Revision Notes: Study of Acids Bases and Salts

Acid

An acid is a compound which when dissolved in water yields hydronium ions (H3O+) as the only positively charged ions.

Examples: 

Classification of Acids

1. Depending on sources 
   Organic acid: Acids which are usually obtained from plants are called organic acids. They contain carbon and hydrogen atoms.
   Examples:
   Inorganic (Mineral) acids: Acids which are obtained from minerals are known as inorganic acids. 
   Examples:
   

2. Depending on strength
   (a) Strength of an acid: The strength of an acid depends on the concentration of the hydronium ions (H₃O⁺) present in the aqueous solution of an acid.
   i. Strong acids: A strong acid vigorously ionises in aqueous solution, thereby producing a high concentration of hydronium ions (H₃O⁺).
   Examples: HNO₃, HCl, H₂SO₄
   ii. Weak acids: Weak acids ionise only partially in aqueous solution to produce ions and molecules. 
   Examples: H₂CO₃, CH₃COOH, HCOOH

3. Depending on basicity
   Basicity of an acid: The number of hydronium ions (H₃O⁺) which can be produced by the ionisation of one molecule of that acid in aqueous solution.
   (i) Monobasic acids: Acids which on ionisation in water produce one hydronium ion (H₃O⁺) per molecule of the acid are known as monobasic acids. 
   Example:
   (ii) Dibasic acids: Acids which on ionisation in water produce two hydronium ions (H₃O⁺) per molecule of the acid are known as dibasic acids.
   Examples:
   (iii) Tribasic acids: Acids which on ionisation in water produce three hydronium ions (H₃O⁺) per molecule of the acid are known as tribasic acids.
   Examples: 
   

4. Depending on concentration: The concentration of an acid means the amount of acid present in a definite amount of its aqueous solution.
   i. Concentrated acid: An acid which contains a very small amount of water or no water is called a concentrated acid.
   ii. Dilute acid: An acid which contains far more amount of water than its own mass is known as a dilute acid.
5. Depending on molecular composition
   i. Hydracids: Acids which contain hydrogen, a non-metallic element and no oxygen are called hydracids. 
   Examples: HCl, H₂S, HBr, HI
   ii. Oxyacids: Acids which contain oxygen, hydrogen and a non-metallic element are called oxyacids. 
   Examples: H₂SO₄, HNO₃, H₂CO₃

Preparation of Acids

1. By synthesis
   H₂  + Cl₂ → 2HCl
2. By the action of water on non-metallic or acidic oxides 
   SO₃ + H₂O  →  H₂SO₄  
   N₂O₅  + H₂O  →   2HNO₃ 
3. By oxidation of non-metals
   S + 6HNO₃ → H₂SO₄ + 2H₂O + 6NO₂
   P + H₃PO₄ → H₃PO₄ + H₂O + 5O₂
4. By displacement
   NaCl + H₂SO₄ → NaHSO₄ + HCl
   NaNO₃ + H₂SO₄ → NaHSO₄ + HNO₃

Properties of Acids

Physical properties
i. Sour in taste in aqueous solution.
ii. Turns blue litmus red.
iii. Some acids are solids and some are liquids at room temperature.
iv. All strong mineral acids have corrosive action on the skin and cause painful burns.
v. They are electrolytes, i.e. they conduct electricity in the aqueous state.

Chemical Properties

1. Reaction with active metals
   Mg + 2HCl → MgCl₂ + H₂
2. Reaction with bases - Neutralisation
   NaOH + H₂SO₄ → NaNO₃ + H₂O
3. Reaction with carbonates and bicarbonates
   CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
4. Reaction with sulphites and bisulphites
   CaSO₃ + 2HCl → CaCl₂ + H₂O + SO₂
   NaHSO₃ + HCl → NaCl + H₂O + SO₂
5. Reaction with sulphides
   ZnS + 2HCl → ZnCl₂ + H₂S
6. Reaction with chlorides
   
7. Reaction with nitrates
   Pb (NO₃)₂ + 2HCl → PbCl₂ + 2HNO₃

Uses of Some Acids

Bases

A base is either a metallic oxide or a metallic hydroxide or ammonium hydroxide which reacts with hydronium ions of an acid to form salt and water only.

Basic Oxide

A basic oxide is a metallic oxide which contains the ion O2- and reacts with an acid to form salt and water.
Alkalis
An alkali is a basic hydroxide which when dissolved in water produces hydroxyl (OH-) ions as the only negatively charged ions.

Note: All alkalis are bases, but all bases are not alkalis.

Classification of Bases

1. On the basis of strength
   i. Strong base: It undergoes almost complete ionisation in aqueous solution to produce a high concentration of OH^- ions.
   Example:
   ii. Weak base: It undergoes only partial ionisation in aqueous solution to produce a low concentration of OH^- in solution.
   Example:
2. On the basis of acidity
   a. Acidity of a base: The number of hydroxyl ions (OH^-) which can be produced per molecule of the base in aqueous solution.
   i. Monoacidic base: Bases which dissociate in aqueous solution to produce one hydroxyl ion (OH-) per molecule of the base are called monoacidic bases.
   Example: NaOH ⇌ Na⁺ + OH^- [Acidity = 1] 
   ii. Diacidic base: Bases which dissociate in aqueous solution to produce two hydroxyl ions (OH ) per molecule of the base are called diacidic bases.
   Example: Ca(OH)₂ ⇌ Ca²⁺ + 2OH^- [Acidity = 2]
   iii. Triacidic base: Bases which dissociate in aqueous solution to produce three hydroxyl ions (OH ) per molecule of the base are called triacidic bases.
   Example: Al (OH)₃ ⇌ Al ³⁺ + 3OH^- [Acidity = 3]
   iv. By oxidation of non-metals
   S + 6HNO₃H₂SO₄ + 2H₂O + 6NO₂
3. On the basis of composition
   Concentrated alkali: It is an alkali with a relatively high percentage of alkali in its aqueous solution. 
   Dilute alkali: It is an alkali with a relatively low percentage of alkali in its aqueous solution.

Preparation of Bases

i. From Metals
2Mg + O₂ → 2MgO
ii. By action of water or steam on reactive metals
2Na + 2H₂O → 2NaOH + H₂
iii. By the action of water on soluble metallic oxides
Na₂O + H₂O → 2NaOH
iv. By double decomposition
FeCl₃ + 3NaOH → Fe (OH)₃ + 3NaCl
v. By the action of oxygen on metal sulphides
2ZnS + 3O₂ → 2ZnO + 2SO₂
vi. By decomposition of salts
CaCO₃ → CaO + CO₂

Properties of Bases

Physical properties

1. They have sharp and bitter taste.
2. They change red litmus blue.
3. Soapy substances, i.e. they are slippery to touch.
4. They are strong electrolytes.
5. They show mild corrosive action on the skin.

Chemical properties

1. Reaction with carbon dioxide
   2NaOH + CO₂ → Na₂CO₃ + H₂O
2. Reaction with acids - Neutralisation
   Ca (OH)₂ + 2HCl → CaCl₂ + 2H₂O
3. Reaction with metallic salts
   CuSO₄ + 2NH₄OH → (NH₄)₂SO₄ + Cu (OH)₂

Uses of Some Bases

pH Value

It represents the strength of acids and alkalis expressed in terms of hydrogen ion concentration.

pH of Solution

pH of a solution is the negative logarithm to the base 10 of the hydrogen ion concentration expressed in mole per litre. 
pH = -log₁₀ (H⁺)

pH Scale

It is a scale showing the relative strength of acids and alkalis.
The normal pH scale ranges from 0 to 14 as shown below.

Indicators

They are complex substances which acquire separate colours in acidic and basic media.

Types of Indicators

a. Acid-base indicators: Common acid-base indicators such as litmus, methyl orange and phenolphthalein can distinguish between acid and basic solutions, but they cannot determine the strength of the solution.
b. Universal indicator: A universal indicator is a mixture of organic dyes which gives a definite colour change over a wide range of pH.

Salts

A salt is a compound formed by the partial or total replacement of the ionisable hydrogen atoms of an acid by a metallic ion or an ammonium ion.

Classification of Salts

1. Normal salts: The salts formed by the complete replacement of the replaceable hydrogen ion of an acid molecule by a basic radical.
   Example: HCl + NaOH → NaCl + H₂O
2. Acid salts: The salts formed by partial replacement of the replaceable hydrogen ion of an acid molecule by a basic radical.
   Example: NaOH + H₂SO₄ → NaHSO₄ + H₂O
3. Basic salts: The salts formed by the partial replacement of the hydroxyl group of a di- or tri-acidic base by an acidic radical.
   Example: Mg (OH)₂ + HCl → Mg (OH)Cl + H₂O 
4. Double salts: The salts formed by the union of two simple salts which dissolve in water and crystallise.
   Example: Potash alum: K₂SO₄. Al₂ (SO₄)₃. 24H₂O 
5. Mixed salts: Mixed salts are those salts which contain more than one basic or acidic radical. 
   Example: Sodium potassium carbonate NaKCO₃
6. Complex salts: Complex salts are those salts which on dissociation give one simple ion and one complex ion.
   Example: Na [Ag (CN)₂] ⇌ Na⁺ + [Ag (CN)₂]^-

Preparation of Soluble Salts

Preparation of Insoluble Salts

1. By direct combination
   Reaction: Pb + S → PbS
2. By combination of an acidic oxide with a basic oxide
   Reaction:  SO₂ + CaO → CaSO₃
3. Double decomposition 
   Reactions: BaCl₂ + H₂SO₄ → BaSO₄ + 2HCl 

Laboratory Preparation of some Normal and Acid Salts

1. Iron (III) chloride or anhydrous ferric chloride
   It is prepared by passing dry chlorine gas over heated iron.
   Fe + Cl₂ → FeCl₃
2. Copper (II) sulphate
   It is prepared by the reaction of copper oxide, copper hydroxides or copper carbonates with dilute sulphuric acid.
   CuO + H₂SO₄   → CuSO₄   + H₂O 
   Cu (OH)₂   +  H₂SO₄  → CuSO₄   +  2H₂O 
   CuCO₃   + H₂SO₄ →  CuSO₄   + H₂O 
   CuSO₄ + 5H₂O → CuSO₄.5H₂O 
3. Zinc sulphate and iron (II) sulphate
   It is prepared by the reaction of metals with dilute sulphuric acid.
   Zn + H₂SO₄ → ZnSO₄ + H₂O
   ZnSO₄ + 7H₂O → FeSO₄.7H₂O
4. Lead chloride
   It is prepared by adding either dilute hydrochloric acid or sodium chloride solution to a solution of lead nitrate.
   Pb (NO₃)₂ + 2HCl → PbCl₂ + 2HNO₃
5. Calcium carbonate
   It is prepared by adding sodium carbonate solution to a hot solution of calcium chloride.
   CaCl₂ + Na₂CO₃ → CaCO₃ + 2NaCl
6. Sodium bicarbonate
   It is prepared by passing excess of carbon dioxide gas through a saturated solution of sodium carbonate.
   Na₂CO₃ + CO₂ + H₂O → 2 NaHCO₃
7. Neutralisation
   It is the process by which H+ ions of an acid react completely with the [OH^-] ions of a base to give salt and water only.
   Example: HCl (Acid) + NaOH (Base) → NaCl (Salt) + H₂O (water)

Water of Crystallisation

It is the amount of water molecules which enter into loose chemical combination with one molecule of the substance on crystallisation from its aqueous solution.

Hydrated Salt

The salts which contain a definite number of water molecules as water of crystallisation are called hydrated salts.
Examples: Na₂CO₃.10H₂O (washing soda), CuSO₄.5H₂O (blue vitriol)

Anhydrous Salt

A salt which does not contain any water of crystallisation is called an anhydrous salt.
Examples: NaCl, NaNO₃, Pb(NO₃)₂

Deliquescence

Water soluble salts which on exposure to the atmosphere absorb moisture from the atmosphere, dissolve in the same and change into a solution. The phenomenon is called deliquescence and the salts deliquescent.
Examples: CaCl₂, MgCl₂, ZnCl₂

Efflorescence

Crystalline hydrated salts which on exposure to the atmosphere lose their water of crystallisation partly or completely and change into a powder. This phenomenon is called efflorescent and the salts efflorescent. 
Examples: CuSO₄.5H₂O, MgSO₄.7H₂O, Na₂CO₃.10H₂O