Gonit Vigyan

 

1. Acids and Bases in Daily Life

• Sour and bitter tastes in food are due to the presence of acids and bases, respectively.
• Example: Lemon juice and vinegar are acidic, while baking soda solution is basic.
• Acidity Remedy: If someone suffers from acidity due to overeating, baking soda solution should be used because it is a base and neutralizes the excess stomach acid.
• Concept Applied: Neutralization reaction – acids and bases cancel each other’s effects.

2. Properties of Acids and Bases

Property	Acids	Bases
Taste	Sour	Bitter
Litmus Test	Turns blue litmus red	Turns red litmus blue
Other Indicators	Methyl orange → Red Phenolphthalein → Colorless	Methyl orange → Yellow Phenolphthalein → Pink
Nature	Corrosive in strong forms	Slippery, soapy feel

3. Acid-Base Indicators

Indicators are substances that help identify whether a solution is acidic or basic.

• Natural Indicators:
  • Litmus (obtained from lichen):
    • Acidic solution: Turns red
    • Basic solution: Turns blue
  • Turmeric:
    • Basic solution: Turns reddish-brown
    • Neutral or Acidic solution: Stays yellow
  • Red cabbage leaves, Hydrangea, Petunia, and Geranium petals also change color in the presence of acids or bases.
• Synthetic Indicators:
  • Methyl Orange:
    • Acid → Red
    • Base → Yellow
  • Phenolphthalein:
    • Acid → Colorless
    • Base → Pink

4. Neutralization Reaction

• Definition: When an acid and a base react, they neutralize each other, forming salt and water.
• Example:
  • Hydrochloric acid (HCl) + Sodium hydroxide (NaOH) → Sodium chloride (NaCl) + Water (H₂O)

5. Practical Applications of Indicators

• Testing pH of substances: Helps determine whether a solution is acidic or basic.
• Curry stain on a white cloth turns reddish-brown when soap (a base) is applied and turns yellow again when washed with water.
• Litmus paper tests are widely used in chemistry labs to classify substances.

6. Key Takeaways

• Acids turn blue litmus red, while bases turn red litmus blue.
• Natural and synthetic indicators help in identifying acids and bases.
• Neutralization reactions are used in medicine (antacids), agriculture (liming soil), and industry.

Question

You have been provided with three test tubes. One contains distilled water, another contains an acidic solution, and the third contains a basic solution. If you are given only red litmus paper, how will you identify the contents of each test tube?

Answer:

To identify the contents of each test tube using only red litmus paper, follow these steps:

1. Dip red litmus paper into each test tube one by one and observe the color change.
2. Interpret the results:
   • If the red litmus remains red, the solution is either acidic or neutral (distilled water).
   • If the red litmus turns blue, the solution is basic.
3. Further differentiation between acid and neutral solutions:
   • Since distilled water is neutral, it does not change the color of red litmus.
   • The test tube in which the red litmus remains red (and is not basic) is either acidic or neutral.
   • To confirm the acidic solution, you would need blue litmus paper (which would turn red in an acid), but since you only have red litmus, you can assume that the test tube where red litmus remains red and is not neutral contains the acid.

Final Identification:

• Test tube where red litmus turns blue → Basic solution
• Test tube where red litmus remains red (not basic) → Either acidic or neutral
• Among the remaining two, the one that is neutral (distilled water) does not affect any litmus paper, while the acidic solution would turn blue litmus red (if available).

2.1 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 Required:

• 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)
• Indicators: Red litmus, Blue litmus, Phenolphthalein, Methyl orange
• Equipment: Watch-glass, dropper

Procedure:

1. Take a drop of each acid and base solution on a watch-glass.
2. Add a drop of different indicators to each solution one by one.
3. Observe the color changes and record them in a table.

Observations (Table 2.1)

Sample Solution	Red Litmus	Blue Litmus	Phenolphthalein	Methyl Orange
Acids	Turns red (no change)	Turns red	Colorless	Turns red
Bases	Turns blue	No change	Pink	Turns yellow

🔹 Conclusion:

• Acids turn blue litmus red, keep red litmus unchanged, turn phenolphthalein colorless, and methyl orange red.
• Bases turn red litmus blue, keep blue litmus unchanged, turn phenolphthalein pink, and methyl orange yellow.
• Indicators help identify whether a solution is acidic or basic based on color changes.

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Activity 2.2: Testing Acids and Bases with Olfactory Indicators

Olfactory Indicators:

Some substances change odour in the presence of acids or bases. Examples: Onion, Vanilla, Clove oil

Procedure:

Part 1: Testing Onion as an Olfactory Indicator

1. Prepare onion strips:
   • Take some finely chopped onions in a plastic bag.
   • Place clean cloth strips inside the bag.
   • Tie the bag tightly and leave it overnight in the fridge.
   • The cloth strips absorb the onion’s smell.
2. Testing the odour change:
   • Take two strips and check their odour.
   • Place one strip on a clean surface and add a few drops of dilute HCl.
   • Place the other strip on another surface and add a few drops of dilute NaOH.
   • Rinse both cloth strips with water and check the odour again.

Observations:

• The onion smell disappears in NaOH but remains in HCl.

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Part 2: Testing Vanilla and Clove Oil as Olfactory Indicators

1. Take two test tubes:
   • One with dilute HCl
   • One with dilute NaOH
2. Add a few drops of dilute vanilla essence to both test tubes, shake well, and observe the odour.
3. Repeat the same process with clove oil and record any changes in odour.

Observations:

Olfactory Indicator	Effect in Acid (HCl)	Effect in Base (NaOH)
Onion	Smell remains	Smell disappears
Vanilla	Smell remains	Smell disappears
Clove Oil	Smell remains	Smell disappears

🔹 Conclusion:

• Onion, vanilla, and clove oil can be used as olfactory indicators because their odour changes in acidic and basic conditions.
• Basic solutions remove the characteristic odour of these substances, while acidic solutions retain it.

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Key Takeaways:

1. Indicators like litmus, phenolphthalein, and methyl orange help identify acids and bases through color changes.
2. Olfactory indicators (onion, vanilla, clove) change their smell in the presence of acids and bases, providing an additional way to distinguish between them.
3. Acids and bases show distinct chemical properties that can be tested through simple lab experiments.

2.1.2 How Do Acids and Bases React with Metals?

Activity 2.3: Reaction of Acids with Metals

Materials Required:

• Dilute acids: Sulphuric acid (H₂SO₄), Hydrochloric acid (HCl), Nitric acid (HNO₃), Acetic acid (CH₃COOH)
• Metal: Zinc granules
• Equipment: Test tube, delivery tube, soap solution, candle/matchstick

Procedure:

1. Set up the apparatus as shown in Fig. 2.1.
2. Take about 5 mL of dilute sulphuric acid in a test tube.
3. Add a few pieces of zinc granules into the acid.
4. Observe the reaction on the zinc surface.
5. Pass the gas evolved through a soap solution.
   • Bubbles form in the soap solution, indicating a gas is being released.
6. Take a burning candle near a gas-filled bubble and observe the effect.
7. Repeat this activity with HCl, HNO₃, and CH₃COOH to compare results.

Observations:

• Zinc reacts with the acid, and bubbles appear on its surface.
• The bubbles in the soap solution confirm that a gas (hydrogen) is evolving.
• When a burning candle is brought near the gas bubble, it produces a ‘pop’ sound, confirming the presence of hydrogen gas.
• The reaction is similar for other acids (HCl, HNO₃, CH₃COOH) but the rate of reaction may vary.

Reaction Equation:

$$\text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen gas}$$

For Sulphuric acid and Zinc:

$$H_2SO_4 + Zn \rightarrow ZnSO_4 + H_2 \uparrow$$

For Hydrochloric acid and Zinc:

$$2HCl + Zn \rightarrow ZnCl_2 + H_2 \uparrow$$

🔹 Conclusion:

• Acids react with metals to form a salt and release hydrogen gas.
• Hydrogen gas burns with a ‘pop’ sound when tested with a burning candle.
• The reactivity depends on the nature of the acid and the metal used.

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Activity 2.4: Reaction of Bases with Metals

Materials Required:

• Metal: Zinc granules
• Base: Sodium hydroxide (NaOH) solution
• Equipment: Test tube, burner, soap solution, candle/matchstick

Procedure:

1. Place a few pieces of granulated zinc metal in a test tube.
2. Add 2 mL of sodium hydroxide (NaOH) solution.
3. Warm the test tube gently.
4. Pass the gas evolved through a soap solution and test with a burning candle.

Observations:

• Zinc reacts with sodium hydroxide, releasing hydrogen gas.
• Bubbles form in the soap solution, confirming gas evolution.
• The gas burns with a ‘pop’ sound when a burning candle is brought near.

Reaction Equation:

$$2NaOH + Zn \rightarrow Na_2ZnO_2 + H_2 \uparrow$$

(Sodium zincate is formed along with hydrogen gas.)

🔹 Conclusion:

• Bases do not react with all metals, but zinc reacts with strong bases like NaOH to release hydrogen gas.
• The reaction produces a complex salt (sodium zincate) and hydrogen gas.

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Key Takeaways:

1. Acids react with metals to form a salt and hydrogen gas.
2. Bases react with certain metals (like zinc) to form a complex salt and hydrogen gas.
3. The ‘pop’ sound test confirms the presence of hydrogen gas.
4. Not all metals react with acids and bases in the same way.
5. Reactivity depends on the nature of the acid/base and the type of metal used.

These reactions help us understand how metals behave in acidic and basic environments and are crucial in industries like metal refining and laboratory analysis.

2.1.3 How Do Metal Carbonates and Metal Hydrogencarbonates React with Acids?

Activity 2.5: Reaction of Sodium Carbonate and Sodium Hydrogencarbonate with Acid

Materials Required:

• Test tubes (2), dilute hydrochloric acid (HCl), sodium carbonate (Na₂CO₃), sodium hydrogencarbonate (NaHCO₃), lime water (Ca(OH)₂ solution), delivery tube.

Procedure:

1. Take two test tubes, label them as A and B.
2. Add about 0.5 g of sodium carbonate (Na₂CO₃) in test tube A.
3. Add about 0.5 g of sodium hydrogencarbonate (NaHCO₃) in test tube B.
4. Pour 2 mL of dilute hydrochloric acid (HCl) into each test tube.
5. Observe the gas evolution in both test tubes.
6. Pass the gas evolved through lime water (Ca(OH)₂ solution) using a delivery tube and note the changes.

Observations:

• Effervescence (bubbles) is observed in both test tubes due to the evolution of carbon dioxide (CO₂) gas.
• When the gas is passed through lime water, it turns milky due to the formation of calcium carbonate (CaCO₃), a white precipitate.
• If excess CO₂ is passed, the milkiness disappears due to the formation of soluble calcium bicarbonate [Ca(HCO₃)₂].

Reaction Equations:

1. Reaction of Sodium Carbonate (Na₂CO₃) with HCl

$$Na_2CO_3 (s) + 2HCl (aq) \rightarrow 2NaCl (aq) + H_2O (l) + CO_2 (g)$$

2. Reaction of Sodium Hydrogencarbonate (NaHCO₃) with HCl

$$NaHCO_3 (s) + HCl (aq) \rightarrow NaCl (aq) + H_2O (l) + CO_2 (g)$$

3. Reaction of CO₂ with Lime Water (Ca(OH)₂)

$$Ca(OH)_2 (aq) + CO_2 (g) \rightarrow CaCO_3 (s) + H_2O (l)$$

(White precipitate of calcium carbonate forms, making lime water milky.)

4. When Excess CO₂ is Passed

$$CaCO_3 (s) + CO_2 (g) + H_2O (l) \rightarrow Ca(HCO_3)_2 (aq)$$

(Milkiness disappears as calcium bicarbonate is soluble in water.)

Conclusion:

• All metal carbonates and metal hydrogencarbonates react with acids to form a salt, carbon dioxide, and water.
• Carbon dioxide turns lime water milky, which is a confirmatory test for CO₂ gas.
• Excess CO₂ makes the solution clear due to the formation of soluble calcium bicarbonate.

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2.1.4 Neutralization Reaction (Acid + Base Reaction)

When an acid and a base react, they form a salt and water. This is called a neutralization reaction.

Example:

$$NaOH (aq) + HCl (aq) \rightarrow NaCl (aq) + H_2O (l)$$

(Sodium hydroxide reacts with hydrochloric acid to form sodium chloride and water.)

General Equation:

$$Base + Acid \rightarrow Salt + Water$$

🔹 Key Point:

• Neutralization helps in reducing the effect of excess acids or bases.
• It is used in antacids to relieve acidity in the stomach.

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2.1.5 Reaction of Metallic Oxides with Acids

Activity 2.7: Reaction of Copper Oxide with Acid

Materials Required:

• Copper oxide (CuO), dilute hydrochloric acid (HCl), beaker, stirring rod.

Procedure:

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

Observations:

• Copper oxide dissolves in HCl, and the solution turns blue-green due to the formation of copper(II) chloride (CuCl₂).

Reaction Equation:

$$CuO (s) + 2HCl (aq) \rightarrow CuCl_2 (aq) + H_2O (l)$$

🔹 General Equation:

$$\text{Metal oxide} + \text{Acid} \rightarrow \text{Salt} + \text{Water}$$

Conclusion:

• Metallic oxides react with acids to form a salt and water.
• Since metal oxides react with acids, they are basic in nature.

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Key Takeaways:

1. Reaction of Metal Carbonates and Hydrogencarbonates with Acids

✅ Metal carbonates (Na₂CO₃) and metal hydrogencarbonates (NaHCO₃) react with acids to form a salt, CO₂ gas, and water.
✅ CO₂ gas turns lime water milky due to calcium carbonate formation.
✅ Excess CO₂ makes the solution clear due to calcium bicarbonate formation.

2. Neutralization Reaction

✅ An acid reacts with a base to form a salt and water.
✅ Example: NaOH + HCl → NaCl + H₂O.
✅ Neutralization reactions help in balancing pH levels in various applications.

3. Reaction of Metallic Oxides with Acids

✅ Metal oxides (CuO) react with acids to form salt and water.
✅ Example: CuO + HCl → CuCl₂ + H₂O (Blue-green solution).
✅ Since metal oxides react with acids, they exhibit basic properties.

These reactions are fundamental to understanding acid-base chemistry and their practical applications in daily life and industries. 🚀

2.1.6 Reaction of a Non-metallic Oxide with a Base

Activity: Reaction of Carbon Dioxide with a Base (Calcium Hydroxide)

Materials Required:

• Calcium hydroxide solution (Ca(OH)₂) – lime water
• Carbon dioxide gas (CO₂) – produced from metal carbonate reaction
• Test tube and delivery tube

Procedure:

1. Take some freshly prepared lime water (Ca(OH)₂ solution) in a test tube.
2. Pass carbon dioxide (CO₂) gas through lime water using a delivery tube.
3. Observe any changes in the solution.
4. Continue passing CO₂ and note any further changes.

Observations:

• Initially, the lime water turns milky due to the formation of insoluble calcium carbonate (CaCO₃).
• If excess CO₂ is passed, the milkiness disappears because calcium bicarbonate [Ca(HCO₃)₂] is formed, which is soluble in water.

Reaction Equations:

1. Reaction of Carbon Dioxide with Lime Water:

$$Ca(OH)_2 (aq) + CO_2 (g) \rightarrow CaCO_3 (s) + H_2O (l)$$

(White precipitate of calcium carbonate forms, making lime water milky.)

2. Effect of Excess CO₂:

$$CaCO_3 (s) + CO_2 (g) + H_2O (l) \rightarrow Ca(HCO_3)_2 (aq)$$

(Milkiness disappears as calcium bicarbonate dissolves in water.)

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Conclusion:

✅ Non-metallic oxides (like CO₂) react with bases to form a salt and water, similar to an acid-base reaction.
✅ This confirms that non-metallic oxides are acidic in nature.
✅ CO₂ reacts with Ca(OH)₂, forming CaCO₃ (insoluble) first, then Ca(HCO₃)₂ (soluble) when excess CO₂ is passed.

Key Concept:

• Since CO₂ reacts with a base in the same way that acids do, it indicates that non-metallic oxides exhibit acidic properties.
• Other non-metallic oxides, such as SO₂, NO₂, and P₂O₅, also form acids or acidic salts when dissolved in water.

This experiment is a key demonstration of the acidic nature of non-metallic oxides and their role in environmental chemistry, such as acid rain formation (SO₂ + H₂O → H₂SO₃, forming sulfurous acid). 🌍

Questions and Answers

1. Why should curd and sour substances not be kept in brass and copper vessels?
✅ Answer: Curd and sour substances contain acids like lactic acid and citric acid. When these acids react with brass (a copper-zinc alloy) or copper vessels, they form toxic salts (like copper acetate or copper lactate), which can be harmful to health.

Reaction Example:

$$Cu + Acid \rightarrow \text{Toxic Copper Salt}$$

This is why curd and sour substances should be stored in glass or food-safe plastic containers instead.

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2. Which gas is usually liberated when an acid reacts with a metal? Illustrate with an example. How will you test for the presence of this gas?
✅ Answer: When an acid reacts with a metal, hydrogen gas (H₂) is usually liberated.

Example Reaction:

$$\text{Zn} (s) + 2HCl (aq) \rightarrow ZnCl_2 (aq) + H_2 (g)$$

(Zinc reacts with hydrochloric acid to form zinc chloride and hydrogen gas.)

Test for Hydrogen Gas:

• Bring a burning matchstick or candle near the gas bubbles.
• If hydrogen is present, it burns with a 'pop' sound, confirming its presence.

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3. Metal compound A reacts with dilute hydrochloric acid to produce effervescence. The gas evolved extinguishes a burning candle. Write a balanced chemical equation for the reaction if one of the compounds formed is calcium chloride.
✅ Answer: The effervescence and extinguishing of a candle suggest that the gas produced is carbon dioxide (CO₂). This means that metal compound A is calcium carbonate (CaCO₃).

Balanced Chemical Equation:

$$CaCO_3 (s) + 2HCl (aq) \rightarrow CaCl_2 (aq) + CO_2 (g) + H_2O (l)$$

Explanation:

• Calcium carbonate (CaCO₃) reacts with hydrochloric acid (HCl) to form calcium chloride (CaCl₂), carbon dioxide (CO₂), and water (H₂O).
• The CO₂ gas turns lime water milky and extinguishes a burning candle.

✅ This confirms the gas evolved is carbon dioxide.

 "What Do All Acids and All Bases Have in Common?"

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1. Common Properties of Acids

• All acids exhibit similar chemical properties because they release hydrogen ions (H⁺) in aqueous solutions.
• When acids react with metals, they produce hydrogen gas (H₂).
• Acids conduct electricity in solution due to the presence of free H⁺ ions.
• Not all compounds containing hydrogen are acidic (e.g., glucose and alcohol contain hydrogen but do not show acidic properties).

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2. Activity 2.8: Testing Conductivity of Acids and Bases

Materials Required:

• Solutions of glucose, alcohol, hydrochloric acid (HCl), sulphuric acid (H₂SO₄), sodium hydroxide (NaOH), calcium hydroxide (Ca(OH)₂)
• Two nails fixed on a cork
• 100 mL beaker
• 6V battery, bulb, and switch

Procedure:

1. Set up the apparatus as shown in Fig. 2.3, with two nails connected to a 6V battery and a bulb.
2. Pour dilute HCl solution into the beaker and turn on the switch.
3. Observe whether the bulb glows.
4. Repeat the experiment with dilute sulphuric acid (H₂SO₄).
5. Now, repeat the experiment with glucose and alcohol solutions.
6. Observe whether the bulb glows in these cases.
7. Repeat the activity using alkalis (sodium hydroxide, calcium hydroxide, etc.).

Observations:

• The bulb glows when using acids (HCl, H₂SO₄) because they dissociate into ions (H⁺ and Cl⁻, H⁺ and SO₄²⁻), allowing electricity to pass through.
• The bulb does NOT glow when using glucose and alcohol, as they do not produce ions in solution and cannot conduct electricity.
• Bases (alkalis) like NaOH and Ca(OH)₂ also conduct electricity because they produce OH⁻ ions in solution.

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3. Explanation: Why Do Acids and Bases Conduct Electricity?

• Acids ionize in water to produce H⁺ ions, which are responsible for conducting electricity.
  • Example: $$HCl (aq) \rightarrow H^+ (aq) + Cl^- (aq)$$ $$H_2SO_4 (aq) \rightarrow 2H^+ (aq) + SO_4^{2-} (aq)$$
• Bases (alkalis) ionize in water to produce OH⁻ ions, which also allow electricity to flow.
  • Example: $$NaOH (aq) \rightarrow Na^+ (aq) + OH^- (aq)$$ $$Ca(OH)_2 (aq) \rightarrow Ca^{2+} (aq) + 2OH^- (aq)$$

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4. Conclusion

• All acids release H⁺ ions in solution, which are responsible for their acidic properties.
• All bases release OH⁻ ions in solution, which are responsible for their basic properties.
• Acids and bases conduct electricity due to the presence of free ions.
• Glucose and alcohol do NOT conduct electricity because they do not produce ions in solution.

✅ Key Takeaway: The presence of H⁺ ions makes a substance acidic, and the presence of OH⁻ ions makes a substance basic.

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

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1. Do Acids Produce Ions Only in Aqueous Solution?

• Acids release H⁺ (hydrogen ions) only in the presence of water.
• Dry hydrogen chloride gas (HCl) does not exhibit acidic properties.
• HCl solution in water is acidic because it dissociates into ions: $$HCl + H_2O → H_3O^+ + Cl^-$$
• Hydrogen ions (H⁺) cannot exist alone; they combine with water molecules to form hydronium ions (H₃O⁺). $$H^+ + H_2O → H_3O^+$$

Activity 2.9: Testing Acidic Nature of HCl

Procedure:

1. Take about 1g solid NaCl in a test tube and set up the apparatus.
2. Add concentrated H₂SO₄ to the test tube.
3. Observe the gas evolved and test it using dry and wet blue litmus paper.
4. Record the changes in color of litmus paper.

Observations:

• No color change in dry blue litmus paper (HCl gas is not acidic in dry form).
• Color changes to red in wet blue litmus paper (HCl gas dissolves in water to form HCl solution, which is acidic).

Inference:

• Dry HCl gas is not acidic because it does not produce H⁺ ions.
• HCl solution is acidic because it releases H₃O⁺ ions in water.

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2. What Happens When a Base is Dissolved in Water?

• Bases release hydroxide ions (OH⁻) in solution.

• Examples of base dissociation in water:

  • Sodium hydroxide (NaOH): $$NaOH (s) \rightarrow Na^+ (aq) + OH^- (aq)$$
  • Potassium hydroxide (KOH): $$KOH (s) \rightarrow K^+ (aq) + OH^- (aq)$$
  • Magnesium hydroxide (Mg(OH)₂): $$Mg(OH)_2 (s) \rightarrow Mg^{2+} (aq) + 2OH^- (aq)$$

• Bases that are soluble in water are called alkalis.

• All alkalis are bases, but not all bases are alkalis.

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3. Neutralization Reaction

• Acid + Base → Salt + Water $$H^+ (aq) + OH^- (aq) → H_2O (l)$$
• The reaction results in the formation of water and a salt.
• This is the basis of acid-base neutralization.

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4. Effect of Mixing Acids or Bases with Water

• Dissolving an acid or a base in water is a highly exothermic reaction.
• Precaution: Always add acid to water slowly while stirring.
• Never add water to acid, as it may cause splashing and burns.

Activity 2.10: Observing Temperature Change During Dilution

Procedure:

1. Take 10 mL of water in a beaker.
2. Add a few drops of concentrated H₂SO₄ slowly and swirl.
3. Touch the beaker and observe if the temperature changes.
4. Repeat the activity with sodium hydroxide (NaOH) pellets and record observations.

Observations:

• The temperature increases, indicating that the reaction is exothermic.
• Heat is released when acids or bases dissolve in water.

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5. Dilution of Acids and Bases

• Dilution: Adding water to an acid or a base reduces the concentration of H₃O⁺ or OH⁻ ions per unit volume.
• The acid or base is said to be diluted when mixed with water.

✅ Key Takeaway:

• Acids release H⁺ ions in water, bases release OH⁻ ions in water.
• Acid or base must be added to water slowly with stirring to prevent accidents.
• Dilution reduces the concentration of the acid or base, making it less reactive.

Answers to the Questions

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1. Why do HCl, HNO₃, etc., show acidic characters in aqueous solutions while solutions of compounds like alcohol and glucose do not show acidic character?

• HCl and HNO₃ are acids, and when dissolved in water, they ionize to produce H⁺ (or H₃O⁺) ions, which are responsible for their acidic properties: $$HCl → H^+ + Cl^-$$ $$HNO_3 → H^+ + NO_3^-$$
• Alcohol and glucose contain hydrogen atoms, but they do not ionize in water to produce H⁺ ions.
• Since acids produce H⁺ ions in water, they show acidic properties, whereas alcohol and glucose do not release H⁺ ions and remain neutral.

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2. Why does an aqueous solution of an acid conduct electricity?

• Acids dissociate into ions in water, and these ions carry electric current.
• Example: Hydrochloric acid (HCl) in water $$HCl → H^+ + Cl^-$$
• The free-moving H⁺ and Cl⁻ ions in the solution act as charge carriers, making the solution electrically conductive.

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3. Why does dry HCl gas not change the colour of the dry litmus paper?

• Dry HCl gas does not contain H⁺ ions, as ionization happens only in the presence of water.
• Since H⁺ ions are responsible for acidity, the dry gas cannot exhibit acidic properties.
• Dry litmus paper does not contain moisture, so HCl gas cannot dissolve in it and does not change its color.

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4. While diluting an acid, why is it recommended that the acid should be added to water and not water to the acid?

• Dilution is a highly exothermic process (releases heat).
• If water is added to acid, the large amount of heat generated can cause acid to splash out, leading to burns and accidents.
• By adding acid slowly to water, the heat is dissipated gradually, preventing accidents.

✅ Rule to follow: Always add acid to water slowly while stirring.

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5. How is the concentration of hydronium ions (H₃O⁺) affected when a solution of an acid is diluted?

• When an acid is diluted (by adding water), the concentration of H₃O⁺ ions per unit volume decreases.
• Reaction in water: $$HCl + H_2O → H_3O^+ + Cl^-$$
• Since more water molecules are added, the number of H₃O⁺ ions per unit volume decreases, reducing the acidity.

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6. How is the concentration of hydroxide ions (OH⁻) affected when excess base is dissolved in a solution of sodium hydroxide?

• Adding more base (NaOH) increases the OH⁻ ion concentration in the solution.
• Dissociation of NaOH in water: $$NaOH → Na^+ + OH^-$$
• When more NaOH is added, more OH⁻ ions are released, increasing the solution’s alkalinity.

✅ Key Point: More base → Higher OH⁻ concentration → Stronger alkaline nature.

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 on pH and Strength of Acids & Bases

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1. What is pH?

• pH is a scale used to measure the concentration of hydrogen ions (H⁺) in a solution.
• The p in pH stands for "potenz", meaning power in German.
• The pH scale generally ranges from 0 to 14.

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2. pH Scale and Its Interpretation

• pH = 7 → Neutral solution (e.g., pure water).
• pH < 7 → Acidic solution (More H⁺ ions).
• pH > 7 → Basic (Alkaline) solution (More OH⁻ ions).
• As the concentration of H₃O⁺ increases, pH decreases, making the solution more acidic.
• As the concentration of OH⁻ increases, pH increases, making the solution more basic.

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3. Universal Indicator and pH Paper

• A universal indicator is a mixture of several indicators that show different colors at different pH values.
• pH paper is impregnated with a universal indicator and changes color based on pH.
• Color variations on pH paper (approximate values):
  • Red (0-3): Strong acid (e.g., HCl, H₂SO₄).
  • Orange/Yellow (4-6): Weak acid (e.g., acetic acid, citric acid).
  • Green (7): Neutral (e.g., pure water).
  • Blue (8-10): Weak base (e.g., ammonia, baking soda).
  • Dark blue/Purple (11-14): Strong base (e.g., NaOH, KOH).

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4. Strong vs. Weak Acids & Bases

• Strong acids: Completely ionize in water, producing more H⁺ ions.
  • Example: HCl → H⁺ + Cl⁻ (Hydrochloric acid).
  • Example: H₂SO₄ → 2H⁺ + SO₄²⁻ (Sulfuric acid).
• Weak acids: Partially ionize, producing fewer H⁺ ions.
  • Example: CH₃COOH ⇌ CH₃COO⁻ + H⁺ (Acetic acid).
• Strong bases: Completely dissociate, releasing more OH⁻ ions.
  • Example: NaOH → Na⁺ + OH⁻ (Sodium hydroxide).
  • Example: KOH → K⁺ + OH⁻ (Potassium hydroxide).
• Weak bases: Partially dissociate, releasing fewer OH⁻ ions.
  • Example: NH₄OH ⇌ NH₄⁺ + OH⁻ (Ammonium hydroxide).

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5. Activity 2.11: Testing pH Values of Common Substances

S. No.	Solution	Color on pH Paper	Approximate pH	Nature
1	Saliva (before meal)	Greenish	7.4	Slightly basic
2	Saliva (after meal)	Yellowish	6.5	Slightly acidic
3	Lemon juice	Reddish	2-3	Strongly acidic
4	Colourless aerated drink	Orange-red	3-4	Acidic
5	Carrot juice	Yellowish-green	5-6	Weakly acidic
6	Coffee	Yellowish-brown	5	Weakly acidic
7	Tomato juice	Orange	4.2	Acidic
8	Tap water	Green	7	Neutral
9	1M NaOH	Blue	14	Strongly basic
10	1M HCl	Red	1	Strongly acidic

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6. Summary

• The pH scale measures how strong an acid or base is based on H⁺ or OH⁻ concentration.
• Lower pH (<7) = Acidic, Higher pH (>7) = Basic, pH = 7 = Neutral.
• Acids and bases differ in strength based on how much they ionize in water.
• Universal indicators and pH papers help determine the pH of substances.

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 pH and Its Importance in Everyday Life

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2.3.1 Importance of pH in Everyday Life

1. Are Plants and Animals pH Sensitive?

• The human body functions within a pH range of 7.0 to 7.8.
• Living organisms can survive only in a narrow pH range.
• pH variations can be harmful:
  • If the pH of blood drops or rises, it can cause severe health issues.
  • Aquatic life is affected when water pH drops below the ideal range due to acid rain or pollution.

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2. Acid Rain and Its Effects

• Acid rain occurs when the pH of rainwater drops below 5.6 due to the presence of acidic gases like SO₂ (sulfur dioxide) and NO₂ (nitrogen dioxide) in the atmosphere.
• Effects of acid rain:
  • Damages buildings and monuments, especially those made of marble (CaCO₃).
  • Lowers the pH of rivers and lakes, making it difficult for aquatic life to survive.
  • Affects soil fertility, reducing crop yield.

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3. pH and Soil for Plant Growth

• Different plants require different soil pH levels for optimal growth.
• Ideal pH for plants:
  • Neutral to slightly acidic (pH 6-7) for most crops.
  • Acidic soil (pH < 6) for plants like tea, potatoes, and strawberries.
  • Alkaline soil (pH > 7) for plants like beets and cabbage.
• Farmers test soil pH to adjust it using fertilizers:
  • If soil is too acidic, they add quicklime (CaO) or slaked lime (Ca(OH)₂).
  • If soil is too basic, they add organic matter (manure or compost) to neutralize it.

Activity 2.12: Testing Soil pH

1. Take 2g of soil in a test tube and add 5mL of water.
2. Shake well and filter the mixture into another test tube.
3. Test the pH of the filtrate using universal indicator paper.
4. Observe and determine the ideal pH range for plant growth in your area.

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4. Acids in Other Planets

• Venus’ atmosphere consists of thick white and yellowish clouds of sulfuric acid (H₂SO₄).
• Extremely acidic conditions make life impossible on Venus.

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5. pH in the Human Digestive System

• The stomach produces hydrochloric acid (HCl) to help digest food.
• Indigestion occurs when excess acid is produced, causing pain and irritation (acidity).
• To neutralize this acid, antacids (mild bases) are used, such as:
  • Magnesium hydroxide (Milk of Magnesia, Mg(OH)₂).
  • Sodium bicarbonate (Baking soda, NaHCO₃).

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6. pH and Tooth Decay

• Tooth enamel is made of calcium hydroxyapatite (Ca₁₀(PO₄)₆(OH)₂), which is corroded when pH falls below 5.5.
• Bacteria in the mouth break down sugar and produce acids.
• This lowers pH, leading to tooth decay.
• Prevention:
  • Brushing teeth with toothpaste, which is basic, neutralizes the acid.
  • Rinsing mouth after eating to remove food particles.

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7. Self-Defense Mechanisms of Animals and Plants (Chemical Warfare)

Bee Sting

• Bee sting injects formic acid (methanoic acid), causing pain and irritation.
• Applying a mild base like baking soda (NaHCO₃) neutralizes the acid, providing relief.

Nettle Plant Sting

• Nettle leaves contain stinging hairs that inject methanoic acid, causing a burning sensation.
• The traditional remedy is rubbing the area with dock leaves, which contain a mild base that neutralizes the acid.

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8. Naturally Occurring Acids

Natural Source	Acid Present
Vinegar	Acetic acid (CH₃COOH)
Curd (Sour milk)	Lactic acid
Orange	Citric acid
Lemon	Citric acid
Tamarind	Tartaric acid
Ant sting	Methanoic acid (Formic acid, HCOOH)
Tomato	Oxalic acid
Nettle sting	Methanoic acid (Formic acid, HCOOH)

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Summary

• Living organisms require a stable pH for survival (7.0 - 7.8 for humans).
• Acid rain (pH < 5.6) harms water bodies, plants, and buildings.
• Soil pH affects plant growth, and farmers adjust it using lime (for acidic soil) or compost (for basic soil).
• The stomach produces HCl for digestion, but too much acid causes indigestion.
• Tooth decay occurs when mouth pH drops below 5.5 due to bacterial acid production.
• Animals and plants use acids for self-defense (bee sting, nettle sting).
• Many common foods and natural substances contain organic acids.

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Questions and Answers 

1. You have two solutions, A and B. The pH of solution A is 6 and pH of solution B is 8. Which solution has more hydrogen ion concentration? Which of these is acidic and which one is basic?

   • Solution A (pH = 6) has a higher hydrogen ion (H⁺) concentration than Solution B (pH = 8).
   • Solution A is acidic (since pH is less than 7), and Solution B is basic (since pH is greater than 7).

2. What effect does the concentration of H⁺ (aq) ions have on the nature of the solution?

   • Higher H⁺ ion concentration makes the solution more acidic (pH < 7).
   • Lower H⁺ ion concentration makes the solution more basic (pH > 7).

3. Do basic solutions also have H⁺ (aq) ions? If yes, then why are these basic?

   • Yes, basic solutions also contain H⁺ ions, but in very small amounts.
   • Basic solutions have a higher concentration of hydroxide ions (OH⁻) than H⁺ ions, which makes them basic.
   • The pH of a basic solution is greater than 7 because the OH⁻ ions dominate over the H⁺ ions.

4. Under what soil condition do you think a farmer would treat the soil of his fields with quick lime (CaO), slaked lime (Ca(OH)₂), or chalk (CaCO₃)?

   • A farmer would treat the soil with these substances if the soil is too acidic (low pH).
   • Quick lime (CaO), slaked lime (Ca(OH)₂), and chalk (CaCO₃) are basic substances that neutralize excess acid in the soil.
   • This helps in maintaining an optimal pH for crop growth, ensuring better yield.

Salts 

What are Salts?

• Salts are ionic compounds formed by the neutralization reaction of an acid and a base.
• They consist of positive ions (cations) from bases and negative ions (anions) from acids.

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2.4.1 Family of Salts

Activity 2.13

1. Writing Chemical Formulae of Given Salts

Salt Name	Chemical Formula
Potassium sulfate	K₂SO₄
Sodium sulfate	Na₂SO₄
Calcium sulfate	CaSO₄
Magnesium sulfate	MgSO₄
Copper sulfate	CuSO₄
Sodium chloride	NaCl
Sodium nitrate	NaNO₃
Sodium carbonate	Na₂CO₃
Ammonium chloride	NH₄Cl

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2. Identifying the Acids and Bases from Which These Salts are Formed

Salt	Acid	Base
Potassium sulfate (K₂SO₄)	Sulfuric acid (H₂SO₄)	Potassium hydroxide (KOH)
Sodium sulfate (Na₂SO₄)	Sulfuric acid (H₂SO₄)	Sodium hydroxide (NaOH)
Calcium sulfate (CaSO₄)	Sulfuric acid (H₂SO₄)	Calcium hydroxide (Ca(OH)₂)
Magnesium sulfate (MgSO₄)	Sulfuric acid (H₂SO₄)	Magnesium hydroxide (Mg(OH)₂)
Copper sulfate (CuSO₄)	Sulfuric acid (H₂SO₄)	Copper hydroxide (Cu(OH)₂)
Sodium chloride (NaCl)	Hydrochloric acid (HCl)	Sodium hydroxide (NaOH)
Sodium nitrate (NaNO₃)	Nitric acid (HNO₃)	Sodium hydroxide (NaOH)
Sodium carbonate (Na₂CO₃)	Carbonic acid (H₂CO₃)	Sodium hydroxide (NaOH)
Ammonium chloride (NH₄Cl)	Hydrochloric acid (HCl)	Ammonium hydroxide (NH₄OH)

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3. Identifying Salt Families

• Salts are classified into families based on having the same cation (positive ion) or anion (negative ion).
• Example of Salt Families:
  • Sodium salts: Na₂SO₄, NaCl, NaNO₃, Na₂CO₃ (contain Na⁺)
  • Sulphate salts: K₂SO₄, Na₂SO₄, CaSO₄, MgSO₄, CuSO₄ (contain SO₄²⁻)
  • Chloride salts: NaCl, NH₄Cl (contain Cl⁻)

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Key Points to Remember

1. Salts are formed by neutralization of acids and bases.
2. Cations come from bases, and anions come from acids.
3. Salts are classified into families based on their common positive or negative ions.
4. Different salts have different properties and uses, depending on their composition.

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2.4.2 pH of Salts

Activity 2.14: Testing pH of Salts

1. Collect the following salt samples:

   • Sodium chloride (NaCl)
   • Potassium nitrate (KNO₃)
   • Aluminium chloride (AlCl₃)
   • Zinc sulphate (ZnSO₄)
   • Copper sulphate (CuSO₄)
   • Sodium acetate (CH₃COONa)
   • Sodium carbonate (Na₂CO₃)
   • Sodium hydrogencarbonate (NaHCO₃)

2. Check solubility in distilled water.

3. Test the solution’s action on litmus paper.

4. Find the pH using pH paper.

Classification of Salts Based on pH

• Neutral Salts (pH = 7): Formed from strong acid + strong base
  • Example: Sodium chloride (NaCl), Potassium nitrate (KNO₃)
• Acidic Salts (pH < 7): Formed from strong acid + weak base
  • Example: Aluminium chloride (AlCl₃), Zinc sulphate (ZnSO₄), Copper sulphate (CuSO₄)
• Basic Salts (pH > 7): Formed from strong base + weak acid
  • Example: Sodium acetate (CH₃COONa), Sodium carbonate (Na₂CO₃), Sodium hydrogencarbonate (NaHCO₃)

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2.4.3 Chemicals from Common Salt

Sources of Common Salt (NaCl)

• Obtained from seawater or rock salt deposits.
• Rock salt is often brown due to impurities.
• Historically significant in India’s freedom struggle (Dandi March by Mahatma Gandhi).

Common Salt as a Raw Material for Chemicals

NaCl is used to produce various important chemicals:

Chemical	Process of Production	Chemical Equation	Uses
Sodium Hydroxide (NaOH)	Chlor-alkali process (electrolysis of brine)	2NaCl + 2H₂O → 2NaOH + Cl₂ + H₂	Used in soap, paper, textiles, and detergents
Bleaching Powder (CaOCl₂)	Action of Cl₂ on slaked lime (Ca(OH)₂)	Ca(OH)₂ + Cl₂ → CaOCl₂ + H₂O	Disinfectant, bleaching agent in textile & paper industries
Baking Soda (NaHCO₃)	Reaction of NaCl, H₂O, CO₂, and NH₃	NaCl + H₂O + CO₂ + NH₃ → NH₄Cl + NaHCO₃	Cooking, antacids, fire extinguishers
Washing Soda (Na₂CO₃.10H₂O)	Heating & recrystallization of baking soda	NaHCO₃ → Na₂CO₃ + H₂O + CO₂	Glass, soap, paper industries, water softening

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1. Sodium Hydroxide (NaOH) – Chlor-Alkali Process

• Electrolysis of brine (saltwater solution).
• Products formed:
  • Sodium hydroxide (NaOH) – near the cathode
  • Chlorine gas (Cl₂) – at the anode
  • Hydrogen gas (H₂) – at the cathode
• Chemical equation: $$2NaCl(aq) + 2H_2O(l) → 2NaOH(aq) + Cl_2(g) + H_2(g)$$
• Uses:
  • Soap and detergent manufacturing
  • Paper and textile industries
  • Water purification

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2. Bleaching Powder (CaOCl₂)

• Made by reacting chlorine gas (Cl₂) with slaked lime (Ca(OH)₂).
• Chemical equation: $$Ca(OH)_2 + Cl_2 → CaOCl_2 + H_2O$$
• Uses:
  • Bleaching cotton, linen, paper, and clothes
  • Disinfecting drinking water
  • Oxidizing agent in chemical industries

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3. Baking Soda (NaHCO₃)

• Also called Sodium Hydrogen Carbonate.

• Produced from:

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

• pH is slightly basic.

• Uses:

  • Baking (used in baking powder with tartaric acid to release CO₂ gas, making cakes soft & spongy)
  • Antacids (neutralizes stomach acid)
  • Fire extinguishers (releases CO₂ when heated)

• Reaction when heated:

  $$2NaHCO_3 → Na_2CO_3 + H_2O + CO_2$$

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4. Washing Soda (Na₂CO₃.10H₂O)

• Formed by recrystallization of sodium carbonate.

• Chemical equation:

  $$Na_2CO_3 + 10H_2O → Na_2CO_3.10H_2O$$

• Uses:

  • Glass, soap, and paper industries
  • Manufacturing of borax
  • Used as a cleaning agent
  • Removes permanent hardness of water

• Why 10H₂O?

  • It is a hydrated salt.
  • The water of crystallization (10H₂O) gives it a crystalline structure but does not make it wet.

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Key Points to Remember

• pH of salts depends on the nature of the acid and base that formed them.
• Chlor-alkali process is used to produce NaOH, Cl₂, and H₂.
• Common salt (NaCl) is the raw material for several industrially important chemicals.
• Bleaching powder, baking soda, and washing soda have important household and industrial applications.

 Water of Crystallisation 

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2.4.4 Are the Crystals of Salts Really Dry?

Activity 2.15: Observing Water of Crystallisation

1. Heat a few crystals of copper sulphate (CuSO₄·5H₂O) in a dry boiling tube.
2. Observe the colour change.
   • Before heating: Blue crystals (hydrated copper sulphate).
   • After heating: White powder (anhydrous copper sulphate).
3. Notice water droplets inside the tube.
   • These droplets come from the water of crystallisation present in the salt.
4. Add a few drops of water to the white powder.
   • Observation: The blue colour reappears, indicating that the salt absorbs water back into its structure.

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What is Water of Crystallisation?

• Definition: The fixed number of water molecules chemically bound in a salt’s crystal structure.

• Example: Copper sulphate (CuSO₄·5H₂O)

  • 5H₂O molecules are attached per formula unit.
  • When heated, it loses water and turns white:
•  
  
  $$\mathrm{<br>}$$
• Adding water restores the blue colour.

• Other Examples:

  • Gypsum (CaSO₄·2H₂O) – Contains 2 water molecules per unit.
  • Washing soda (Na₂CO₃·10H₂O) – Contains 10 water molecules per unit.

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Plaster of Paris (POP)

• Derived from Gypsum (CaSO₄·2H₂O)
• Formation:
  • Heating gypsum at 373 K removes part of its water and forms Plaster of Paris (CaSO₄·½H₂O).
  

  
  
  $$CaSO_4·2H_2O \xrightarrow{\text{heat at 373K}} CaSO_4·\frac{1}{2}H_2O + 1\frac{1}{2}H_2O$$
• When mixed with water, it sets into hard gypsum again:
• $$CaSO_4·\frac{1}{2}H_2O + \frac{3}{2}H_2O → CaSO_4·2H_2O$$

Uses of Plaster of Paris

• Medical use: Supports fractured bones.
• Construction: Smooths walls and ceilings.
• Art & Decoration: Used in making toys, statues, and decorative materials.

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Why is it Called "Plaster of Paris"?

• Named after Paris, France, where it was first manufactured from large gypsum deposits.

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Key Takeaways

• Hydrated salts contain water of crystallisation, making them appear dry.
• Heating removes this water, changing the salt’s colour and structure.
• Plaster of Paris is made from gypsum and hardens when mixed with water.
• Water of crystallisation is responsible for the physical appearance of many salts.

Questions and Answer 

1. What is the common name of the compound CaOCl₂?

   • Answer: Bleaching powder

2. Name the substance which on treatment with chlorine yields bleaching powder.

   • Answer: Slaked lime (Calcium hydroxide) – Ca(OH)₂
   • Reaction: $$Ca(OH)_2 + Cl_2 → CaOCl_2 + H_2O$$

3. Name the sodium compound which is used for softening hard water.

   • Answer: Washing soda (Sodium carbonate – Na₂CO₃·10H₂O)

4. What will happen if a solution of sodium hydrogencarbonate is heated? Give the equation of the reaction involved.

   • Answer: It decomposes to form sodium carbonate, water, and carbon dioxide gas.
   • Reaction: $$$$
     

5. Write an equation to show the reaction between Plaster of Paris and water.

• Answer: When mixed with water, Plaster of Paris (CaSO₄·½H₂O) absorbs water and forms gypsum (CaSO₄·2H₂O), which hardens.
• Reaction: $$CaS{O}_4\cdot \frac{1}{2}{H}_{2}O+\frac{3}{2}{H}_{2}O\to CaS{O}_4\cdot 2H_{2}O$$

Exercise 

Questions and Answers:

1. A solution turns red litmus blue; its pH is likely to be:
   (a) 1
   (b) 4
   (c) 5
   (d) 10

   Answer: (d) 10

   • A solution that turns red litmus blue is basic. A pH of 10 indicates a moderately strong base.

2. A solution reacts with crushed egg-shells to give a gas that turns lime-water milky. The solution contains:
   (a) NaCl
   (b) HCl
   (c) LiCl
   (d) KCl

   Answer: (b) HCl

   • Eggshells contain calcium carbonate (CaCO₃), which reacts with HCl to form carbon dioxide (CO₂). CO₂ turns limewater milky due to the formation of calcium carbonate precipitate.
   • Reaction: $$CaCO_3 + 2HCl → CaCl_2 + H_2O + CO_2$$

3. 10 mL of a solution of NaOH is found to be completely neutralised by 8 mL of a given solution of HCl. If we take 20 mL of the same solution of NaOH, the amount of HCl solution (the same solution as before) required to neutralise it will be:
   (a) 4 mL
   (b) 8 mL
   (c) 12 mL
   (d) 16 mL

   Answer: (d) 16 mL

   • Since 10 mL of NaOH requires 8 mL of HCl, doubling NaOH to 20 mL will require double the amount of HCl, i.e., 16 mL.

4. Which one of the following types of medicines is used for treating indigestion?
   (a) Antibiotic
   (b) Analgesic
   (c) Antacid
   (d) Antiseptic

   Answer: (c) Antacid

   • Indigestion is caused by excess stomach acid. Antacids, such as magnesium hydroxide (Mg(OH)₂) and sodium bicarbonate (NaHCO₃), neutralize excess acid and relieve discomfort.

5. Write word equations and then balanced equations for the reaction taking place when –
   (a) Dilute sulphuric acid reacts with zinc granules.
   Word Equation: Zinc + Sulphuric acid → Zinc sulphate + Hydrogen gas
   Balanced Equation:

   $$Zn + H_2SO_4 → ZnSO_4 + H_2$$

   (b) Dilute hydrochloric acid reacts with magnesium ribbon.
   Word Equation: Magnesium + Hydrochloric acid → Magnesium chloride + Hydrogen gas
   Balanced Equation:

   $$Mg + 2HCl → MgCl_2 + H_2$$

   (c) Dilute sulphuric acid reacts with aluminium powder.
   Word Equation: Aluminium + Sulphuric acid → Aluminium sulphate + Hydrogen gas
   Balanced Equation:

   $$2Al + 3H_2SO_4 → Al_2(SO_4)_3 + 3H_2$$

   (d) Dilute hydrochloric acid reacts with iron filings.
   Word Equation: Iron + Hydrochloric acid → Iron chloride + Hydrogen gas
   Balanced Equation:

   $$Fe + 2HCl → FeCl_2 + H_2$$

6. Compounds such as alcohols and glucose also contain hydrogen but are not categorised as acids. Describe an Activity to prove it.
   Answer:

   • Activity:
     • Set up a circuit with a bulb, battery, and electrodes in a beaker of water.
     • Add dilute HCl → Bulb glows (HCl releases H⁺ ions, conducting electricity).
     • Add glucose/alcohol solution → Bulb does not glow (No ionization, no conduction).
   • Conclusion: Acids release H⁺ ions in water, while glucose and alcohol do not, so they are not acids.

7. Why does distilled water not conduct electricity, whereas rainwater does?
   Answer:

   • Distilled water has no ions, so it does not conduct electricity.
   • Rainwater contains dissolved CO₂ and minerals, forming ions that conduct electricity.

8. Why do acids not show acidic behaviour in the absence of water?
   Answer:

   • Acids release H⁺ ions only in water. Without water, they remain in their molecular form and do not exhibit acidity.

9. Five solutions A, B, C, D, and E showed pH values of 4, 1, 11, 7, and 9, respectively. Identify:

   • (a) Neutral: D (pH = 7)
   • (b) Strongly alkaline: C (pH = 11)
   • (c) Strongly acidic: B (pH = 1)
   • (d) Weakly acidic: A (pH = 4)
   • (e) Weakly alkaline: E (pH = 9)
   • Increasing order of hydrogen-ion concentration: C < E < D < A < B

10. Equal lengths of magnesium ribbons are taken in test tubes A and B. Hydrochloric acid (HCl) is added to test tube A, while acetic acid (CH₃COOH) is added to test tube B. Amount and concentration are the same. In which test tube will the fizzing occur more vigorously and why?
    Answer:

• Fizzing is more vigorous in Test Tube A (HCl) because HCl is a strong acid, producing more H⁺ ions. Acetic acid is a weak acid, so the reaction is slower.

11. Fresh milk has a pH of 6. How do you think the pH will change as it turns into curd? Explain your answer.
    Answer:

• As milk turns into curd, lactic acid forms, making it more acidic. The pH decreases from 6 to below 6.

12. A milkman adds a very small amount of baking soda to fresh milk.
    (a) Why does he shift the pH of the fresh milk from 6 to slightly alkaline?

• Answer: Baking soda (NaHCO₃) is a mild base, which neutralizes acidity and prevents milk from curdling quickly.

(b) Why does this milk take a long time to set as curd?

• Answer: A higher pH slows down the action of lactic acid bacteria, delaying curdling.

13. Plaster of Paris should be stored in a moisture-proof container. Explain why.
    Answer:

• Plaster of Paris (CaSO₄·½H₂O) absorbs moisture, turning into gypsum (CaSO₄·2H₂O), which hardens and becomes unusable.

14. What is a neutralisation reaction? Give two examples.
    Answer:

• Neutralisation: Reaction of acid + base → salt + water
• Examples:
  • HCl + NaOH → NaCl + H₂O
  • H₂SO₄ + Mg(OH)₂ → MgSO₄ + 2H₂O

15. Give two important uses of washing soda and baking soda.
    Answer:

• Washing Soda (Na₂CO₃·10H₂O):
  1. Softens hard water
  2. Used in glass and soap manufacturing
• Baking Soda (NaHCO₃):
  1. Used in baking (CO₂ release makes dough fluffy)
  2. Used as an antacid to relieve acidity