentration of H+ (or H3O+) ions. Examples include HCl, HBr, HI, HNO3, HClO4, and H2SO4 (first ionization only).

Weak Acids

Weak acids only partially ionize in solution, resulting in an equilibrium between the acid and its ions. Most carboxylic acids are weak acids.

Strong Bases

Strong bases completely dissociate in solution to produce OH− ions. Examples include NaOH, KOH, and Ba(OH)2.

Weak Bases

Weak bases only partially react with water to produce OH− ions. Ammonia (NH3) and organic amines are common weak bases.

Water: Acid and Base in One

Self-Ionization of Water

Water is amphoteric, meaning it can act as both an acid and a base. In pure water, a small amount of self-ionization occurs:

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

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

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

The pH and pOH Scales

pH Scale

• pH = –log[H3O+]

• pH < 7: acidic; pH = 7: neutral; pH > 7: basic (at 25°C)

• The pH scale is logarithmic: a decrease of 1 unit means a tenfold increase in [H3O+].

pOH Scale

• pOH = –log[OH−]

• pH + pOH = 14 (at 25°C)

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. Human blood is a natural buffer system, maintaining pH between 7.36 and 7.40.

Chemistry and Health: Acid Rain and Antifreeze Poisoning

Acid Rain

Acid rain results from sulfur oxides and nitrogen oxides reacting with water in the atmosphere to form acids, which can damage buildings and ecosystems.

Antifreeze Poisoning

Ethylene glycol (antifreeze) is metabolized to glycolic acid, which can overwhelm the body's buffer system, leading to dangerously low blood pH and potentially fatal consequences.