NCERT Science Notes - Class 10
Chapter 2 - Acids, Bases and Salts

Welcome to our AJs Chalo Seekhen. This webpage is dedicated to Class 10 | Science | Chapter 2 - Acids, Bases and Salts. This chapter provides an insightful exploration of these crucial chemical substances. It introduces students to the fundamental concepts of acids and bases, their properties, and the reactions they undergo. The chapter also delves into the formation and characteristics of salts, a product of reactions between acids and bases. This comprehensive overview is designed to build a strong foundation in understanding these essential components of chemistry, making it a vital part of the curriculum for Class 10th students.

Class 10 NCERT Acids bases and salts notes ajs, cbse notes class 10 ajslearning, cbse notes ajs, ajs notes class 10, ajslearning, ajs chalo seekhen

NCERT Science Notes - Class 10
Chapter 2 - Acids, Bases and Salts

1. Sour and Bitter Tastes: Acids cause sour taste; bases cause bitter taste.
2. Remedy for Acidity: Baking soda solution is suggested due to its basic nature.
3. Acid-Base Neutralization: Acids and bases can nullify each other’s effects.
4. Indicators:
   • Litmus: Natural indicator; blue litmus turns red in acid, red litmus turns blue in base.
   • Turmeric: Turns reddish-brown with bases.
   • Methyl Orange and Phenolphthalein: Synthetic indicators.
5. Litmus Solution: Purple dye from lichen; purple when neutral.
6. Other Natural Indicators: Red cabbage leaves, turmeric, certain flower petals.

2.1.2 How do Acids and Bases React with Metals?

Activity 2.3

Activity 2.5

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

1. Objective: To observe the reaction of Sodium Carbonate (Na₂CO₃) and Sodium Hydrogencarbonate (NaHCO₃) with dilute Hydrochloric Acid (HCl).
2. Procedure:
   1. Sodium Carbonate is placed in test tube A and Sodium Hydrogencarbonate in test tube B.
   2. 2 mL of dilute HCl is added to both test tubes.
   3. The gas produced in each test tube is passed through lime water.
3. Chemical Reactions:
   1. Test tube A: Na₂CO₃ + 2HCl(aq) → 2NaCl + H₂O(l) + CO₂(g)
   2. Test tube B: NaHCO₃(s) + HCl(aq) → NaCl(aq) + H₂O(l) + CO₂(g)
   3. Reaction with lime water: Ca(OH)2(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)
   4. Reaction with excess CO₂: CaCO₃(s) + H₂O(l) + CO₂(g) → Ca(HCO₃)2(aq)
4. Observations:
   1. Bubbling observed in both test tubes, indicating the release of CO₂ gas.
   2. Passing the gas through lime water turns it milky due to the formation of calcium carbonate.
   3. Further passage of CO₂ dissolves the precipitate, indicating the formation of calcium hydrogencarbonate.
5. Conclusion: Metal carbonates and hydrogencarbonates react with acids to form a salt, carbon dioxide, and water.
6. Forms of Calcium Carbonate: Limestone, chalk, and marble are different forms of calcium carbonate, each with distinct physical properties but the same chemical composition: CaCO₃.
7. Reaction of Metal Carbonates and Hydrogencarbonates with Acids
   1. General Reaction: Both metal carbonates and metal hydrogencarbonates react with acids to form a corresponding salt, carbon dioxide, and water. The general reaction can be written as:
      Metal carbonate/Metal hydrogencarbonate + Acid → Salt + Carbon dioxide + Water
   2. Example Reaction: A typical example is the reaction of Sodium Carbonate (Na₂CO₃) or Sodium Hydrogencarbonate (NaHCO₃) with Hydrochloric Acid (HCl):
   3. Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂ (using Sodium Carbonate)
   4. NaHCO₃ + HCl → NaCl + H₂O + CO₂ (using Sodium Hydrogencarbonate)

2.1.4 How do Acids and Bases React with each other?

1. Objective of Activity 2.6: To demonstrate how acids and bases react with each other, specifically observing the neutralization reaction.
2. Experiment Setup:
   1. Begins with dilute NaOH (a base) in a test tube.
   2. Addition of phenolphthalein, an indicator that turns pink in basic solutions.
3. Observations and Interpretations:
   • Initial pink color of the solution: Indicates the basic nature of the NaOH solution.
   • Gradual addition of dilute HCl (an acid): The pink color fades, demonstrating the neutralization of the base by the acid.
   • The color change of phenolphthalein: Phenolphthalein changes color based on the pH of the solution. It's pink in basic solutions and colorless in acidic environments. The fading of the pink color indicates the shift of the solution's pH from basic towards neutral.
   • Re-addition of NaOH: The reappearance of the pink color upon adding more NaOH suggests that the solution's pH has shifted back to basic.
4. Chemical Reaction Involved:

   The reaction of NaOH with HCl is a classic example of a neutralization reaction where an acid (HCl) and a base (NaOH) react to form a salt (NaCl) and water (H2O).

   The general equation for a neutralization reaction is: Base + Acid → Salt + Water.

   Chemical Equation: NaOH (aq) + HCl (aq) → NaCl (aq) + H₂O (l)

5. Concept of Neutralization:

   Neutralization is the process where an acid and a base react to form water and a salt, effectively neutralizing each other's properties. This reaction is important in various chemical processes and in everyday life, such as in medicine and environmental management.

6. Indicator's Role:

   Phenolphthalein serves as an indicator to visually demonstrate the pH change during the reaction. It's a practical tool to understand the acid-base balance in a given solution.

7. Application and Significance:

   This activity illustrates the fundamental principle of acid-base chemistry and the practical use of indicators. Understanding neutralization is crucial in various fields, including chemistry, biology, medicine, and environmental science. This experiment effectively demonstrates the neutralizing effect of acids and bases on each other and the practical application of indicators in visualizing these reactions.

Decomposition Reaction

Let's write and balance the chemical equation for the reaction described in Activity 2.7, where copper oxide reacts with hydrochloric acid.

Given Reaction:

• Copper oxide ( $\text{CuO}$) reacts with dilute hydrochloric acid ( $\text{HCl}$) to form copper(II) chloride ( ${\mathrm{<span\; style="font-size:\; 1.07rem;">CuCl2)\; and\; water\; (H₂O)</span>}}_{}$

Unbalanced Chemical Equation: CuO (s) + HCl (aq) → CuCl₂ ​(aq) + H₂​O (l) $\text{<span style="font-size: 21.4px;"><br></span> <span style="font-size: 21.4px;"><b>Balancing the Equation:</b></span>}$

1. Copper (Cu): One atom on both sides – balanced.
2. Oxygen (O): One atom on both sides – balanced.
3. Chlorine (Cl): There are two chlorine atoms in ${\text{CuCl}}_2$but only one in $\text{HCl}$. Therefore, we need to have two $\text{HCl}$ molecules to balance chlorine.
4. Hydrogen (H): After balancing chlorine, we will have two hydrogen atoms from the two $\text{HCl}$ molecules, which are balanced with the two hydrogen atoms in H₂O.

Reaction of a Non-metallic Oxide with Base

1. Example Reaction (Activity 2.5): The reaction between carbon dioxide (CO₂, a non-metallic oxide) and calcium hydroxide (lime water, a base) is used as an example.
2. Chemical Process:
   • Reaction Formula: Ca(OH)₂ (aq) + CO₂ (g) → CaCO₃ (s) + H₂O (l)
   • This reaction involves calcium hydroxide reacting with carbon dioxide to produce calcium carbonate (a salt) and water.
3. Conclusion on Non-metallic Oxides:
   • Non-metallic oxides, like CO₂, are acidic in nature as they react with bases to form salts and water, similar to the reaction of acids.
4. Significance:
   • Understanding the behavior of non-metallic oxides is crucial for various chemical processes, environmental studies, and industrial applications.
   • This knowledge assists in categorizing compounds as acidic or basic, fundamental in the study of chemistry.

What Do All Acids and All Bases Have in Common?

Activity 2.8 : Overview

1. Objective: To determine if compounds containing hydrogen are acidic.
2. Experiment Setup: Solutions of glucose, alcohol, hydrochloric acid (HCl), and sulphuric acid (H₂SO₄) were tested. Two nails were fixed on a cork in a beaker and connected to a 6-volt battery through a bulb and switch.
3. Procedure:
   • Dilute HCl and sulphuric acid were poured into the beaker, and the current was switched on.
   • The experiment was repeated separately with glucose and alcohol solutions.
4. Observations:
   • The bulb glowed with acid solutions, indicating the flow of electric current.
   • Glucose and alcohol solutions did not conduct electricity, as the bulb did not glow.
5. Conclusion: The glowing bulb in the presence of acids suggests that acids conduct electricity due to ionization. Acids contain H⁺ ions as cations, which are responsible for their acidic properties.

The experiment concludes that the presence of H⁺ ions in a solution is indicative of acidic properties, differentiating acids from non-acidic compounds like glucose and alcohol.

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

1. Common Feature in Acids:
   • All acids generate hydrogen gas on reacting with metals, indicating that hydrogen (H⁺) ions are common to all acids.
   • In Activity 2.8, when acids like hydrochloric acid (HCl) and sulfuric acid (H₂SO₄) are tested, they conduct electricity, as evidenced by the glowing of a bulb in an electrical setup. This conductivity is due to the presence of ions in the solution.
2. Acidic Solutions and Conductivity:
   • The conductivity of acidic solutions is due to the dissociation of acids into their respective ions in solution. For example, HCl dissociates into H⁺ and Cl⁻ ions.
   • This ionization is responsible for the acidic properties of these compounds.
3. Testing Non-Acidic Compounds:
   • When solutions of non-acidic compounds like glucose and alcohol are tested, the bulb does not glow, indicating no electrical conductivity. This is because these solutions do not ionize into charged particles in water.
4. Experiment with Bases:
   • Repeating the same activity with alkalis (bases) like sodium hydroxide (NaOH) or calcium hydroxide (Ca(OH)₂) will show that the bulb glows, indicating that bases also conduct electricity in solution.
5. Conclusion from Activity:
   • The common feature in acids is the production of hydrogen ions (H⁺) in solution, leading to their acidic properties.
   • Bases, on the other hand, generate hydroxide ions (OH⁻) in solution.
   • Both acids and bases conduct electricity in their aqueous solutions due to the presence of these ions.

Acids, Bases, and their Behavior in Water Solutions

1. Ion Production in Acids and Bases:
   • Acids and bases produce ions in aqueous solutions.
   • Acids generate hydrogen ions (H⁺) in water, which exist as hydronium ions (H₃O⁺).
   • Bases generate hydroxide ions (OH⁻) in water.
   • Example: HCl + H₂O → H3O⁺ + Cl^–
2. Activity 2.9: Acid in a Non-Aqueous Environment:
   • Testing dry HCl gas with both dry and wet blue litmus paper shows that acidic character is prominent only in the presence of water.
   • This suggests that hydrogen ions in HCl are produced only in the presence of water.
3. Alkalis: Bases soluble in water:
   • Not all bases dissolve in water. Those that do are called alkalis (e.g., NaOH, KOH).
   • Alkalis are soapy to touch, bitter, and corrosive.
4. Neutralization Reaction:
   • The neutralization reaction can be viewed as the reaction of H⁺ ions from the acid and OH⁻ ions from the base to form water (H₂O).
   • Example: ${}^{}$

     Acid + Base → Salt + Water
     H [ X + M ] OH → MX + HOH
     H ⁺ (aq) + OH ^– (aq) → H ₂O(l)

Dissolving Acids and Bases in Water

2.3 HOW STRONG ARE ACID OR BASE SOLUTIONS?

pH Scale and Universal Indicator:

1. pH Scale:
   • The pH scale measures the concentration of hydrogen ions in a solution.
   • It ranges from 0 (very acidic) to 14 (very alkaline), with 7 being neutral.
   • The term 'pH' comes from 'potenz' in German, meaning power or potential.
   • Lower pH values indicate higher hydronium ion concentration and hence more acidic solutions.
   • Higher pH values indicate higher OH⁻ ion concentration, suggesting stronger alkaline solutions.
2. Universal Indicator:
   • A mixture of several indicators that shows different colors at different hydrogen ion concentrations.
   • It's commonly used for measuring pH, often impregnated in paper strips.

Activity 2.11: Testing pH Values

• Objective: To test and record the pH values of various solutions and determine their acidic or basic nature.
• Method: Using pH paper impregnated with a universal indicator, the pH of different substances like saliva, lemon juice, coffee, tap water, and solutions of NaOH and HCl is tested.
• Expected Outcome: The color change on the pH paper will indicate the pH value, helping to classify each substance as acidic, basic, or neutral.

Strength of Acids and Bases:

• Based on Ion Production:
  • Strong Acids: Acids that produce more H⁺ ions in solution are considered strong acids (e.g., hydrochloric acid).
  • Weak Acids: Acids that produce fewer H⁺ ions are considered weak acids (e.g., acetic acid).
  • Strong Bases: Bases that produce more OH⁻ ions are considered strong bases (e.g., sodium hydroxide).
  • Weak Bases: Bases that produce fewer OH⁻ ions are considered weak bases.

Conclusion:

• The strength of an acid or base is determined by its ability to produce ions in solution.
• The pH scale provides a quantitative measure of this strength, indicating the acidic or basic nature of the solution.
• Understanding these concepts is crucial for various applications in chemistry, biology, and environmental science.

Importance of pH in Everyday Life

Understanding the significance of pH in our daily lives reveals its crucial role in various environmental and biological processes.pH Sensitivity in Living Organisms:

1. Human Body:
   • The human body operates optimally within a pH range of 7.0 to 7.8.
   • This pH balance is essential for various bodily functions and maintaining overall health.
2. Plants and Animals:
   • Both plants and animals are sensitive to pH changes.
   • They can only survive and thrive within a narrow pH range.

1. Acid Rain:
   • Rainwater with a pH less than 5.6 is classified as acid rain.
   • Acid rain can lower the pH of river water, adversely affecting aquatic life.
2. Soil pH and Plant Growth:
   • The pH of soil is a critical factor for the healthy growth of plants.
   • Different plants require different pH levels in the soil for optimal growth.

• Acids in Other Planets:
  • For instance, the atmosphere of Venus is composed of thick clouds of sulfuric acid.
  • Such extreme acidic conditions raise questions about the possibility of life on planets with harsh environments.

• Activity 2.12: Involves collecting soil samples from various locations and testing their pH.
• Purpose: To determine the pH range required for the healthy growth of specific plants and understand the relationship between soil pH and plant species in different regions.

Activity 2.12

In Activity 2.12, you are conducting an experiment to determine the pH of the soil in your region. Here are the steps and conclusions:

1. Put about 2 g of soil in a test tube.
2. Add 5 mL of water to the test tube.
3. Shake the contents of the test tube to mix the soil and water.
4. Filter the contents to remove solid particles and collect the filtrate in another test tube.
5. Check the pH of the filtrate using universal indicator paper.
   
   

1. pH in our digestive system: The stomach produces hydrochloric acid to aid in digestion, and antacids like magnesium hydroxide are used to neutralize excess acid when indigestion occurs.
   
   
2. pH change as the cause of tooth decay: Tooth decay begins when the pH in the mouth drops below 5.5, leading to the corrosion of tooth enamel. Regular cleaning of the mouth, including using basic toothpaste, can help prevent tooth decay.
   
   
3. Self-defense by animals and plants through chemical warfare: Animals like honey-bees and plants like nettles use chemical substances with varying pH levels for defense. For example, bee stings inject an acidic substance, and using a mild base like baking soda can provide relief. Nettle stings contain methanoic acid, and dock plants growing nearby can provide a remedy by having properties that neutralize the acid.
   
   
4. Nature provides neutralization options: The example of nettles and dock plants demonstrates how nature often provides natural remedies for chemical irritations and stings. Other remedies might exist for similar situations, depending on the region and the specific irritants involved.

2.4 MORE ABOUT SALTS

2.4.1 Family of Salts

Here are the chemical formulae of the salts mentioned in Activity 2.13:

1. Potassium Sulphate: K₂SO₄
2. Sodium Sulphate: Na₂SO₄
3. Calcium Sulphate: CaSO₄
4. Magnesium Sulphate: MgSO₄
5. Copper Sulphate: CuSO₄
6. Sodium chloride: NaCl
7. Sodium nitrate: NaNO₃
8. Sodium carbonate: Na₂CO₃
9. Ammonium chloride: NH₄Cl

Now, let's identify the acids and bases from which these salts may be obtained:

1. Potassium Sulphate: Obtained from Sulphuric acid (H₂SO₄) and Potassium hydroxide (KOH).
2. Sodium Sulphate: Obtained from Sulphuric acid (H₂SO₄) and Sodium hydroxide (NaOH).
3. Calcium Sulphate: Obtained from Sulphuric acid (H₂SO₄) and Calcium hydroxide (Ca(OH)₂).
4. Magnesium Sulphate: Obtained from Sulphuric acid (H₂SO₄) and Magnesium hydroxide (Mg(OH)₂).
5. Copper Sulphate: Obtained from Sulphuric acid (H₂SO₄) and Copper oxide (CuO).
6. Sodium chloride: Obtained from Hydrochloric acid (HCl) and Sodium hydroxide (NaOH).
7. Sodium nitrate: Obtained from Nitric acid (HNO₃) and Sodium hydroxide (NaOH).
8. Sodium carbonate: Obtained from Carbonic acid (H₂CO₃) and Sodium hydroxide (NaOH).
9. Ammonium chloride: Obtained from Hydrochloric acid (HCl) and Ammonia (NH₃).

1. Sodium Salts:
   • Sodium Sulphate (Na₂SO₄)
   • Sodium chloride (NaCl)
   • Sodium nitrate (NaNO₃)
   • Sodium carbonate (Na₂CO₃)
2. Potassium Salts:
   • Potassium Sulphate (K₂SO₄)
3. Calcium Salts:
   • Calcium Sulphate (CaSO₄)
4. Magnesium Salts:
   • Magnesium Sulphate (MgSO₄)
5. Copper Salts:
   • Copper Sulphate (CuSO₄)
6. Ammonium Salt:
   • Ammonium chloride (NH₄Cl)

2.4 MORE ABOUT SALTS

2.4.2 - pH of Salts

Activity 2.14
- 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), and sodium hydrogencarbonate (NaHCO₃).
- Check their solubility in water (use distilled water only).
- Check the action of these solutions on litmus and find the pH using a pH paper.
- Which of the salts are acidic, basic or neutral? Identify the acid or base used to form the salt.
- Report your observations in Table 2.4.
Salts of a strong acid and a strong base are neutral with pH value of 7. On the other hand, salts of a strong acid and weak base are acidic with pH value less than 7 and those of a strong base and weak acid are basic in nature, with pH value more than 7.

In this table:

• Solubility indicates whether the salt dissolves in water.
• Effect on Litmus tells you whether the solution turns red (acidic), blue (basic), or has no change (neutral) on litmus paper.
• pH Value gives the approximate pH of the solution.
• Nature describes whether the salt is acidic, basic, or neutral based on the pH.
• The last column identifies the acid (if any) and the base used to form the salt.

• Salt Formation: The combination of hydrochloric acid and sodium hydroxide solution results in the formation of sodium chloride, commonly known as table salt. It is a staple in food.
• Neutral Salt: Sodium chloride, as observed in the previous activity, is classified as a neutral salt because it has a pH close to 7.
• Seawater Composition: Seawater contains a mixture of various dissolved salts, including sodium chloride, along with other minerals and elements.
• Sodium Chloride Separation: Sodium chloride is separated from other salts present in seawater through various methods.
• Rock Salt: Solid deposits of salt, known as rock salt, are found in various regions around the world. These salt crystals are often brown due to impurities.
• Formation of Rock Salt: Beds of rock salt were created during ancient times when seas dried up, leaving behind these deposits.
• Mining Process: Rock salt is extracted from the Earth in a manner similar to coal mining.
• Historical Significance: Sodium chloride, also known as common salt, held significant symbolic importance during Mahatma Gandhi's Dandi March, which was a pivotal event in India's struggle for independence.

1. Common Salt as a Raw Material:

• Common salt serves as a crucial raw material for various everyday substances, including:
1. Sodium hydroxide
2. Baking soda
3. Washing soda
4. Bleaching powder

2. Sodium Hydroxide Production:

• Sodium hydroxide (NaOH) is produced by passing electricity through an aqueous solution of sodium chloride (brine) in a process known as the chlor-alkali process.
• The name "chlor-alkali" refers to the products formed:
1. "Chlor" for chlorine
2. "Alkali" for sodium hydroxide

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

• During this process:
1. Chlorine gas (Cl2) is released at the anode.
2. Hydrogen gas (H2) is released at the cathode.
3. Sodium hydroxide solution (NaOH) is formed near the cathode.
• All three products have various useful applications.
  
  

3. Bleaching Powder Production:

• Chlorine gas produced during the electrolysis of aqueous sodium chloride is used for manufacturing bleaching powder.
• Bleaching powder is produced by the action of chlorine on dry slaked lime [Ca(OH)₂].

Chemical Equation: Ca(OH)₂ + Cl₂ → CaOCl₂ + H₂O

• Uses of Bleaching Powder: Bleaching powder has several applications:
1. It's used for bleaching cotton and linen in the textile industry.
2. It's employed for bleaching wood pulp in paper factories.
3. It's used for bleaching washed clothes in laundry.
4. It serves as an oxidizing agent in many chemical industries.
5. It's utilized to purify drinking water, making it free from harmful germs.

4. Baking Soda Usage:

Baking soda, chemically known as sodium hydrogencarbonate (NaHCO3), is commonly used in kitchens for various purposes. It is used to make crispy pakoras, expedite cooking, and add a pleasant taste to food.

Production of Baking Soda:

• Baking soda is produced using sodium chloride (NaCl) as one of its raw materials. The reaction involves the following chemical equation:
• NaCl + H₂O + CO₂ + NH₃ → NH₄Cl + NaHCO₃ (Ammonium chloride + Sodium hydrogencarbonate)

pH of Sodium Hydrogencarbonate:

• Sodium hydrogencarbonate is mildly basic and non-corrosive. In Activity 2.14, its pH can be observed to be alkaline. This property makes it suitable for neutralizing acids.

Heating of Sodium Hydrogencarbonate:

• When sodium hydrogencarbonate is heated during cooking, it undergoes a reaction that produces sodium carbonate (Na2CO3), water (H2O), and carbon dioxide (CO2): 2NaHCO₃ → Na₂CO₃ + H₂O + 2CO₂

Household Uses of Baking Soda:

• Baking Powder: Baking soda is used to make baking powder, which is a mixture of baking soda and a mild edible acid like tartaric acid. When heated or mixed with water, this reaction occurs:
• NaHCO₃ + H⁺ → CO₂ + H₂O⁺ Sodium salt of acid (From any acid)

The release of carbon dioxide gas causes bread or cake to rise, making them soft and spongy.

• Antacid: Sodium hydrogencarbonate is an ingredient in antacids. Its alkaline nature allows it to neutralize excess stomach acid, providing relief from indigestion.
• Fire Extinguishers: Sodium hydrogencarbonate is used in soda-acid fire extinguishers for firefighting purposes.

5. Washing Soda 

The ${\text{10H}}_2\text{O}$ in the compound ${\text{Na}}_2{\text{CO}}_3\cdot 10{\text{H}}_2\text{O}$ (washing soda) signifies that there are 10 water molecules associated with each sodium carbonate ( ${\text{Na}}_2{\text{CO}}_3$) molecule. This is referred to as water of crystallization, which is water that is included in the crystal structure of a compound. These water molecules are not just physically trapped, but they are a part of the crystal structure of the compound, giving it a specific geometric shape.
This does not mean that the ${\text{Na}}_2{\text{CO}}_3$ is wet. Rather, the water molecules are chemically bonded to the sodium carbonate, and the compound is typically in the form of a powder. This water of crystallization is important for the compound's properties and stability. When you heat washing soda, these water molecules can be driven off, leaving anhydrous sodium carbonate, which is without water.As for its uses, washing soda is an important industrial and household chemical:

1. In Industries:
   • Glass Industry: Washing soda is used in the production of glass.
   • Soap and Detergent Industry: It's a major component in soap and detergent manufacturing.
   • Paper Industry: It's used in the process of making paper.
2. Manufacture of Sodium Compounds:
   • Washing soda is used as a starting material for producing other sodium compounds like borax.
3. Cleaning Agent:
   • It is used as a cleaning agent for domestic purposes due to its ability to cut grease, remove stains, and neutralize odors.
4. Water Softening:
   • Washing soda is used to remove the 'permanent hardness' of water, which is caused by the presence of dissolved calcium and magnesium salts. It does so by precipitating these salts out of the water.

2.4.4 Are the Crystals of Salts really Dry?

1. Activity with Copper Sulphate (CuSO₄·5H₂O):
   • Heating Copper Sulphate Crystals: When you heat copper sulphate crystals, which are blue due to water of crystallization, the water molecules are driven off as the heat provides enough energy to break the bonds holding the water in the crystal structure.
   • Color Change: Upon heating, the color of hydrated copper sulphate changes from blue to white, which is the color of anhydrous copper sulphate (CuSO₄ without water).
   • Observation of Water Droplets: The water droplets observed in the boiling tube come from the water of crystallization that has been released from the copper sulphate crystals due to heating.
   • Rehydration: When water is added to the anhydrous white copper sulphate, the blue color is restored, indicating that the water of crystallization has been reabsorbed into the crystal structure, reforming the hydrated compound.
2. Water of Crystallisation:
   • This is the term for water molecules that are integrated into the crystal structure of a salt in a fixed ratio. It is not liquid water, but part of the solid crystal lattice.
   • Copper Sulphate: CuSO₄·5H₂O means there are five water molecules per formula unit of copper sulphate.
   • Gypsum: Similarly, gypsum (CaSO₄·2H₂O) contains two water molecules per formula unit of calcium sulphate.
3. Na₂CO₃ · 10H₂O : With this understanding, we can say that the crystals of Na₂CO₃ · 10H₂O (washing soda) are not "wet" in the usual sense. They contain water of crystallization, which is part of their crystalline structure. So, even though the crystals appear dry, they contain water molecules within their crystal lattice.
4. Gypsum and Its Uses:
   • Gypsum, or calcium sulfate dihydrate (CaSO₄·2H₂O), has two water molecules as water of crystallization.
   • When gypsum is heated, it loses water and becomes calcium sulfate hemihydrate, commonly known as plaster of Paris (CaSO₄·½H₂O).
   • Gypsum is used in the construction industry for making drywall or plasterboard and as a soil conditioner in agriculture. It's also a primary material in the creation of plaster of Paris, which is used for castings and moldings.

CBSE Class 10 Science Chapter 1 - Chemical Reactions and Equations Notes

Class 10 CBSE Important Questions and Answers Chapter 1 - Chemical Reactions and Equations