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 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.

• 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): Main component of vinegar; a carboxylic acid.

• Carboxylic Acids: Organic acids containing the –COOH group, such as citric acid and malic acid.

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.

• 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): 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 donor (gives H+).

• Base: Proton acceptor (receives H+).

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

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.

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

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

Reactions with Metals and Metal Oxides

• Acids react with many metals to produce hydrogen gas and a salt.

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

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.

Example Calculation

• Given: 10.00 mL HCl solution, 12.54 mL of 0.100 M NaOH required for neutralization.

• Find: Concentration of HCl.

Solution steps:

1. Calculate moles of NaOH used: $12.54\ \mathrm{mL} \times \frac{1\ \mathrm{L}}{1000\ \mathrm{mL}} \times 0.100\ \mathrm{mol/L} = 1.25 \times 10^{-3}\ \mathrm{mol}$

2. From the reaction stoichiometry (1:1), moles of HCl = moles of NaOH.

3. Calculate molarity: $\frac{1.25 \times 10^{-3}\ \mathrm{mol}}{0.01000\ \mathrm{L}} = 0.125\ \mathrm{M}$

Strong and Weak Acids and Bases

Strong Acids

• Completely ionize in solution.

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

Weak Acids

• Partially ionize in solution; equilibrium exists between ionized and unionized forms.

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

Strong Bases

• Completely dissociate in solution to give OH− ions.

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

Weak Bases

• Partially react with water to produce OH− ions.

• Examples: Ammonia (NH3), organic amines.

Water: Acid and Base in One

Self-Ionization of Water

Water can act as both an acid and a base, undergoing self-ionization to produce equal concentrations of H3O+ and OH− at 25°C:

• $[\mathrm{H_3O^+}] = [\mathrm{OH^-}] = 1.0 \times 10^{-7}\ \mathrm{M}$

• $K_w = [\mathrm{H_3O^+}][\mathrm{OH^-}] = 1.0 \times 10^{-14}$

The pH and pOH Scales

Definition and Calculation

• pH: $\mathrm{pH} = -\log [\mathrm{H_3O^+}]$

• pOH: $\mathrm{pOH} = -\log [\mathrm{OH^-}]$

• At 25°C: $\mathrm{pH} + \mathrm{pOH} = 14$

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

Buffers

Definition and Function

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.

Environmental and Health Connections

Acid Rain

Acid rain results from the reaction of sulfur and nitrogen oxides (from fossil fuel combustion) with water, forming sulfuric and nitric acids. Acid rain can damage buildings and ecosystems by dissolving metals and metal oxides.

Health Example: Antifreeze Poisoning

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