⚛️ Class 10 Chemistry — Chapter 6: Salts

1. Salts are ionic compounds formed by the electrostatic attraction between oppositely charged ions.

2. The base/metal supplies the cation (positively charged ion).

3. The acid supplies the anion (negatively charged ion/acid radical).

4. The combination forms a neutral compound.

1. The ions are held together by a strong electrostatic force of attraction (ionic bond).

2. The positive charge supplied by the cation equals the negative charge supplied by the anion.

3. Since the net charges balance, the salt is electrically neutral.

1. The cation (e.g., Na⁺) is derived from the base/metal and replaces the hydrogen ion of the acid.

2. The anion (e.g., Cl⁻) is the acid radical derived from the acid.

3. The salt forms by combining the metal part of the base and the non-metal part of the acid.

1. Salts form a rigid crystal lattice structure at STP.

2. Ions are bound by very strong electrostatic forces (ionic bonds).

3. A large amount of energy is required to break these bonds, causing high melting points.

1. Ionic compounds exist as crystalline solids under normal conditions.

2. Their ions arrange in a regular 3D structure called a crystal lattice.

3. Each ion is surrounded by oppositely charged ions, maximizing attraction and providing rigidity.

1. High melting point: due to strong electrostatic attractions.

2. Brittle nature: when layers shift, like charges align and repel, causing the crystal to fracture.

3. Hardness: strong forces lock ions into fixed positions.

1. Solid salts: ions are fixed in the lattice and immobile, so they do not conduct electricity.

2. Molten salts: heating frees ions, making them mobile and able to conduct electricity.

3. Mobile ions carry charge, so molten salts are good conductors.

1. When a salt dissolves, it dissociates into ions.

2. Cations move to the cathode and anions move to the anode.

3. This coordinated movement of charged ions constitutes the electric current.

1. In solid NaCl ions are fixed in the lattice.

2. In aqueous solution, polar water molecules pull ions out of the lattice.

3. Ions become mobile and can carry current.

1. All salts of Sodium (Na), Potassium (K), and Ammonium (NH₄) are soluble.

2. All salts containing the Nitrate ion (NO₃) are soluble.

3. Most Chlorides are soluble except for Lead (Pb²⁺) and Silver (Ag⁺).

1. Most carbonates are insoluble in water.

2. Exceptions: carbonates of Na⁺ and K⁺ are soluble (e.g., Na₂CO₃).

3. CaCO₃ (limestone) is an example of an insoluble carbonate.

1. Precipitation occurs when two soluble salts react and an insoluble product forms.

2. Solubility rules predict which product is insoluble and will separate as precipitate.

3. The insoluble product appears as a solid that settles out of the solution.

1. Filtration: remove unreacted excess solid from the solution.

2. Evaporation: heat the solution to concentrate it (saturated solution).

3. Crystallization: cool slowly to form pure crystals.

4. Drying: wash and dry the crystals to obtain the final product.

1. Adding excess ensures the limiting reactant is fully consumed.

2. It prevents contamination of the final product with unreacted acid.

3. Filtration removes the unreacted excess solid, giving a pure aqueous salt solution.

1. Using metal/carbonate: add excess solid and then filter — straightforward purification.

2. Using soluble alkali: perform titration to get a neutral solution — requires precise measurement.

3. The excess method is easier because the excess reagent is insoluble and can be filtered out.