Unit 2: Stoichiometry

Unit 2: Stoichiometry (8 Teaching Hours)

1. Dalton’s Atomic Theory and its Postulates

Dalton's atomic theory, proposed in 1808, laid the foundation of modern chemistry. The key postulates are:

• Atoms are indivisible particles: They cannot be created, divided, or destroyed.
• Atoms of the same element are identical: They have the same mass and properties.
• Atoms of different elements are different: They differ in mass and properties.
• Compounds are formed by a combination of atoms: Atoms of different elements combine in simple whole-number ratios.
• Chemical reactions involve rearrangement of atoms: Atoms are neither created nor destroyed during chemical reactions.

2. Laws of Stoichiometry

Stoichiometry is based on fundamental laws of chemical combinations:

• Law of Conservation of Mass: In a chemical reaction, the total mass of reactants is equal to the total mass of products.
• Law of Definite Proportions: A chemical compound always contains the same elements in the same proportion by mass.
• Law of Multiple Proportions: If two elements combine to form more than one compound, the mass of one element combines with a fixed mass of the other in simple whole-number ratios.

3. Avogadro’s Law and Some Deductions

• Avogadro's Law: Equal volumes of gases, at the same temperature and pressure, contain an equal number of molecules.

  Deductions from Avogadro's Law:

  • Molecular Mass and Vapour Density: The molecular mass of a gas is twice its vapour density.
  • Molecular Mass and Volume of Gas: One mole of any gas occupies 22.4 liters at standard temperature and pressure (STP).
  • Molecular Mass and Number of Particles: The number of particles (atoms, molecules) in one mole is 6.022 × 10²³ (Avogadro's number).

4. Mole and Its Relationship with Mass, Volume, and Number of Particles

• Mole: A unit representing 6.022 × 10²³ particles of any substance.
• Mass: 1 mole of a substance has a mass equal to its molecular or atomic mass in grams.
• Volume: 1 mole of a gas at STP occupies 22.4 liters.
• Number of Particles: 1 mole contains 6.022 × 10²³ atoms, molecules, or ions.

5. Calculations Based on Mole Concept

• Mass of a substance: $\text{Mass} = \text{Number of moles} \times \text{Molar mass}$
• Number of moles: $\text{Moles} = \frac{\text{Given mass}}{\text{Molar mass}}$
• Number of particles: $\text{Particles} = \text{Moles} \times 6.022 \times 10^{23}$
• Volume of gas: $\text{Volume} = \text{Number of moles} \times 22.4 \, \text{liters}$ (at STP)

6. Limiting Reactant and Excess Reactant

• Limiting Reactant: The reactant that is completely consumed in a chemical reaction, limiting the amount of product formed.
• Excess Reactant: The reactant that remains unreacted after the limiting reactant is used up.

7. Theoretical Yield, Experimental Yield, and Percentage Yield

• Theoretical Yield: The maximum amount of product that can be formed from the limiting reactant, calculated based on stoichiometric ratios.
• Experimental Yield: The actual amount of product obtained from an experiment.
• Percentage Yield: $\text{Percentage Yield} = \left( \frac{\text{Experimental Yield}}{\text{Theoretical Yield}} \right) \times 100$

8. Calculation of Empirical and Molecular Formula from Percent Composition

• Empirical Formula: The simplest whole-number ratio of atoms in a compound.
  • Steps:
    1. Convert the percentage composition of each element to mass (in grams).
    2. Convert the masses to moles by dividing by the atomic mass.
    3. Find the simplest mole ratio and write the empirical formula.
• Molecular Formula: The actual number of atoms of each element in a molecule.
  • It can be found by multiplying the empirical formula by the ratio of the molecular mass to the empirical formula mass.

9. Solving Related Numerical Problems

• Calculations involving the mole concept, limiting and excess reactants, percentage yield, and empirical/molecular formulas are essential for mastering stoichiometry.