Acids, Bases, and Salts

Acids

Acids are substances often associated with sour tastes, such as those found in unripe fruits like lemons, limes, and oranges. Palm wine turns sour when left open due to the formation of acid.

An acid is a substance that produces hydrogen ions (H⁺) as the only positive ions when dissolved in water. For instance, hydrochloric acid (HCl) dissociates in water to produce hydrogen ions and chloride ions:

HCl → H⁺ + Cl⁻

This process is known as ionization.

Classes of Acids

• Organic acids: Naturally found in plants and animals (e.g., citric acid, ethanoic acid).
• Inorganic acids: Synthesized from minerals (e.g., HCl, H₂SO₄).

Strength of Acids

The strength of an acid refers to how completely it ionizes in water.

Strong Acids

These acids ionize completely in water and are also strong electrolytes. Examples include:

• Hydrochloric acid (HCl)
• Sulphuric acid (H₂SO₄)
• Nitric acid (HNO₃)

HCl → H⁺ + Cl⁻ (100% ionization)

Weak Acids

These ionize only partially in water and are weak electrolytes. Examples include:

• Ethanoic acid (CH₃COOH)
• Carbonic acid (H₂CO₃)
• Hydrofluoric acid (HF)

CH₃COOH ⇌ CH₃COO⁻ + H⁺ (partial ionization)

Basicity of an Acid

The basicity of an acid is the number of replaceable hydrogen ions (H⁺) in one molecule of the acid.

Physical Properties of Acids

• They turn blue litmus paper red.
• They taste sour.
• Concentrated strong acids are corrosive.

Preparation of Acids

1. By dissolving non-metallic oxides in water:
   CO₂(g) + H₂O(l) → H₂CO₃(aq)
2. Direct combination of elements:
   H₂(g) + Cl₂(g) → 2HCl(g) (in presence of catalyst)
3. Hydrogen and bromine reaction (with platinum catalyst):
   H₂(g) + Br₂(g) → 2HBr(g)
4. Displacement of a weak acid from its salt by a stronger acid:
   NaCl(s) + H₂SO₄(aq) → NaHSO₄(aq) + HCl(g)

Chemical Reactions of Acids

1. Reaction with Metals

• Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)
• Mg(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂(g)

Note: Dilute HNO₃ does not produce hydrogen gas with metals.

2. Reaction with Bases and Alkalis (Neutralization)

• HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)
• H₂SO₄(aq) + CaO(aq) → CaSO₄(aq) + H₂O(l)

3. Reaction with Carbonates and Hydrogen Carbonates

• Na₂CO₃(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g)
• NaHCO₃(s) + HCl(aq) → NaCl(aq) + H₂O(l) + CO₂(g)

Uses of Acids

Bases

A base is a substance that neutralizes an acid to form only salt and water. Bases are typically oxides or hydroxides of metals, such as sodium oxide (Na₂O) or magnesium hydroxide (Mg(OH)₂).

An alkali is a base that dissolves in water. All alkalis are bases, but not all bases are alkalis. Examples of alkalis include sodium hydroxide (NaOH), potassium hydroxide (KOH), and calcium hydroxide (Ca(OH)₂).

Oxides of heavy metals like PbO, ZnO, and CuO are insoluble in water and are considered bases, not alkalis. Slightly soluble oxides like CaO and MgO are also alkalis.

Strength of an Alkali

The strength of an alkali is determined by how completely it ionizes in water:

• Strong alkalis ionize completely in water (e.g., NaOH, KOH, Ca(OH)₂, Sr(OH)₂).
• Weak alkalis ionize only partially in water (e.g., aqueous ammonia, NH₄OH).

Characteristics of Bases

• They feel soapy to the touch (e.g., NaOH).
• They have a bitter taste (e.g., lime water).
• They turn red litmus paper blue.
• Concentrated alkalis like NaOH and KOH are corrosive.
• Bases conduct electricity (they are electrolytes).

Preparation of Bases

1. By burning reactive metals in air:
   2Ca + O₂ → 2CaO
2. By reacting metals with water (steam):
   Ca + H₂O → Ca(OH)₂ + H₂
3. By thermal decomposition of metal hydroxides:
   Ca(OH)₂ → CaO + H₂O
   Cu(OH)₂ → CuO + H₂O
4. By precipitation (double decomposition):
   CuSO₄ + 2NaOH → Cu(OH)₂ + Na₂SO₄
5. By dissolving metallic oxides in water:
   Na₂O + H₂O → 2NaOH
   K₂O + H₂O → 2KOH

Reactions of Bases

1. Reaction with Acids (Neutralization)

NaOH + HCl → NaCl + H₂O

2. Thermal Decomposition of Hydroxides

• Zn(OH)₂ → ZnO + H₂O

3. Reaction with Ammonium Salts (with heat)

• 2NH₄Cl + Ca(OH)₂ → CaCl₂ + 2NH₃ + 2H₂O
• NH₄NO₃ + NaOH → NaNO₃ + NH₃ + H₂O

Uses of Bases

• NaOH: Used as a drying agent due to its high solubility.
• Ca(OH)₂: Used to neutralize soil acidity and in cement, whitewash, plaster of Paris, and sugar refining.
• Mg(OH)₂: Found in toothpaste.
• NH₄OH: Used as a mild cleanser and grease remover.
• NH₃: Used in making fertilizers and detergents.
• Bases in general: Used in soap production.

The pH Scale

The term pH stands for the "hydrogen ion index." It is a scale from 0 to 14 used to measure the acidity or alkalinity of a solution:

• pH < 7 → acidic solution
• pH = 7 → neutral solution
• pH > 7 → alkaline solution

The pH meter is an instrument used to measure pH.

Logarithmic pH Scale

The logarithmic pH scale was introduced by Sørensen in 1909 to simplify the representation of hydrogen ion concentration. He defined:

pH = -log₁₀[H⁺]

For example, if [H⁺] = 1 × 10⁻⁵ mol/dm³:

• pH = -log(1 × 10⁻⁵) = 5

Relationship Between [H⁺] and [OH⁻]

[H⁺] × [OH⁻] = 1 × 10⁻¹⁴

pH + pOH = 14

pOH = 14 − pH

Note: A higher pH value means a lower [H⁺] concentration (weaker acidity), while a higher [OH⁻] concentration indicates stronger alkalinity.

Salts

A salt is a compound formed when all or part of the replaceable hydrogen ions in an acid are substituted with metallic or ammonium ions. Salts are made up of positively charged metal ions and negatively charged acid ions.

Characteristics of Salts

1. Water of Crystallization: Some salts chemically combine with a specific amount of water molecules when forming crystals. These are called hydrated salts. When heated, they lose their water and become anhydrous, losing their crystalline form.
   Examples:
   • Sodium trioxosulphate(IV) decahydrate: Na₂SO₃·10H₂O
   • Iron(II) tetraoxosulphate(VI) heptahydrate: FeSO₄·7H₂O
   • Sodium trioxocarbonate(IV) decahydrate: Na₂CO₃·10H₂O
2. Efflorescence: This is the loss of some or all of the water of crystallization from a salt when exposed to air. Example: Na₂CO₃·10H₂O (washing soda).
3. Deliquescence: Substances that absorb moisture from the air and dissolve in it. Examples: CaCl₂, FeCl₃, CuCl₂, ZnCl₂.
4. Hygroscopicity: Substances that absorb moisture from the air without dissolving. They only become moist or sticky. Examples: sodium trioxonitrate(V), potassium trioxonitrate(V).

Hygroscopic substances are often used as drying agents. However, drying agents must not react with the gas being dried. For instance, concentrated H₂SO₄ cannot dry ammonia gas because they react:

2NH₃ + H₂SO₄ → (NH₄)₂SO₄

Preparation of Salts

The method of preparation depends on the salt’s solubility in water and its thermal stability.

A. Preparation of Soluble Salts

1. Acid + Metal: A reactive metal displaces hydrogen from an acid.
   Zn + 2HCl → ZnCl₂ + H₂
2. Acid + Alkali: A neutralization reaction between an acid and an alkali forms a salt.
   KOH + HNO₃ → KNO₃ + H₂O
3. Acid + Insoluble Base: Heating a dilute acid and adding an insoluble base until no more dissolves.
   2HCl + CuO → CuCl₂ + H₂O

B. Preparation of Insoluble Salts

1. Double Decomposition: Mixing two soluble salts to precipitate an insoluble one.
   H₂SO₄ + CuO → CuSO₄ + H₂O
2. Combination of Elements: Direct reaction between the elements forming the salt.
   Fe + S → FeS
   2Fe + 3Cl₂ → 2FeCl₃

Types of Salts

1. Normal Salts: Formed when all the hydrogen ions in an acid are replaced by metal ions.
   H₂SO₄ + ZnO → ZnSO₄ + H₂O
   HCl + NaOH → NaCl + H₂O
2. Acid Salts: Formed when only part of the replaceable hydrogen ions are replaced. They turn blue litmus red.
   H₂SO₄ + NaOH → NaHSO₄ + H₂O
3. Basic Salts: Formed when a base is not fully neutralized by an acid.
   Ca(OH)₂ + HCl → Ca(OH)Cl + H₂O
   Zn(OH)₂ + HCl → Zn(OH)Cl + H₂O
4. Double Salts: Contain more than one type of positive ion. Example:
   Potash alum (KAl(SO₄)₂·12H₂O):
   KAl(SO₄)₂·12H₂O → K⁺ + Al³⁺ + 2SO₄²⁻ + 12H₂O
5. Complex Salts: Contain complex ions.
   Examples:
   • Sodium tetrahydroxozincate(II): Na₂[Zn(OH)₄]
   • Potassium hexacyanoferrate(II): K₄[Fe(CN)₆]

Uses of Salts

• Sodium chloride is used in food preservation.
• Some salts serve as drying agents and antifreeze.
• Salts help in soap making, separating soap from glycerin.
• Sodium carbonate decahydrate (Na₂CO₃·10H₂O) is used for water softening.
• Salts are used in stabilizing unpaved roads.
• Salts are ingredients in toothpaste production.

Hydrolysis of Salts

When salts dissolve in water, they can form acidic or basic solutions through a process called hydrolysis. This happens when the ions of the salt interact with water molecules.

If a salt is formed from a strong acid and a weak base, its solution will be acidic. If it’s formed from a weak acid and a strong base, the solution will be basic.

Example:

NH₄Cl + H₂O → NH₄OH + HCl