Acids, Bases, and Salts (Chemistry)

01. About Acids

Definition of an Acid

  • A substance that reacts with a base to form salt and water, typically described as:

    • Corrosive - Can damage or destroy other materials.

    • pH Level - Less than 7 indicates acidity.

    • Litmus Test - Turns blue litmus paper red.

    • Taste - Often has a sour taste.

    • Electrolytic Properties - Can conduct electric current.

Examples of Acids

  • Carbonic Acid - Used in aerated drinks.

  • Sulphuric Acid - Used in car batteries and paint.

  • Tartaric Acid - Important in wine manufacturing.

  • Nitric Acid - Commonly used in fertilizers.

  • Hydrochloric Acid - Plays a role in digestion.

  • Salicylic Acid - Found in aspirin.

  • Acetic Acid - Known for vinegar.

  • Citric Acid - Present in citrus fruits.

  • Laboratory Acids - Used in various lab settings.

Composition of Acids

  • All acids contain Hydrogen (e.g., HCl, HF, HNO3).

  • Can exist in anhydrous forms (i.e not dissolved in water) as solid, liquid, or gas at room temperature:

    • Solids: Citric and tartaric acid.

    • Liquids: Nitric and sulfuric acid.

    • Gas: Hydrogen chloride.

Ionization of Acids

  • To carry an electric current, ions must exist and move freely. In water:

    • Acids ionize to form H+ (hydrogen ions) and negative anions.

    • Hydronium Ions form when H+ ions associate with water molecules.

Example of Ionization

  • Hydrochloric Acid Ionization:

    • Reaction: HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq).

    • Simplified formula: HCl(aq) ⟶ H+(aq) + Cl-(aq).

Other Ionization Examples

  • Nitric Acid: HNO3(aq) ⟶ H+(aq) + NO3-(aq).

  • Sulfuric Acid: H₂SO4(aq) ⟶ 2H+(aq) + SO4^(2-)(aq).

  • Phosphoric Acid: H3PO4(aq) ⟶ 3H+(aq) + PO4^(3-)(aq).

  • Acetic Acid: CH3COOH(aq) ⟶ CH3COO-(aq) + H+(aq).

Acids as Proton Donors

  • Acids donate protons (H+) in reactions, known as proton donors.

    • When hydrogen loses an electron, it becomes a proton.

Classification of Acids

  • Acids are categorized as:

    • Strong vs Weak Acids - Complete vs partial ionization in water.

    • Dilute vs Concentrated Acids - Amount of solute in solution.

    • Basicity - Monobasic, dibasic, tribasic.

    • Acid Anhydrides - Compounds forming acids when reacting with water.

Strong vs Weak Acids

  • Strong Acids: Fully ionize in water (e.g., sulfuric, nitric, Hydrobromic

    Acid, Hydroiodic Acid, Hydrochloric Acid, Perchloric Acid).

  • Weak Acids: Partially ionize (e.g., hydrofluoric, citric, Ethanoic

    Acid, Carbonic Acid, Phosphoric acid).

Dilute vs Concentrated Acids

  • Dilute Acid: High water content, e.g., hydrochloric acid (0.1 mol dm⁻³).

Basicity/Proticity

  • Basicity refers to the number of H+ ions produced per molecule in solution.

Acid Anhydrides

  • Compounds that react with water to form acids. Examples include carbon dioxide and sulfur trioxide.

Organic vs Inorganic Acids

  • Inorganic Acids: Contain hydrogen and non-metallic elements (e.g., HCl, H2SO4).

  • Organic Acids: Contain carboxyl groups (COOH), e.g., hydrofluoric acid, acetic acid.

02. About Bases

Definition of a Base

  • Reacts with an acid to form salt and water (neutralization).

    • Includes metal oxides, metal hydroxides, and metal carbonates.

Reactions with Acids

  • Bases act as proton acceptors, forming water when OH^- ions combine with H+ ions.

  • Classifications include strong and weak bases.

Strong vs Weak Bases

  • Strong Bases: Fully ionize in solution. Note all strong bases contain OH ions (alkalis)

  • Weak Bases: Partially ionize; e.g., ammonia (NH3).

Differences Between Bases and Alkalis

  • Bases: Insoluble in water (e.g. CuO)

  • Alkalis: Soluble bases (e.g., NaOH).

Properties of Aqueous Alkalis

  • Presence of OH- ions confers characteristic properties:

    • Bitter taste, changes red litmus to blue, pH > 7.

    • Corrosive and feel soapy when touched.

Amphoteric Oxides and Hydroxides

  • Certain metal oxides/hydroxides that react as either acids or bases (amphoteric).

Examples and Reactions of Amphoteric Oxides

  • Aluminium (Oxide & Hydroxide)

  • Zinc (Oxide & Hydroxide)

  • Lead (Oxide & Hydroxide)

Classification of Oxides

  • Acidic Oxides: React with bases.

  • Basic Oxides: React with acids.

  • Neutral Oxides: Do not react with acids or bases.

  • Amphoteric Oxides: React with both.

Acidic Oxides

  • Non-metals that produce salts and water upon reaction with bases.

  • Example reactions include:

    • NaOH + CO2 ⟶ Na2CO3 + H2O.

Basic Oxides

  • Metal oxides responding with acids to yield salts and water.

  • Example: CuO + 2HNO3 ⟶ Cu(NO3)2 + H2O.

Amphoteric Oxides

  • Metal oxides reacting with both acids and strong alkalis.

  • Example: PbO with HNO3 and NaOH producing salts and water.

Neutral Oxides

  • Do not participate in acid or base reactions (e.g., CO, NO).

03. pH Scale

Understanding the pH Scale

  • Measures acidity/alkalinity; ranges from 0 to 14:

    • pH < 7: Acidic (stronger with lower values).

    • pH = 7: Neutral.

    • pH > 7: Alkaline (stronger with higher values).

Measurement Methods

  • Universal Indicator: Color indicates pH level.

  • pH Meter: An electronic instrument providing precise pH readings.

  • Indicators: Determine whether a solution is acidic/alkaline.

Indicators Used in Laboratory

  • Litmus: Red in acid, blue in alkaline.

  • Phenolphthalein: Colorless in acidic, pink in alkaline.

  • Methyl Orange: Red in acid, yellow in alkaline.

  • Screen Methyl Orange: Red in acid, green in alkaline.

  • Bromothymol Blue: Yellow in acid, blue in alkaline.