class 11 chemistry language of chemistry notes || class 11 chemistry unit 1

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🧪 Chapter: Language of Chemistry – Class 11

🔹 1. Symbols of Elements

• Definition: A symbol is the shorthand representation of an element.

• Modern Symbols: Based on the element's English/Latin name.

  • E.g., Hydrogen → H, Sodium → Na (from Natrium), Iron → Fe (Ferrum).

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🔹 2. Chemical Formula

• Definition: It represents the composition of a compound using element symbols and their ratios.

• Types:

  • Empirical Formula: Shows simplest ratio of atoms (e.g., CH₂O for glucose).

  • Molecular Formula: Actual number of atoms (e.g., C₆H₁₂O₆).

  • Structural Formula: Shows arrangement of atoms.

Examples:

• Water: H₂O

• Carbon Dioxide: CO₂

• Methane: CH₄

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🔹 3. Valency

• Definition: Combining capacity of an element.

• Based on: Number of electrons lost, gained, or shared.

• Common Valencies:

  • H = 1, O = 2, N = 3, C = 4

Criss-Cross Method for formula:

• Al³⁺ and O²⁻ → Al₂O₃

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🔹 4. Radicals

• Definition: A group of atoms that behave like a single unit and have a charge.

• Types:

  • Cation (+ve): NH₄⁺, Na⁺, Ca²⁺

  • Anion (-ve): NO₃⁻, SO₄²⁻, Cl⁻

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🔹 5. Chemical Equations

• Definition: Shorthand representation of a chemical reaction.

• Types:

  • Word Equation: Hydrogen + Oxygen → Water

  • Symbolic Equation: H₂ + O₂ → H₂O (needs balancing)

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🔹 6. Balancing Chemical Equations

• Law of Conservation of Mass: Total mass of reactants = Total mass of products.

• Steps:

  1. Write unbalanced equation.

  2. Count atoms of each element.

  3. Balance one by one using coefficients.

Example: Unbalanced: H₂ + O₂ → H₂O
Balanced: 2H₂ + O₂ → 2H₂O

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🔹 7. Types of Chemical Reactions

1. Combination: A + B → AB
   Example: 2H₂ + O₂ → 2H₂O

2. Decomposition: AB → A + B
   Example: 2KClO₃ → 2KCl + 3O₂

3. Displacement: A + BC → AC + B
   Example: Zn + CuSO₄ → ZnSO₄ + Cu

4. Double Displacement: AB + CD → AD + CB
   Example: NaCl + AgNO₃ → AgCl + NaNO₃

5. Neutralization: Acid + Base → Salt + Water
   Example: HCl + NaOH → NaCl + H₂O

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🔹 8. Atomic Mass & Molecular Mass

• Atomic Mass Unit (amu): 1/12th of mass of one atom of carbon-12.

• Atomic Mass: Relative mass of one atom.

• Molecular Mass = Sum of atomic masses in a molecule.

Example (H₂O):

• H = 1, O = 16 → Molecular Mass = 2(1) + 16 = 18 amu

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🔹 9. Mole Concept

• 1 mole = 6.022 × 10²³ particles (Avogadro's number)

• Molar Mass = Mass of 1 mole of substance (g/mol)

• Formula:

  $$\text{Number of moles} = \frac{\text{Given mass (g)}}{\text{Molar mass (g/mol)}}$$

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🔹 10. Concentration Terms

• Molarity (M): Moles of solute / Litre of solution

  $$M = \frac{\text{n}}{\text{V (L)}}$$

• Molality (m): Moles of solute / kg of solvent

• Normality (N): Gram equivalents / Litre of solution

• Percentage Composition:

  \text{Mass % of element} = \frac{\text{Mass of element}}{\text{Molar mass}} \times 100

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🔹 11. Empirical and Molecular Formula Calculations

• Steps:

1. Find % of each element.

2. Divide by atomic masses.

3. Get simplest ratio → empirical formula.

4. Molecular Formula = (Empirical Formula) × n, where

   $$n = \frac{\text{Molar mass}}{\text{Empirical formula mass}}$$

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