Notes :: Chapter 2: Acids , Bases and Salts

Acids, Bases, and Indicators

1. Acids and Bases in Daily Life

• Sour taste in food is due to acids.
• Bitter taste in food is due to bases.
• Acids turn blue litmus red.
• Bases turn red litmus blue.

2. Remedy for Acidity (Indigestion)

Baking soda solution is used to neutralize acidity. Reason: Acidity is caused by excess acid in the stomach, and baking soda (a base) neutralizes it.

Concept: Acids and bases nullify each other’s effects.

3. Acid-Base Indicators

Natural Indicators:

• Litmus solution (extracted from lichen):
  • Acidic solution → Red
  • Basic solution → Blue
• Turmeric:
  • Basic solution → Reddish-brown
  • Returns to yellow when washed with water
• Red cabbage leaves, Hydrangea, Petunia, Geranium petals – Used to detect acids and bases.

Synthetic Indicators:

• Methyl orange
• Phenolphthalein

4. Litmus Solution

A purple dye obtained from lichen (Thallophyta plant). Used to test acidity and basicity. Neutral solution remains purple.

5. Key Concept

Acids and bases react with each other and neutralize their effects. Indicators help identify whether a substance is acidic or basic without tasting it.

Understanding the Chemical Properties of Acids and Bases

2.1.1 Acids and Bases in the Laboratory

Activity 2.1: Testing Acids and Bases with Indicators

Materials Collected:

Acids:

• Hydrochloric acid (HCl)
• Sulphuric acid (H₂SO₄)
• Nitric acid (HNO₃)
• Acetic acid (CH₃COOH)

Bases:

• Sodium hydroxide (NaOH)
• Calcium hydroxide (Ca(OH)₂)
• Potassium hydroxide (KOH)
• Magnesium hydroxide (Mg(OH)₂)
• Ammonium hydroxide (NH₄OH)

Procedure:

Place a drop of each solution on a watch glass. Add one drop of the given indicators (red litmus, blue litmus, phenolphthalein, and methyl orange) to each sample. Observe and record the colour changes in the table.

Observations with Indicators:

Sample Solution	Red Litmus	Blue Litmus	Phenolphthalein	Methyl Orange
Acids (HCl, H₂SO₄, HNO₃, CH₃COOH)	Turns red (no change)	Turns red	Colourless	Turns red
Bases (NaOH, Ca(OH)₂, KOH, Mg(OH)₂, NH₄OH)	Turns blue	No change (remains blue)	Turns pink	Turns yellow
Neutral (Water)	No change	No change	No change	No change

Conclusion:

• Acids turn blue litmus red, while bases turn red litmus blue.
• Phenolphthalein remains colourless in acids and turns pink in bases.
• Methyl orange turns red in acids and yellow in bases.

Activity 2.2: Olfactory Indicators

Olfactory Indicators: Substances whose odour changes in acidic or basic solutions.

Materials Used:

• Finely chopped onions
• Vanilla essence
• Clove oil
• Dilute HCl solution (acid)
• Dilute NaOH solution (base)

Procedure:

Onion Test:

• Place finely chopped onion in a plastic bag with clean cloth strips.
• Leave overnight in the fridge.
• Take out two strips, smell them, and note the odour.
• Add a few drops of dilute HCl to one strip and dilute NaOH to the other.
• Observe if there is any change in odour.
• Rinse both strips with water and check for odour changes again.

Vanilla and Clove Test:

• Add a few drops of vanilla essence to test tubes containing dilute HCl and dilute NaOH.
• Shake well and check if the odour changes.
• Repeat with clove oil and record observations.

Observations:

• Onion smell fades in basic (NaOH) solution but remains in acidic (HCl) solution.
• Vanilla essence loses its smell in basic solution but remains in acidic solution.
• Clove oil’s odour changes in basic solution but remains in acidic solution.

Conclusion:

Onion, vanilla, and clove can be used as olfactory indicators. Their odour changes in acidic and basic media.

Key Takeaways:

• Acids and bases can be identified using indicators (litmus, phenolphthalein, methyl orange).
• Olfactory indicators like onion, vanilla, and clove help detect acids and bases by changes in smell.
• Acids and bases react differently with different indicators, which helps in their identification.

How do Acids and Bases React with Metals?

Activity 2.3: Reaction of Acids with Metals

Procedure:

Set up the apparatus as shown in Figure 2.1. Take 5 mL of dilute sulphuric acid (H₂SO₄) in a test tube. Add a few pieces of zinc (Zn) granules to it. Observe the surface of the zinc granules. Pass the gas evolved through a soap solution. Bring a burning candle near a gas-filled bubble. Repeat the experiment with other acids:

• Hydrochloric acid (HCl)
• Nitric acid (HNO₃)
• Acetic acid (CH₃COOH)

Observations:

• Effervescence (bubbles) is observed on the surface of zinc granules.
• The bubbles in the soap solution indicate the release of a gas.
• When a burning candle is brought near, the gas burns with a ‘pop’ sound, confirming the presence of hydrogen gas (H₂).

The reaction follows this general equation: Acid + Metal → Salt + Hydrogen gas

Example Reactions:

• Zinc + Sulphuric Acid: Zn + H₂SO₄ → ZnSO₄ + H₂
• Zinc + Hydrochloric Acid: Zn + 2HCl → ZnCl₂ + H₂
• Zinc + Nitric Acid: Zn + 2HNO₃ → Zn(NO₃)₂ + H₂
• Zinc + Acetic Acid: Zn + 2CH₃COOH → Zn(CH₃COO)₂ + H₂

Conclusion:

Metals react with acids to produce hydrogen gas and a salt. The gas evolved burns with a pop sound, confirming the presence of hydrogen.

Activity 2.4: Reaction of Bases with Metals

Procedure:

Take a test tube and place a few pieces of granulated zinc metal. Add 2 mL of sodium hydroxide (NaOH) solution. Warm the contents of the test tube gently. Observe the reaction. Pass the gas evolved through a soap solution and test with a burning candle, as in Activity 2.3.

Observation:

Effervescence occurs, indicating the release of hydrogen gas (H₂). The hydrogen gas burns with a pop sound when tested with a flame. The reaction can be written as:

2NaOH(aq) + Zn(s) → Na₂ZnO₂(aq) + H₂(g) (Sodium zincate is formed along with hydrogen gas.)

Conclusion:

Metals react with strong bases (alkalis) like NaOH to produce hydrogen gas, but such reactions do not occur with all metals. Zinc reacts with NaOH to form sodium zincate and hydrogen gas.

How do Metal Carbonates and Metal Hydrogencarbonates React with Acids?

Activity 2.5: Reaction of Metal Carbonates and Metal Hydrogencarbonates with Acids

Procedure:

Take two test tubes and label them as A and B. Add 0.5 g of sodium carbonate (Na₂CO₃) to test tube A. Add 0.5 g of sodium hydrogencarbonate (NaHCO₃) to test tube B. Pour 2 mL of dilute hydrochloric acid (HCl) into both test tubes. Observe the reaction and the gas evolved. Pass the gas produced through lime water (Ca(OH)₂ solution).

Observations:

• Effervescence is observed in both test tubes, indicating gas evolution.
• The gas turns lime water milky, confirming it is carbon dioxide (CO₂).

Balanced Chemical Equations:

• Test Tube A: Sodium Carbonate + Acid → Salt + Carbon Dioxide + Water

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

• Test Tube B: Sodium Hydrogencarbonate + Acid → Salt + Carbon Dioxide + Water

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

Reaction with Lime Water:

CO₂ + Lime Water → White Precipitate of Calcium Carbonate

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

Excess CO₂ → Calcium Bicarbonate (Soluble in Water)

CaCO₃(s) + H₂O(l) + CO₂(g) → Ca(HCO₃)₂(aq)

Conclusion:

All metal carbonates and metal hydrogencarbonates react with acids to form a salt, carbon dioxide, and water. CO₂ gas turns lime water milky, confirming its presence. Limestone, chalk, and marble are different forms of calcium carbonate (CaCO₃).

2.1.4 Neutralisation Reaction

A base reacts with an acid to form a salt and water.

Example Reaction: NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)

General Formula: Base + Acid → Salt + Water

This reaction neutralises the effect of acids and bases.

2.1.5 Reaction of Metallic Oxides with Acids

Activity 2.7: Reaction of Copper Oxide with Hydrochloric Acid

Procedure:

Take a small amount of copper(II) oxide (CuO) in a beaker. Add dilute hydrochloric acid (HCl) slowly while stirring. Observe the change in colour of the solution.

Observations:

Copper oxide dissolves in acid. The solution turns blue-green, indicating the formation of copper(II) chloride (CuCl₂).

Balanced Chemical Equation:

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

Conclusion:

Metallic oxides are basic in nature. Metal oxides react with acids to form a salt and water, following this general reaction: Metal Oxide + Acid → Salt + Water

Summary of Reactions:

Type of Reaction	General Equation	Example
Metal Carbonate + Acid	Metal Carbonate + Acid → Salt + CO₂ + H₂O	Na₂CO₃ + 2HCl → 2NaCl + CO₂ + H₂O
Metal Hydrogencarbonate + Acid	Metal Hydrogencarbonate + Acid → Salt + CO₂ + H₂O	NaHCO₃ + HCl → NaCl + CO₂ + H₂O
Metal Oxide + Acid	Metal Oxide + Acid → Salt + Water	CuO + 2HCl → CuCl₂ + H₂O
Neutralisation Reaction	Base + Acid → Salt + Water	NaOH + HCl → NaCl + H₂O

How do Acids and Bases React with Each Other?

Activity 2.6: Reaction of an Acid with a Base (Neutralisation Reaction)

Procedure:

Take 2 mL of dilute sodium hydroxide (NaOH) solution in a test tube. Add two drops of phenolphthalein indicator to the test tube. Observe and note the colour of the solution. Slowly add dilute hydrochloric acid (HCl) drop by drop to the solution while stirring. Observe the colour change in the solution. Now, add a few more drops of NaOH to the same mixture and observe the change in colour.

Observations:

• Initially, the solution turns pink due to phenolphthalein in a basic medium (NaOH solution).
• Upon adding HCl, the pink colour disappears, indicating the neutralisation of the base.
• When NaOH is added again, the pink colour reappears, confirming that the solution has become basic again.

Explanation:

Phenolphthalein is a pH indicator that is pink in basic solutions and colourless in acidic solutions. When HCl (acid) is added, it neutralises NaOH (base), forming NaCl (salt) and water, causing the colour to disappear. When more NaOH (base) is added, the solution becomes basic again, and the pink colour reappears.

Balanced Chemical Equation for Neutralisation Reaction:

NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)

Conclusion:

Acids and bases neutralise each other to form a salt and water. This reaction is called a neutralisation reaction.

General equation for neutralisation:

Base + Acid → Salt + Water

Summary Table of Neutralisation Reaction:

Component	Before Neutralisation	After Adding Acid	After Adding Base Again
Solution Type	Basic (NaOH)	Neutralised (Salt + Water)	Basic (NaOH) Again
Phenolphthalein Colour	Pink	Colourless	Pink Again
Reaction Type	Base	Neutralisation	Base Reappears

Key Takeaways:

• Neutralisation occurs when an acid reacts with a base to form a salt and water.
• Phenolphthalein is pink in a basic solution and colourless in an acidic solution.
• The reaction follows the general equation: Acid + Base → Salt + Water
• This process is used in antacids, soil treatment, and industrial applications.

Reaction of Metallic Oxides with Acids

Activity 2.7: Reaction of Copper Oxide with Hydrochloric Acid

Procedure:

Take a small amount of copper oxide (CuO) in a beaker. Slowly add dilute hydrochloric acid (HCl) to the beaker while stirring. Observe the colour change in the solution.

Observations:

Copper oxide (CuO) is black in colour. When HCl is added, the black CuO dissolves and forms a blue-green solution.

Balanced Chemical Equation for the Reaction:

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

Explanation:

Copper oxide (CuO) is a metallic oxide. Metallic oxides react with acids to form salt and water, similar to the reaction of bases with acids. This shows that metallic oxides are basic in nature.

General Reaction of a Metallic Oxide with an Acid:

Metal oxide + Acid → Salt + Water

Key Takeaways:

• Metallic oxides react with acids to form salts and water.
• Since metallic oxides neutralise acids, they are considered basic in nature.
• Example: Copper(II) oxide (CuO) reacts with hydrochloric acid (HCl) to form copper(II) chloride (CuCl₂) and water.
• This reaction is similar to the reaction of a base with an acid (neutralisation reaction).

2.1.6 Reaction of a Non-Metallic Oxide with a Base

Example Reaction: Carbon Dioxide (CO₂) and Calcium Hydroxide [Ca(OH)₂]

In Activity 2.5, carbon dioxide (CO₂) was passed through lime water (Ca(OH)₂). A white precipitate of calcium carbonate (CaCO₃) was formed.

Balanced Chemical Equation:

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

On passing excess CO₂, the white precipitate dissolves, forming calcium bicarbonate:

CaCO₃(s) + CO₂(g) + H₂O(l) → Ca(HCO₃)₂(aq)

Conclusion:

• Non-metallic oxides react with bases to form salts and water.
• This reaction is similar to the neutralisation reaction between acids and bases.
• Therefore, non-metallic oxides are acidic in nature.
• Example: CO₂ (a non-metallic oxide) reacts with Ca(OH)₂ (a base) to form CaCO₃ (a salt) and water.

What Do All Acids and Bases Have in Common?

Common Properties of Acids:

• All acids exhibit similar chemical properties due to the presence of hydrogen ions (H⁺).
• Acids react with metals to produce hydrogen gas (H₂).
• Acids conduct electricity in aqueous solutions due to the presence of free H⁺ ions.

Activity 2.8: Investigating Acids and Bases with Electricity

Materials Required:

• Solutions of glucose, alcohol, HCl, H₂SO₄ (sulphuric acid), etc.
• A 100 mL beaker, two nails fixed on a cork.
• 6V battery, bulb, switch, wires for the electric circuit.

Procedure:

Pour dilute HCl into the beaker and connect the circuit. Switch on the current and observe whether the bulb glows. Repeat the experiment with dilute H₂SO₄. Perform the experiment again using glucose and alcohol solutions. Observe whether the bulb glows in all cases.

Observations:

• The bulb glows in the case of acids (HCl, H₂SO₄), indicating electric current flow.
• The bulb does not glow in glucose and alcohol solutions, meaning they do not conduct electricity.

Conclusion:

H⁺ ions are responsible for the acidic nature of acids. Acids release H⁺ ions in aqueous solutions, making them electrolytes.

Testing Bases for Conductivity:

When the experiment is repeated with bases (NaOH, Ca(OH)₂, etc.), the bulb also glows. This indicates that bases release ions (OH⁻) in solution, which help conduct electricity.

Key Takeaways:

• All acids produce H⁺ ions in aqueous solutions, which gives them their acidic properties.
• Bases produce OH⁻ ions, which give them their basic properties.
• Only substances that ionize in water can conduct electricity.

What Happens to an Acid or a Base in a Water Solution?

Do Acids Produce Ions Only in Aqueous Solution?

Experiment (Activity 2.9):

Solid NaCl is taken in a test tube. Concentrated H₂SO₄ is added. The evolved gas is tested with dry and wet blue litmus paper.

Observations:

• The gas does not change the color of dry blue litmus.
• The gas turns wet blue litmus red, indicating acidic properties.

Inference:

Dry HCl gas is not acidic, as it does not release H⁺ ions. HCl in water becomes acidic, as it ionizes to form H₃O⁺ (hydronium ions).

Hydrogen ions (H⁺) do not exist alone, but combine with water to form hydronium ions:

HCl + H₂O → H₃O⁺ + Cl⁻

H⁺ + H₂O → H₃O⁺

Dissolution of Bases in Water

Reaction of Bases with Water:

• NaOH(s) → Na⁺(aq) + OH⁻(aq)
• KOH(s) → K⁺(aq) + OH⁻(aq)
• Mg(OH)₂(s) → Mg²⁺(aq) + 2OH⁻(aq)

Bases release OH⁻ ions in aqueous solution. Bases that dissolve in water are called alkalis.

Neutralization Reaction

Acid + Base → Salt + Water

H⁺(aq) + OH⁻(aq) → H₂O(l)

Effect of Mixing Acids and Bases with Water (Activity 2.10)

Procedure:

10 mL of water is taken in a beaker. A few drops of concentrated H₂SO₄ are added slowly while stirring. The temperature of the beaker is noted. The experiment is repeated with sodium hydroxide pellets.

Observations:

Temperature increases during both reactions. The process is highly exothermic.

Precautions:

• Always add acid to water slowly, never water to acid.
• This prevents splashing and burns caused by the heat released.
• Containers of concentrated acids and bases carry warning signs.

Dilution of Acids and Bases

Dilution is the process of adding water to an acid or base to decrease the concentration of H₃O⁺ or OH⁻ ions per unit volume. Diluted acids and bases are less reactive than concentrated ones.

Strength of Acid and Base Solutions (pH Scale):

Acid-Base Indicators:

Used to distinguish between acids and bases.

Universal Indicator: A mixture of several indicators that shows different colors at different pH levels.

pH Scale:

Measures the hydrogen ion (H⁺) concentration in a solution. Ranges from 0 to 14:

• pH less than 7 → Acidic solution (Higher H₃O⁺ concentration).
• pH = 7 → Neutral solution (e.g., pure water).
• pH > 7 → Alkaline (basic) solution (Higher OH⁻ concentration).

Strength of Acids and Bases:

• Strong acids ionize completely in water and produce more H⁺ ions (e.g., HCl, H₂SO₄).
• Weak acids ionize partially, producing fewer H⁺ ions (e.g., acetic acid).
• Strong bases produce more OH⁻ ions (e.g., NaOH, KOH).
• Weak bases produce fewer OH⁻ ions (e.g., NH₄OH).

Testing pH:

pH paper or a universal indicator is used to determine the approximate pH value of a solution.

Example:

• Lemon juice (acidic) → pH around 2-3
• Tap water (neutral) → pH around 7
• NaOH (strong base) → pH around 13-14

pH and Concentration Relationship:

Higher H₃O⁺ concentration = Lower pH (strong acid)
Higher OH⁻ concentration = Higher pH (strong base)

Importance of pH in Everyday Life:

pH Sensitivity in Living Organisms

The human body functions within a pH range of 7.0 to 7.8. Living organisms can survive only within a narrow pH range. Acid rain (pH less than 5.6) lowers the pH of river water, making survival difficult for aquatic life. Venus atmosphere contains thick clouds of sulphuric acid, making it unsuitable for life.

pH and Soil for Plants

Plants require a specific pH range for healthy growth.

Activity to check soil pH:

Mix 2g soil with 5mL water in a test tube. Shake and filter the solution. Test the pH using universal indicator paper. Ideal p H varies for different plants.

pH in Digestion

The stomach produces hydrochloric acid (HCl) to help digest food. Excess acid production causes pain and irritation (indigestion). Antacids (mild bases like magnesium hydroxide) neutralize excess acid.

pH and Tooth Decay

Tooth decay starts when pH less that 5.5. Tooth enamel (made of calcium hydroxyapatite) dissolves at low pH. Bacteria in the mouth produce acid from sugar, leading to decay. Toothpaste (basic in nature) neutralizes acid and prevents decay.

Self-Defense in Animals and Plants (Chemical Warfare)

Bee stings contain acid, causing pain and irritation. Remedy: Apply a mild base (baking soda) to neutralize it. Nettle plant stings contain methanoic acid, causing burning pain. Remedy: Dock plant leaves (basic in nature) neutralize the sting.

Naturally Occurring Acids

Natural Source	Acid	Natural Source	Acid
Vinegar	Acetic acid	Sour milk (Curd)	Lactic acid
Orange	Citric acid	Lemon	Citric acid
Tamarind	Tartaric acid	Ant sting	Methanoic acid
Tomato	Oxalic acid	Nettle sting	Methanoic acid

Salts

1. Definition of Salts:

Salts are ionic compounds formed when an acid reacts with a base during a neutralization reaction.

General reaction: Acid + Base → Salt + Water

2. Chemical Formulae of Some Common Salts:

Salt Name	Chemical Formula	Acid Used	Base Used
Potassium sulphate	K₂SO₄	H₂SO₄ (Sulfuric acid)	KOH (Potassium hydroxide)
Sodium sulphate	Na₂SO₄	H₂SO₄	NaOH (Sodium hydroxide)
Calcium sulphate	CaSO₄	H₂SO₄	Ca(OH)₂ (Calcium hydroxide)
Magnesium sulphate	MgSO₄	H₂SO₄	Mg(OH)₂ (Magnesium hydroxide)
Copper sulphate	CuSO₄	H₂SO₄	Cu(OH)₂ (Copper hydroxide)
Sodium chloride	NaCl	HCl (Hydrochloric acid)	NaOH
Sodium nitrate	NaNO₃	HNO₃ (Nitric acid)	NaOH
Sodium carbonate	Na₂CO₃	H₂CO ₃ (Carbonic acid)	NaOH
Ammonium chloride	NH₄Cl	HCl	NH₄OH (Ammonium hydroxide)

3. Families of Salts:

Salts with the same positive ion (cation) or negative ion (anion) belong to the same family.

Identified families from the given salts:

• Sodium Salts Family: NaCl, Na₂SO₄, NaNO₃, Na₂CO₃
• Sulphate Salts Family: K₂SO₄, Na₂SO₄, CaSO₄, MgSO₄, CuSO₄
• Chloride Salts Family: NaCl, NH₄Cl
• Carbonate Salts Family: Na₂CO₃

4. Key Properties of Salts:

• Solubility: Some salts are soluble in water (e.g., NaCl, K₂SO₄), while others are insoluble (e.g., PbSO₄, BaSO₄).
• Conductivity: Salts dissolve in water to form electrolytes, which conduct electricity.
• pH Nature:
  • Neutral salts (e.g., NaCl) are formed from strong acid and strong base.
  • Acidic salts (e.g., NH₄Cl) are formed from strong acid and weak base.
  • Basic salts (e.g., Na₂CO₃) are formed from weak acid and strong base.

5. Uses of Salts:

• Sodium chloride (NaCl): Used in cooking, food preservation, and making chlorine.
• Sodium carbonate (Na₂CO₃): Used in glass manufacturing and washing soda.
• Copper sulphate (CuSO₄): Used in fungicides and electroplating.
• Calcium sulphate (CaSO₄): Used in Plaster of Paris and cement.

pH of Salts

1. Understanding the pH of Salts:

The pH of a salt solution depends on the strength of the acid and base that formed it.

Types of salts based on pH:

• Neutral Salts (pH = 7): Formed from a strong acid and a strong base.
• Acidic Salts (pH less than 7): Formed from a strong acid and a weak base.
• Basic Salts (pH > 7): Formed from a weak acid and a strong base.

2. Testing the pH of Salt Solutions (Activity 2.14):

Salt	Solubility in Water	Effect on Litmus	Approximate pH	Nature of Salt	Formed from
Sodium chloride (NaCl)	Soluble	No change	7	Neutral	HCl + NaOH
Potassium nitrate (KNO₃)	Soluble	No change	7	Neutral	HNO₃ + KOH
Aluminium chloride (AlCl₃)	Soluble	Red litmus stays red, blue litmus turns red	less than 7	Acidic	HCl + Al(OH)₃
Zinc sulphate (ZnSO₄)	Soluble	Red litmus stays red, blue litmus turns red	less than 7	Acidic	H₂SO₄ + Zn(OH)₂
Copper sulphate (CuSO₄)	Soluble	Red litmus stays red, blue litmus turns red	less than7	Acidic	H₂SO₄ + Cu(OH)₂
Sodium acetate (CH₃COONa)	Soluble	Blue litmus stays blue, red litmus turns blue	> 7	Basic	CH₃COOH + NaOH
Sodium carbonate (Na₂CO₃)	Soluble	Blue litmus stays blue, red litmus turns blue	> 7	Basic	H₂CO₃ + NaOH
Sodium bicarbonate (NaHCO₃)	Soluble	Blue litmus stays blue, red litmus turns slightly blue	Slightly > 7	Weakly Basic	H₂CO₃ + NaOH

3. Summary of Acidic, Basic, and Neutral Salts:

• Neutral Salts (pH = 7): Formed from a strong acid + strong base (e.g., NaCl, KNO₃).
• Acidic Salts (pH less than 7): Formed from a strong acid + weak base (e.g., AlCl₃, ZnSO₄, CuSO₄).
• Basic Salts (pH > 7): Formed from a weak acid + strong base (e.g., Na₂CO₃, CH₃COONa, NaHCO₃).

4. Importance of pH in Salt Solutions:

pH affects chemical reactions, soil fertility, and biological processes. Neutral salts (like NaCl) do not alter the pH of a solution. Acidic salts lower pH and can cause corrosion or acidity in the environment. Basic salts increase pH and are used in antacids, water treatment, and industrial applications.

Chemicals from Common Salt

1. Common Salt (Sodium Chloride - NaCl) as a Raw Material:

Sodium chloride (NaCl) is obtained from seawater and rock salt deposits. It is used in the food industry and serves as a raw material for many industrial chemicals.

2. Chemicals Derived from Common Salt:

(i) Sodium Hydroxide (NaOH) - The Chlor-Alkali Process

Formed by the electrolysis of brine (NaCl solution).

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

Products and Uses:

• Sodium hydroxide (NaOH): Used in soap, paper, textile, and detergent industries.
• Chlorine (Cl₂): Used in disinfectants, plastics (PVC), and bleaching powder.
• Hydrogen (H₂): Used in fuels and ammonia production.

( ii) Bleaching Powder (CaOCl₂)

Made by reacting chlorine gas (from electrolysis) with slaked lime (Ca(OH)₂).

Reaction: Ca(OH)₂ + Cl₂ → CaOCl₂ + H₂O

Uses:

• Bleaching agent in textiles, paper, and laundry.
• Oxidizing agent in industries.
• Disinfectant for making drinking water germ-free.

(iii) Baking Soda (Sodium Hydrogen Carbonate - NaHCO₃)

Formed using sodium chloride as a raw material.

Reaction: NaCl + H₂O + CO₂ + NH₃ → NH₄Cl + NaHCO₃

Uses:

• Baking powder (mix of NaHCO₃ + edible acid). When heated, it releases CO₂, making food soft and spongy.
• Antacid: Neutralizes stomach acid.
• Used in fire extinguishers (releases CO₂ when heated).

Decomposition Reaction of Baking Soda:

2NaHCO₃ → Heat → Na₂CO₃ + CO₂ + H₂O

(iv) Washing Soda (Sodium Carbonate - Na₂CO₃·10H₂O)

Formed by recrystallization of sodium carbonate from baking soda.

Reaction: Na₂CO₃ + 10H₂O → Na₂CO₃·10H₂O

Uses:

• Used in glass, soap, and paper industries.
• Used to manufacture sodium compounds like borax.
• Used as a cleaning agent in households.
• Helps remove permanent hardness of water.

Note: The 10H₂O molecules indicate water of crystallization, which does not make washing soda wet but helps form crystals.

3. Summary Table of Chemicals from Common Salt

Chemical	Formula	Formation	Key Uses
Sodium Hydroxide	NaOH	Electrolysis of brine	Soap, detergents, paper, textiles
Bleaching Powder	CaOCl₂	Reaction of Cl₂ with Ca(OH)₂	Disinfectant, bleaching agent
Baking Soda	NaHCO₃	Reaction of NaCl, H₂O, CO₂, NH₃	Baking, antacid, fire extinguisher
Washing Soda	Na₂CO₃·10H₂O	Recrystallization of Na₂CO₃	Cleaning, water softening, glass & soap making

Water of Crystallisation & Plaster of Paris

1. Water of Crystallisation in Salts

Definition: The fixed number of water molecules present in one formula unit of a salt.

Example: Copper sulphate (CuSO₄·5H₂O) contains five water molecules. Gypsum (CaSO₄·2H₂O) contains two water molecules.

2. Activity: Heating Copper Sulphate Crystals

Observation:

• Before heating: Blue crystals of CuSO₄·5H₂O.
• After heating: White powder of anhydrous CuSO₄ (loses water).
• Water droplets form on the boiling tube (released from CuSO₄·5H₂O).
• Adding water back: Blue colour is restored.

Conclusion:

Hydrated copper sulphate contains water of crystallisation. Removing water changes its colour from blue to white. Rehydration restores the blue colour.

3. Water of Crystallisation in Gypsum & Plaster of Paris

Gypsum (CaSO₄·2H₂O) contains two water molecules as water of crystallisation.

Uses: Making Plaster of Paris (POP) and cement.

Plaster of Paris (POP) - CaSO₄·½H₂O

Formed by heating gypsum at 373 K: CaSO₄·2H₂O → 373K CaSO₄·½H₂O + 1½H₂O

Properties & Uses:

• White powder that hardens when mixed with water (reverts to gypsum).
• Used for fracture casts, moldings, and decorative materials.
• Helps make surfaces smooth.

4. Why is it called Plaster of Paris?

It is named after Paris, France, where large deposits of gypsum were found and used to make the material.