Acids, Bases and Salts: Complete Guide for Class 10 CBSE Notes

Introduction to Acids, Bases and Salts

Acids, bases, and salts are fundamental chemical substances that we encounter daily. From the sour taste of lemon juice (citric acid) to the slippery feel of soap (basic), these compounds play crucial roles in our everyday life. Understanding their properties, reactions, and applications is essential for Class 10 CBSE chemistry.

The study of acids and bases has evolved through various definitions, starting with simple taste-based identification to sophisticated electronic theories. This comprehensive guide covers all aspects of acids, bases, and salts as per the CBSE Class 10 curriculum.

Theories of Acids and Bases

1. Arrhenius Concept of Acids and Bases

Swedish chemist Svante Arrhenius proposed this theory in 1887, focusing on ionization in water.

Arrhenius Acid:

• A substance containing hydrogen that releases H⁺ ions when dissolved in water
• Example: HCl(aq) → H⁺(aq) + Cl⁻(aq)

Arrhenius Base:

• A substance containing hydroxyl groups that releases OH⁻ ions when dissolved in water
• Example: NaOH(aq) → Na⁺(aq) + OH⁻(aq)

Strong vs. Weak Acids/Bases:

• Strong acids (HCl, H₂SO₄, HNO₃) ionize almost completely in water
• Weak acids (CH₃COOH, H₃PO₄) ionize partially in water
• Strong bases (NaOH, KOH) ionize completely in water
• Weak bases (NH₄OH, Mg(OH)₂) ionize partially in water

Neutralization according to Arrhenius: H⁺(aq) + OH⁻(aq) → H₂O(l)

2. Bronsted-Lowry Concept of Acids and Bases

Proposed independently by J.N. Bronsted and J.M. Lowry in 1923, this theory provides a more general definition.

• Bronsted-Lowry Acid: A proton (H⁺) donor
• Bronsted-Lowry Base: A proton (H⁺) acceptor

Conjugate Acid-Base Pairs:

When an acid donates a proton, it forms its conjugate base. When a base accepts a proton, it forms its conjugate acid.

Example:

HCl(aq) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)(acid) (base) (conjugate (conjugate acid) base)

Important Examples:

1. HNO₃ + H₂O ⇌ H₃O⁺ + NO₃⁻
2. NH₄⁺ + H₂O ⇌ H₃O⁺ + NH₃
3. H₂O + NH₃ ⇌ NH₄⁺ + OH⁻

Advantage: This theory explains acid-base behavior in non-aqueous solutions and includes ionic species as acids or bases.

3. Lewis Concept of Acids and Bases

G.N. Lewis proposed the most fundamental definition based on electron pairs.

• Lewis Acid: An electron pair acceptor (has empty orbitals)
• Lewis Base: An electron pair donor (has lone pairs of electrons)
• General Reaction: A⁺ + :B → B:A
• Types of Lewis Acids:

1. Molecules with incomplete octet: BF₃, AlCl₃, FeCl₃
2. Simple cations: Na⁺, K⁺, H⁺, Ag⁺
3. Molecules that can expand their octet: SiF₄, SnCl₄

Types of Lewis Bases:

1. Neutral molecules with lone pairs: NH₃, H₂O
2. Anions with negative charge: F⁻, Cl⁻, OH⁻
3. Molecules with π bonds: C₂H₄, C₂H₂

Examples:

1. H⁺ + NH₃ → NH₄⁺
2. BF₃ + NH₃ → H₃N→BF₃
3. AlCl₃ + Cl⁻ → [AlCl₄]⁻

Differences Between Arrhenius, Bronsted-Lowry and Lewis Definitions

Progressive Evolution:

• Arrhenius → Focused on water solutions
• Bronsted-Lowry → Extended to proton transfer reactions
• Lewis → Universal application based on electron pairs

Classification of Acids

A. Based on Source

1. Mineral/Inorganic Acids:

• Obtained from minerals or rocks
• Examples: HCl, H₂SO₄, HNO₃, H₂CO₃, H₃PO₄

2. Organic Acids:

• Obtained from plants and animals
• Examples: HCOOH (formic acid), CH₃COOH (acetic acid), C₆H₅COOH (benzoic acid), C₆H₈O₇ (citric acid)

B. Based on Basicity (Number of H⁺ ions produced)

1. Monobasic Acids:

• Produce one H⁺ ion per molecule
• Examples: HCl, HBr, HNO₃

2. Dibasic Acids:

• Produce two H⁺ ions per molecule
• Examples: H₂SO₄, H₂CO₃

3. Tribasic Acids:

• Produce three H⁺ ions per molecule
• Examples: H₃PO₄, C₆H₈O₇ (citric acid)

C. Based on Strength

1. Strong Acids:

• Undergo complete ionization
• Examples: HCl, H₂SO₄, HNO₃

2. Weak Acids:

• Undergo incomplete ionization
• Examples: CH₃COOH, HCOOH, H₂CO₃

D. Based on Concentration

1. Concentrated Acids: Contain less water, more acid

2. Dilute Acids: Contain large volume of water, less acid

Related Helpful Resources

Classification of Bases

A. Based on Acidity (Number of OH⁻ ions produced)

1. Monoacidic Bases:

• Produce one OH⁻ ion per molecule
• Examples: NaOH, KOH

2. Diacidic Bases:

• Produce two OH⁻ ions per molecule
• Examples: Ca(OH)₂, Mg(OH)₂

3. Triacidic Bases:

• Produce three OH⁻ ions per molecule
• Examples: Al(OH)₃, Fe(OH)₃

B. Based on Strength

1. Strong Bases:

• Undergo complete ionization
• Examples: NaOH, KOH, Ba(OH)₂

2. Weak Bases:

• Undergo incomplete ionization
• Examples: Ca(OH)₂, Mg(OH)₂, NH₄OH

C. Based on Concentration

1. Concentrated Bases: Contain less water

2. Dilute Bases: Contain excess water

Important Note: Bases that are soluble in water are called alkalis. All alkalis are bases, but all bases are not alkalis.

Common Examples of Acids, Bases and Salts in Daily Life

Acids in Daily Life

Bases in Daily Life

Salts in Daily Life

How to Test Substances Using Litmus, pH and Indicators

1. Litmus Test

Litmus is a purple dye extracted from lichen (a plant). It is available as:

• Blue litmus paper
• Red litmus paper
• Litmus solution (purple)

Testing Procedure:

Practical Application:

1. Dip litmus paper in the test solution
2. Observe color change
3. Record the nature of the substance

2. Phenolphthalein Indicator

Properties:

• Colorless in neutral or acidic solutions
• Pink in basic solutions

Testing Method:

• Add 2-3 drops to the test solution
• Colorless → Acidic or Neutral
• Pink → Basic

3. Methyl Orange Indicator

Properties:

• Orange in neutral medium
• Red in acidic medium
• Yellow in basic medium

4. Universal Indicator and pH Paper

Universal indicator is a mixture of several indicators that shows different colors at different pH values.

pH Paper Testing:

1. Dip pH paper in the test solution
2. Compare the color with the pH color chart
3. Determine the pH value (0-14)

pH Color Chart:

• 0-3: Red (Strongly acidic)
• 4-6: Orange-Yellow (Weakly acidic)
• 7: Green (Neutral)
• 8-10: Blue (Weakly basic)
• 11-14: Purple-Violet (Strongly basic)

5. Olfactory Indicators

Some substances change their odor in acidic or basic medium.

Examples:

Onion Extract:

• Natural smell in neutral/basic medium
• Smell disappears in acidic medium

Vanilla Essence:

• Characteristic smell in acidic medium
• Smell changes in basic medium

Clove Oil:

• Distinct smell in acidic medium
• Smell changes in basic medium

Testing Procedure:

1. Prepare onion-scented cloth strips
2. Add acid/base to separate strips
3. Rinse with water and smell
4. Compare odors to identify acid/base

The pH Scale

Understanding pH

The pH scale was introduced by Danish biochemist S. Sorensen to measure the strength of acids and bases. pH stands for "power of hydrogen" (potenz in German).

pH Scale Characteristics:

• Range: 0 to 14
• pH < 7: Acidic solution
• pH = 7: Neutral solution
• pH > 7: Basic/Alkaline solution

Mathematical Definition: pH = -log[H⁺] = -log[H₃O⁺]

pH Values of Common Substances

Relationship Between pH and H⁺ Concentration

Important Relationships:

• pH + pOH = 14
• pOH = -log[OH⁻]
• Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ (at 25°C)

Importance of pH in Daily Life

1. Human Body:

• Blood pH: 7.35-7.45 (slightly basic)
• Stomach pH: 1.5-3.0 (for digestion)
• Saliva pH: 6.5-7.5

2. Agriculture:

• Most plants prefer pH 6.0-7.5
• Acidic soil: Add lime Ca(OH)₂
• Basic soil: Add organic matter

3. Aquatic Life:

• Fish survive best at pH 6.5-8.5
• pH changes affect aquatic ecosystems

4. Industrial Applications:

• pH control in manufacturing
• Water treatment processes
• Food processing

Chemical Properties of Acids and Bases

1. Reaction of Acids with Metals

General Reaction: Metal + Acid → Metal Salt + Hydrogen gas

Examples:

1. Zinc with hydrochloric acid: Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)↑
2. Magnesium with sulphuric acid: Mg(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂(g)↑
3. Sodium with hydrochloric acid: 2Na(s) + 2HCl(aq) → 2NaCl(aq) + H₂(g)↑

Test for Hydrogen Gas:

• Brings a burning candle near the gas
• Hydrogen burns with a "pop" sound
• Forms soap bubbles that burst when ignited

Important Note: Copper, mercury, and silver do not react with dilute HCl or H₂SO₄.

2. Reaction of Bases with Metals

Some metals react with strong alkalis to produce hydrogen gas.

Examples:

1. Zinc with sodium hydroxide: Zn(s) + 2NaOH(aq) → Na₂ZnO₂(aq) + H₂(g)↑ (sodium zincate)
2. Aluminum with sodium hydroxide: 2Al(s) + 2NaOH(aq) + 2H₂O(l) → 2NaAlO₂(aq) + 3H₂(g)↑ (sodium meta-aluminate)
3. Tin with sodium hydroxide: Sn(s) + 2NaOH(aq) → Na₂SnO₂(aq) + H₂(g)↑ (sodium stannite)

3. Reaction of Acids with Metal Oxides

General Reaction: Metal Oxide + Acid → Salt + Water

Metal oxides are basic in nature and neutralize acids.

Examples:

1. Copper(II) oxide with hydrochloric acid: CuO(s) + 2HCl(aq) → CuCl₂(aq) + H₂O(l) (black) → (bluish-green solution)
2. Calcium oxide with hydrochloric acid: CaO(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l)
3. Magnesium oxide with sulphuric acid: MgO(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂O(l)

4. Reaction of Bases with Non-Metal Oxides

General Reaction: Base + Non-metal Oxide → Salt + Water

Non-metal oxides are acidic in nature.

Examples:

1. Sodium hydroxide with carbon dioxide: 2NaOH(aq) + CO₂(g) → Na₂CO₃(aq) + H₂O(l) (sodium carbonate)
2. Calcium hydroxide with carbon dioxide: Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l) (lime water) → (milky precipitate) Excess CO₂: CaCO₃(s) + CO₂(g) + H₂O(l) → Ca(HCO₃)₂(aq) (milkiness disappears) (soluble)
3. Potassium hydroxide with carbon dioxide: 2KOH(aq) + CO₂(g) → K₂CO₃(aq) + H₂O(l)

Note: CO₂, SO₂, SO₃ are non-metallic oxides and are acidic in nature.

5. Reaction of Acids with Metal Carbonates and Bicarbonates

General Reactions:

• Metal Carbonate + Acid → Salt + Water + Carbon Dioxide
• Metal Bicarbonate + Acid → Salt + Water + Carbon Dioxide

Examples:

1. Calcium carbonate with hydrochloric acid: CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)↑
2. Sodium carbonate with sulphuric acid: Na₂CO₃(s) + H₂SO₄(aq) → Na₂SO₄(aq) + H₂O(l) + CO₂(g)↑
3. Sodium bicarbonate with sulphuric acid: 2NaHCO₃(s) + H₂SO₄(aq) → Na₂SO₄(aq) + 2H₂O(l) + 2CO₂(g)↑

Test for Carbon Dioxide:

• Pass the gas through lime water
• Lime water turns milky
• Ca(OH)₂ + CO₂ → CaCO₃↓ + H₂O

Interesting Facts:

• Limestone (CaCO₃), chalk, marble are different forms of calcium carbonate
• Egg shells contain calcium carbonate

Neutralization Reaction: Equation Examples and Applications

Definition and Concept

Neutralization is the chemical reaction between an acid and a base to form salt and water.

General Equation: Acid + Base → Salt + Water

Ionic Equation: H⁺(aq) + OH⁻(aq) → H₂O(l)

The reaction is called neutralization because the properties of both acid and base are neutralized.

Neutralization Reaction Examples

1. Hydrochloric acid with sodium hydroxide: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l) (acid) + (base) → (salt) + (water)

2. Sulphuric acid with potassium hydroxide: H₂SO₄(aq) + 2KOH(aq) → K₂SO₄(aq) + 2H₂O(l)

3. Nitric acid with sodium hydroxide: HNO₃(aq) + NaOH(aq) → NaNO₃(aq) + H₂O(l)

4. Acetic acid with ammonium hydroxide: CH₃COOH(aq) + NH₄OH(aq) → CH₃COONH₄(aq) + H₂O(l)

5. Phosphoric acid with potassium hydroxide: H₃PO₄(aq) + 3KOH(aq) → K₃PO₄(aq) + 3H₂O(l)

Practical Applications of Neutralization

1. Medical Applications:

Treating Acidity:

• Problem: Excess HCl in stomach causes acidity
• Solution: Antacids (bases) neutralize excess acid
• Antacids contain: NaHCO₃, Mg(OH)₂, Al(OH)₃
• Reaction: HCl + NaHCO₃ → NaCl + H₂O + CO₂

Treating Indigestion:

• Milk of magnesia: Mg(OH)₂(aq) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l)

2. Insect Bites and Stings:

Bee/Ant Stings (contain formic acid):

• Apply base like baking soda paste
• NaHCO₃ neutralizes formic acid
• Soap (contains NaOH) can also be used

Wasp Stings (contain alkali):

• Apply weak acid like vinegar
• Acetic acid neutralizes the alkaline sting

3. Agricultural Applications:

Acidic Soil Treatment:

• Problem: Excess acid in soil harms crops
• Solution: Add slaked lime Ca(OH)₂ or limestone CaCO₃
• Neutralization raises soil pH to optimal range (6-7)

Basic Soil Treatment:

• Problem: Excess alkali in soil
• Solution: Add organic matter or gypsum (CaSO₄)

4. Industrial Applications:

Factory Waste Treatment:

• Acidic waste: Neutralized with lime Ca(OH)₂
• Basic waste: Neutralized with dilute acids
• Prevents environmental damage

Water Treatment:

• Adjust pH of water for safe consumption
• Use lime for acidic water
• Use alum for basic water

5. Fire Extinguishers:

• Soda-acid fire extinguisher
• Contains NaHCO₃ and H₂SO₄
• Reaction produces CO₂: 2NaHCO₃ + H₂SO₄ → Na₂SO₄ + 2H₂O + 2CO₂

How to Prepare and Handle Dilute Solutions Safely

Dilution of Acids and Bases

Definition: Dilution is the process of decreasing the concentration of an acid or base by adding water.

Exothermic Nature: The dilution of concentrated acids and bases is highly exothermic (releases heat).

Safety Rules for Dilution

Golden Rule of Dilution

CORRECT METHOD:"Always add acid to water, never water to acid"

Procedure:

1. Take water in a large beaker
2. Add concentrated acid drop by drop with constant stirring
3. Use a glass rod for stirring
4. Heat dissipates gradually and safely

WRONG METHOD:Never add water to concentrated acid

Dangers:

• Massive heat generation
• Mixture may splash out causing burns
• Glass container may crack
• Acid vapors cause pollution and respiratory problems
• Can cause severe burns to skin and eyes

Practical Dilution Procedure

Step-by-Step Guide:

For Diluting Sulphuric Acid:

1. Preparation:
   • Wear safety goggles and gloves
   • Use a large heat-resistant beaker
   • Keep cold water nearby for emergency
2. Dilution Process:
   • Pour water into the beaker (fill 2/3)
   • Add concentrated H₂SO₄ drop by drop
   • Stir continuously with a glass rod
   • Add acid slowly along the side of the beaker
   • Allow cooling between additions
3. Observation:
   • Beaker becomes warm (exothermic reaction)
   • Never touch the beaker directly
   • Let it cool before handling

For Diluting Sodium Hydroxide:

1. Take water in a beaker
2. Add NaOH pellets or solution slowly
3. Stir continuously
4. Heat is released, so handle carefully
5. NaOH is deliquescent (absorbs moisture from air)

Safety Equipment and Precautions

Personal Protective Equipment (PPE):

• Safety goggles (mandatory)
• Lab coat or apron
• Chemical-resistant gloves
• Closed-toe shoes
• Face shield for concentrated acids

Laboratory Precautions:

1. Work in well-ventilated area or fume hood
2. Keep fire extinguisher nearby
3. Have neutralizing agents ready:
   • Baking soda for acid spills
   • Dilute acid for base spills
4. Never smell chemicals directly
5. Label all containers clearly
6. Store acids and bases separately
7. Keep containers tightly closed

First Aid for Chemical Burns

Acid Burns:

1. Immediately wash with plenty of water (15-20 minutes)
2. Apply baking soda paste (weak base)
3. Seek medical attention
4. Do not apply oil or ointments

Base Burns:

1. Wash thoroughly with water
2. Apply dilute vinegar or lemon juice (weak acid)
3. Seek medical attention
4. Do not rub the affected area

Eye Contact:

1. Rinse eyes immediately with water for 15 minutes
2. Use eye wash station if available
3. Keep eyes open while rinsing
4. Seek immediate medical help

Storage Guidelines

Acid Storage:

• Store in cool, dry place
• Keep away from bases
• Use corrosion-resistant shelves
• Store below eye level
• Keep away from metals

Base Storage:

• Store in airtight containers
• Protect from air and moisture
• Keep away from acids
• Use plastic or glass containers

Warning Signs:

• Display appropriate hazard symbols
• Corrosive warning signs
• Emergency contact information
• Safety data sheets (SDS) accessible

Salts: Formation, Properties and Applications

What are Salts?

Definition: Salts are ionic compounds consisting of:

• Cation (positive ion): From base
• Anion (negative ion): From acid

General Formation: Acid + Base → Salt + Water

Classification of Salts

1. Based on pH

a) Neutral Salts (pH = 7):

• Formed from strong acid + strong base
• Example: NaCl (HCl + NaOH)
• No effect on litmus paper

b) Acidic Salts (pH < 7):

• Formed from strong acid + weak base
• Example: NH₄Cl (HCl + NH₄OH)
• Turn blue litmus red

c) Basic Salts (pH > 7):

• Formed from weak acid + strong base
• Example: Na₂CO₃ (H₂CO₃ + NaOH)
• Turn red litmus blue

d) Amphoteric Salts (pH ≈ 7):

• Formed from weak acid + weak base
• Example: CH₃COONH₄
• Can act as acid or base

2. Based on Composition

a) Normal Salts:

• Complete neutralization
• No H⁺ or OH⁻ ions
• Example: NaCl, K₂SO₄

b) Acid Salts:

• Incomplete neutralization of polybasic acid
• Contain replaceable H⁺
• Example: NaHSO₄, NaHCO₃

c) Basic Salts:

• Incomplete neutralization of polyacidic base
• Contain OH⁻ ions
• Example: Mg(OH)Cl

d) Mixed Salts:

• More than one cation or anion
• Example: NaKSO₄, CaOCl₂

e) Double Salts:

• Two salts crystallized together
• Example: Mohr's salt FeSO₄.(NH₄)₂SO₄.6H₂O

Family of Salts

Salts with the same cation or anion belong to the same family.

Sodium Salt Family (Na⁺):

• NaCl, Na₂SO₄, NaNO₃, Na₂CO₃

Chloride Family (Cl⁻):

• NaCl, KCl, NH₄Cl, CaCl₂

Sulphate Family (SO₄²⁻):

• Na₂SO₄, K₂SO₄, CuSO₄, MgSO₄

Methods of Preparation of Salts

1. Neutralization: Acid + Base → Salt + Water Example: HCl + NaOH → NaCl + H₂O

2. Metal + Acid: Metal + Acid → Salt + Hydrogen Example: Zn + H₂SO₄ → ZnSO₄ + H₂↑

3. Metal Oxide + Acid: Metal Oxide + Acid → Salt + Water Example: CuO + 2HCl → CuCl₂ + H₂O

4. Metal Carbonate + Acid: Metal Carbonate + Acid → Salt + Water + CO₂ Example: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂↑

5. Metal + Alkali: Metal + Base → Salt + Hydrogen Example: Zn + 2NaOH → Na₂ZnO₂ + H₂↑

Important Salts and Their Uses

1. Sodium Chloride (NaCl) - Common Salt

Sources:

• Sea water (main source)
• Rock salt deposits
• Salt lakes

Properties:

• White crystalline solid
• Neutral salt (pH = 7)
• Highly soluble in water
• Essential for life

Uses:

1. Food seasoning and preservation
2. Raw material for chemicals:
   • Sodium hydroxide (NaOH)
   • Washing soda (Na₂CO₃)
   • Baking soda (NaHCO₃)
   • Bleaching powder (CaOCl₂)
3. De-icing roads in winter
4. Manufacturing soaps and detergents

Preparation from Sea Water:

1. Evaporation of sea water in shallow pools
2. Crystallization of NaCl
3. Separation and purification

2. Sodium Hydroxide (NaOH) - Caustic Soda

Preparation: Chlor-Alkali Process

2NaCl(aq) + 2H₂O(l) → 2NaOH(aq) + Cl₂(g)↑ + H₂(g)↑ electrolysis

Products:

• At cathode: Hydrogen gas (H₂)
• At anode: Chlorine gas (Cl₂)
• Near cathode: Sodium hydroxide solution

Uses:

1. Soap and detergent manufacturing
2. Paper manufacturing
3. Artificial fibers production
4. Degreasing metals
5. Petroleum refining

Products from Chlor-Alkali Process:

Hydrogen (at cathode):

• Fuel
• Margarine production
• Ammonia for fertilizers

Chlorine (at anode):

• Bleaching powder
• Hydrochloric acid
• Household bleaches
• PVC plastics
• Pesticides
• Medicines

Sodium Hydroxide:

• All the uses mentioned above

3. Bleaching Powder (CaOCl₂)

Preparation:

Ca(OH)₂(s) + Cl₂(g) → CaOCl₂(s) + H₂O(l)(slaked lime) (chlorine) (bleaching powder)

Properties:

1. Yellowish-white powder
2. Strong smell of chlorine
3. Soluble in water (leaves lime residue)
4. Decomposes in air to release chlorine

Chemical Reactions:

With excess dilute acid:

CaOCl₂ + H₂SO₄ → CaSO₄ + H₂O + Cl₂↑CaOCl₂ + 2HCl → CaCl₂ + H₂O + Cl₂↑

With CO₂ (decomposition in air):

CaOCl₂ + CO₂ → CaCO₃ + Cl₂↑

Uses:

1. Bleaching:
   • Cotton and linen textiles
   • Wood pulp (paper industry)
   • Not suitable for silk, wool (damages them)
2. Disinfectant:
   • Water sterilization (drinking water, swimming pools)
   • Sanitation
3. Oxidizing agent:
   • Chemical manufacturing
4. Other uses:
   • Making wool unshrinkable
   • Treating wastewater

4. Baking Soda (Sodium Bicarbonate - NaHCO₃)

Preparation:

NaCl + H₂O + CO₂ + NH₃ → NH₄Cl + NaHCO₃(Solvay process)

Properties:

• White crystalline powder
• Mildly alkaline (pH ≈ 8.5)
• Decomposes on heating:

2NaHCO₃ --heat--> Na₂CO₃ + H₂O + CO₂↑

Uses:

1. Baking:
   • Baking powder ingredient
   • Releases CO₂, makes cake/bread fluffy
   • NaHCO₃ + Acid (tartaric acid) → CO₂ + Salt + Water
2. Antacid:
   • Neutralizes stomach acid
   • NaHCO₃ + HCl → NaCl + H₂O + CO₂
3. Fire Extinguisher:
   • Soda-acid type
   • Generates CO₂ which extinguishes fire
4. Cleaning agent:
   • Removes stains
   • Neutralizes odors
5. Toothpaste:
   • Mild abrasive
   • Neutralizes mouth acids

5. Washing Soda (Sodium Carbonate - Na₂CO₃.10H₂O)

Preparation:

2NaHCO₃ --heat--> Na₂CO₃ + H₂O + CO₂↑Na₂CO₃ + 10H₂O → Na₂CO₃.10H₂O (washing soda crystals)

Recrystallization of sodium carbonate: One formula unit of Na₂CO₃ combines with 10 water molecules (water of crystallization).

Properties:

• Translucent white crystals
• Strongly alkaline (pH ≈ 11)
• Loses water on exposure to air (efflorescence)
• Soluble in water

Uses:

1. Cleaning agent:
   • Washing clothes
   • Removes permanent hardness of water
2. Glass manufacturing:
   • Important raw material
3. Paper industry:
   • Pulp processing
4. Soap and detergent manufacturing
5. Water softening:
   • Removes Ca²⁺ and Mg²⁺ ions
   • Na₂CO₃ + CaSO₄ → CaCO₃↓ + Na₂SO₄

6. Plaster of Paris (CaSO₄.½H₂O)

Preparation:

CaSO₄.2H₂O --heat 373K--> CaSO₄.½H₂O + 1½H₂O(gypsum) (plaster of Paris)

Properties:

• White powder
• Sets into hard mass when mixed with water
• Setting reaction:

CaSO₄.½H₂O + 1½H₂O → CaSO₄.2H₂O (gypsum)

Uses:

1. Medical:
   • Plaster casts for fractured bones
   • Dental moulds
2. Construction:
   • Decorative materials
   • Smooth surfaces on walls
3. Sculpture and artwork
4. Toy making and pottery

Chemical Formulas

Practice Questions and Concepts

Important Points to Remember

1. Acid-Base Indicators:

• Litmus: Red in acid, blue in base
• Phenolphthalein: Colorless in acid, pink in base
• Methyl orange: Red in acid, yellow in base

2. pH Memory Aid:

• pH < 7: Acidic (like vinegar, lemon)
• pH = 7: Neutral (pure water)
• pH > 7: Basic (soap, baking soda)

3. Strong vs. Weak:

• Strong acids: HCl, H₂SO₄, HNO₃
• Weak acids: CH₃COOH, H₂CO₃
• Strong bases: NaOH, KOH
• Weak bases: NH₄OH, Mg(OH)₂

4. Safety Rules:

• Always add acid to water, never reverse
• Wear safety equipment
• Work in ventilated areas

5. Salt Properties:

• Strong acid + Strong base = Neutral salt
• Strong acid + Weak base = Acidic salt
• Weak acid + Strong base = Basic salt

Common Misconceptions Clarified

Misconception 1: "All compounds containing hydrogen are acids" Reality: Only those that release H⁺ in water are acids. Ethanol (C₂H₅OH) and glucose (C₆H₁₂O₆) contain H but are not acids.

Misconception 2: "Dilution changes the nature of acid/base" Reality: Dilution only changes concentration and pH, not the acidic/basic nature.

Misconception 3: "pH can be any value" Reality: While pH scale is 0-14, some substances can have pH < 0 or > 14 in special cases.

Misconception 4: "All salts are neutral" Reality: Salts can be acidic, basic, or neutral depending on parent acid and base.

Misconception 5: "Dry HCl gas is acidic" Reality: Dry HCl gas doesn't show acidic properties. Only aqueous solution does.

Conclusion

Understanding acids, bases, and salts is fundamental to chemistry and has immense practical importance in daily life. From the food we eat to the medicines we take, from agriculture to industry, these concepts are everywhere.

1. Three major theories (Arrhenius, Bronsted-Lowry, Lewis) provide progressively broader definitions of acids and bases
2. Indicators and pH scale help identify and quantify acidity or basicity
3. Chemical reactions of acids and bases with metals, metal oxides, carbonates follow predictable patterns
4. Neutralization is the basis of many practical applications including medicine, agriculture, and industry
5. Safety precautions are essential when handling concentrated acids and bases
6. Salts formed from acid-base reactions have diverse properties and uses