Atoms, Molecules and Stoichiometry

This chapter builds foundational understanding in chemistry by exploring atomic and molecular masses, the mole concept, chemical formulas, and stoichiometric calculations involving masses, volumes, and solutions.

📘 2.1 Relative Masses of Atoms and Molecules

🔹 Unified Atomic Mass Unit (u)

• Definition: 1 unified atomic mass unit (1 u) is defined as 1/12 of the mass of a carbon-12 atom.
• It’s a standard for expressing atomic and molecular masses.

🔹 Definitions

• Relative Atomic Mass (Ar):
  Weighted average mass of an atom of an element compared to 1/12 of the mass of a carbon-12 atom.
• Relative Isotopic Mass:
  Mass of a specific isotope relative to 1/12 of the mass of a carbon-12 atom.
• Relative Molecular Mass (Mr):
  Sum of the relative atomic masses of all atoms in a molecule.
• Relative Formula Mass:
  Similar to Mr, but for ionic compounds, calculated using formula units.

📘 2.2 The Mole and the Avogadro Constant

🔹 Definition of the Mole

• A mole is the amount of a substance that contains Avogadro’s number (6.022 × 10²³) of particles (atoms, molecules, ions, etc.).
• The Avogadro constant is the number of specified particles in one mole of a substance.

📘 2.3 Formulas

🔹 Writing Formulas

1. Ionic Compounds:
   • Use oxidation numbers or positions in the periodic table to predict charges.
   • Write correct formulas by balancing the total charges of cations and anions.
   • Examples of important ions:
     • Nitrate (NO₃⁻), Carbonate (CO₃²⁻), Sulfate (SO₄²⁻), Hydroxide (OH⁻),
       Ammonium (NH₄⁺), Zinc (Zn²⁺), Silver (Ag⁺), Bicarbonate (HCO₃⁻),
       Phosphate (PO₄³⁻)
2. Balanced Equations:
   • Include both molecular and ionic equations.
   • Exclude spectator ions from ionic equations.
   • State symbols: (s), (l), (g), (aq)

🔹 Empirical and Molecular Formulas

• Empirical Formula: Simplest whole number ratio of atoms of each element in a compound.
• Molecular Formula: Actual number of atoms of each element in a molecule.
• Calculated using experimental data (mass or % composition).

🔹 Hydrated Compounds

• Anhydrous: Substance without water.
• Hydrated: Compound with water molecules in its crystalline structure.
• Water of Crystallisation: Fixed number of water molecules associated with each formula unit (e.g., CuSO₄·5H₂O)

🔹 Formula Calculations

• Use experimental data to determine empirical and molecular formulas.
• Apply mass ratios and molar masses.

📘 2.4 Reacting Masses and Volumes (of Solutions and Gases)

🔹 Molar Calculations

1. Reacting Masses:
   • Use balanced equations and molar ratios.
   • Include percentage yield:
     % Yield = (Actual Yield / Theoretical Yield) × 100
2. Gases:
   • At RTP (Room Temp. and Pressure), 1 mole of any gas occupies 24 dm³.
   • Apply volume ratios from balanced equations.
3. Solutions:
   • Use formula:
     moles = concentration (mol/dm³) × volume (dm³)
4. Limiting and Excess Reagents:
   • Identify the limiting reagent (reactant completely used up).
   • The other is in excess and doesn’t limit the reaction yield.
5. Stoichiometry:
   • Determine mole ratios from chemical equations.
   • Deduce relationships from given masses, volumes, or concentrations.

✅ Tips for Success:

• Always balance equations before performing calculations.
• Use the correct number of significant figures.
• Clearly show units in calculations.
• For ionic equations, remove spectator ions.

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