CBSE Class 10 Science Acids, Bases and Salts Last Minute Revision

CBSE Class 10 Science Acids, Bases and Salts: Acids, bases, and salts are some of the most important chemical substances we study in Class 10 Science. They play a major role not only in laboratory reactions but also in our everyday lives. 

From the food we eat to the medicines we use and the cleaning products at home, these compounds are everywhere. Acids release H⁺ ions and usually taste sour, while bases release OH⁻ ions and feel bitter or soapy. 

When acids and bases react, they form salts and water. Understanding their properties, reactions, and the pH scale helps build a strong foundation in chemistry.

CBSE Class 10 Science Acids, Bases and Salts

Revising CBSE Class 10 Science Acids, Bases and Salts is important for scoring well in exams. 

This chapter explains the properties, reactions, and uses of acids, bases, and common salts along with the pH scale and indicators. To quickly understand all key concepts and important points, check below for a clear and simple revision summary.

Introduction to Acids

Acids are substances showing distinct properties:

• They are sour in taste.

• They release H⁺ ions (or hydronium ions, H₃O⁺) in an aqueous solution. Substances releasing H⁺ ions in water are defined as acids.

• Most acids are corrosive, causing burning sensations and damaging materials.

Classification of Acids

Acids are classified by their strength and concentration.

1. Based on Strength (Degree of Dissociation):

• Strong Acids: These acids completely dissociate in water, producing a high concentration of H⁺ ions.

• Examples: Hydrochloric acid (HCl), Sulfuric acid (H₂SO₄).

• Weak Acids: These acids partially dissociate in water, releasing fewer H⁺ ions.

• Examples: Acetic acid (CH₃COOH), Oxalic acid.

2. Based on Concentration (Amount of Water):

• Concentrated Acids: Solutions with a large amount of acid and a small amount of water.

• Dilute Acids: Solutions with a small amount of acid and a large amount of water.

Dilution of Acids

Dilution is mixing acid with water. This process is highly exothermic, releasing significant heat.

Correct Procedure for Dilution:

• Always add acid slowly to water, with continuous stirring.

• NEVER add water to a concentrated acid. This can generate intense heat, causing the acid to splash or the container to break, leading to severe accidents.

Introduction to Bases

Bases are substances with characteristic properties:

• They are bitter in taste.

• They feel soapy to the touch.

• They release hydroxide ions (OH⁻) in an aqueous solution. Substances releasing OH⁻ ions in water are defined as bases.

Classification of Bases

• Strong Bases: Completely dissociate in water to produce a high concentration of OH⁻ ions.

• Examples: Sodium hydroxide (NaOH), Potassium hydroxide (KOH).

• Weak Bases: Partially dissociate in water.

• Example: Ammonium hydroxide (NH₄OH).

• Alkalis: Bases that are soluble in water are called alkalis.

Indicators

Indicators are substances that signal whether a solution is acidic or basic, usually by changing color or smell.

1. Natural Indicators

These are obtained from natural sources.

2. Synthetic Indicators

These are man-made indicators.

3. Olfactory Indicators

These indicators identify acidic or basic solutions by a change in their smell (odor).

Chemical Properties of Acids and Bases

1. Reaction with Metals

• Acids: React with metals to produce a salt and hydrogen gas (H₂).

• Acid + Metal → Salt + Hydrogen Gas

• Example: 2HCl (aq) + Zn (s) → ZnCl₂ (aq) + H₂ (g)

• Bases: Also react with certain metals to produce a salt and hydrogen gas.

• Base + Metal → Salt + Hydrogen Gas

• Example: 2NaOH (aq) + Zn (s) → Na₂ZnO₂ (s) [Sodium Zincate] + H₂ (g)

Test for Hydrogen Gas: Hydrogen gas burns with a characteristic pop sound when a burning matchstick is brought near it.

2. Reaction with Metal Carbonates and Metal Hydrogen Carbonates

• Acids: React with both metal carbonates and metal hydrogen carbonates, forming a salt, carbon dioxide (CO₂), and water (H₂O).

• Acid + Metal Carbonate → Salt + CO₂ + H₂O

• Acid + Metal Hydrogen Carbonate → Salt + CO₂ + H₂O

• Examples: 2HCl + Na₂CO₃ → 2NaCl + CO₂ + H₂O and HCl + NaHCO₃ → NaCl + CO₂ + H₂O

• Bases: Do not react with metal carbonates or metal hydrogen carbonates.

Test for Carbon Dioxide (Lime Water Test):

• CO₂ gas turns lime water (Ca(OH)₂) milky due to the formation of insoluble calcium carbonate (CaCO₃).

• Ca(OH)₂ (aq) + CO₂ (g) → CaCO₃ (s) + H₂O (l)

• Passing excess CO₂ makes the milkiness disappear as CaCO₃ converts to soluble calcium bicarbonate (Ca(HCO₃)₂).

• CaCO₃ (s) + H₂O (l) + CO₂ (g) [excess] → Ca(HCO₃)₂ (aq) [Colorless]

3. Reaction of Acids and Bases with Each Other (Neutralization)

• Acids and bases react to form a salt and water in a neutralization reaction.

• Acid + Base → Salt + Water

4. Reaction of Metallic Oxides with Acids

• Metallic oxides are basic in nature.

• They react with acids in a neutralization reaction to form a salt and water.

• Metallic Oxide (Base) + Acid → Salt + Water

• Example: CuO (s) + 2HCl (aq) → CuCl₂ (aq) + H₂O (l)

5. Reaction of Non-Metallic Oxides with Bases

• Non-metallic oxides are acidic in nature.

• They react with bases to form a salt and water.

• Non-Metallic Oxide (Acid) + Base → Salt + Water

• Example: CO₂ (g) + Ca(OH)₂ (aq) → CaCO₃ (s) + H₂O (l)

The pH Scale and Universal Indicators

• Universal Indicator: A mixture of indicators showing different colors at various hydrogen ion concentrations. It measures pH.

• pH Scale: A scale measuring hydrogen ion concentration in a solution, ranging from 0 to 14.

• pH < 7: Acidic solution

• pH = 7: Neutral solution

• pH > 7: Basic (or alkaline) solution

Important Salts

A salt is an ionic compound formed from the neutralization of an acid and a base.

1. Sodium Hydroxide (NaOH)

• Preparation (Chlor-alkali process): Produced by the electrolysis of aqueous NaCl (brine).

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

• Chlorine gas (Cl₂) is produced at the anode.

• Hydrogen gas (H₂) is produced at the cathode.

• Uses: Manufacturing soaps, detergents, and paper.

2. Bleaching Powder (CaOCl₂)

• Preparation: Action of chlorine gas on dry slaked lime (Ca(OH)₂).

• Ca(OH)₂ + Cl₂ → CaOCl₂ + H₂O

• Uses:

• Bleaching cotton and linen in textiles.

• Bleaching wood pulp in paper factories.

• As an oxidizing agent.

• Disinfecting drinking water to make it germ-free.

3. Baking Soda (Sodium Hydrogen Carbonate, NaHCO₃)

• Preparation: Uses NaCl as a raw material.

• NaCl + H₂O + CO₂ + NH₃ → NH₄Cl + NaHCO₃

• Action of Heat (During Cooking): When heated, baking soda decomposes to form sodium carbonate, water, and carbon dioxide. The CO₂ gas produced causes bread and cakes to rise, making them soft and spongy.

• 2NaHCO₃ (s) --(Heat)→ Na₂CO₃ (s) + H₂O (l) + CO₂ (g)

• Other Uses:

• Ingredient in antacids to neutralize stomach acid.

• Used in soda-acid fire extinguishers.

• Used to make baking powder (baking soda with a mild edible acid like tartaric acid).

4. Washing Soda (Sodium Carbonate Decahydrate, Na₂CO₃·10H₂O)

• Preparation: Recrystallization of sodium carbonate with water.

• Na₂CO₃ (s) + 10H₂O (l) → Na₂CO₃·10H₂O (s)

• Uses:

• In glass, soap, and paper industries.

• Manufacturing sodium compounds like Borax.

• As a cleaning agent for domestic purposes.

• For removing the permanent hardness of water.

5. Plaster of Paris (POP) (Calcium Sulfate Hemihydrate, CaSO₄·½H₂O)

• Preparation: Produced by heating Gypsum (CaSO₄·2H₂O) at 373 K (100°C). Gypsum loses water molecules.

• CaSO₄·2H₂O --(Heat at 373 K)→ CaSO₄·½H₂O + 1½H₂O

• Properties: A white powder that, when mixed with water, changes back into gypsum, setting into a hard solid mass.

• CaSO₄·½H₂O + 1½H₂O → CaSO₄·2H₂O (Gypsum)

• Uses:

• Supporting fractured bones.

• Making toys and decorative materials.

• Creating smooth surfaces and false ceilings.