Atoms and Molecules
Introduction
Ancient Indian and Greek philosophers proposed that matter is made up of tiny, indivisible particles. This chapter explores the fundamental units of matter - atoms and molecules.
Laws of Chemical Combination
Several laws govern how elements combine to form compounds:
1. Law of Conservation of Mass
"Mass can neither be created nor destroyed in a chemical reaction."
In other words: Total mass of reactants = Total mass of products
Example: When 12g of carbon reacts with 32g of oxygen to form carbon dioxide, the mass of carbon dioxide formed is 44g (12g + 32g = 44g).
2. Law of Constant Proportions (Law of Definite Proportions)
"In a chemical compound, the elements are always present in definite proportions by mass."
Example: Pure water (H₂O) always contains hydrogen and oxygen in the ratio of 1:8 by mass, regardless of its source (river, sea, rain, etc.).
Dalton's Atomic Theory (1808)
John Dalton proposed the atomic theory to explain the laws of chemical combination:
- All matter is made up of tiny, indivisible particles called atoms
- Atoms of the same element are identical in mass and properties
- Atoms of different elements have different masses and properties
- Atoms combine in the ratio of small whole numbers to form compounds
- Atoms can neither be created nor destroyed in a chemical reaction
Limitations of Dalton's Atomic Theory
- Atoms are divisible into electrons, protons, and neutrons
- Atoms of the same element can have different masses (isotopes)
- Atoms can combine in ratios which are not whole numbers (in complex organic compounds)
What is an Atom?
Atoms are extremely small - their radius is measured in nanometers (nm):
1 nanometer = 10⁻⁹ meters
Atomic Mass
Examples of Atomic Masses:
- Hydrogen (H) = 1 u
- Carbon (C) = 12 u
- Nitrogen (N) = 14 u
- Oxygen (O) = 16 u
- Sodium (Na) = 23 u
- Chlorine (Cl) = 35.5 u
What is a Molecule?
Types of Molecules
1. Molecules of Elements
Molecules formed by atoms of the same element.
a) Monoatomic: Contains only one atom
- Noble gases: Helium (He), Neon (Ne), Argon (Ar)
b) Diatomic: Contains two atoms
- Hydrogen (H₂), Oxygen (O₂), Nitrogen (N₂), Chlorine (Cl₂)
c) Triatomic: Contains three atoms
- Ozone (O₃)
d) Polyatomic: Contains more than three atoms
- Phosphorus (P₄), Sulphur (S₈)
2. Molecules of Compounds
Molecules formed by atoms of different elements.
- Water (H₂O) - 2 hydrogen + 1 oxygen
- Carbon dioxide (CO₂) - 1 carbon + 2 oxygen
- Ammonia (NH₃) - 1 nitrogen + 3 hydrogen
- Methane (CH₄) - 1 carbon + 4 hydrogen
Chemical Formula
Rules for Writing Chemical Formulas:
- Write symbols of elements
- Write valency below each symbol
- Interchange the valencies
- Simplify if possible
Molecular Mass
Example: Water (H₂O)
Molecular mass of H₂O = 2 × (Atomic mass of H) + 1 × (Atomic mass of O)
= 2 × 1 + 1 × 16 = 2 + 16 = 18 u
Formula Unit Mass
Example: Sodium chloride (NaCl)
Formula unit mass = Atomic mass of Na + Atomic mass of Cl
= 23 + 35.5 = 58.5 u
Ions
Types of Ions:
1. Cation (Positive Ion): Formed by loss of electrons
- Na → Na⁺ + e⁻ (Sodium ion)
- Mg → Mg²⁺ + 2e⁻ (Magnesium ion)
- Al → Al³⁺ + 3e⁻ (Aluminium ion)
2. Anion (Negative Ion): Formed by gain of electrons
- Cl + e⁻ → Cl⁻ (Chloride ion)
- O + 2e⁻ → O²⁻ (Oxide ion)
- N + 3e⁻ → N³⁻ (Nitride ion)
Polyatomic Ions
Common Polyatomic Ions:
| Ion Name | Formula | Charge |
|---|---|---|
| Ammonium | NH₄ | +1 |
| Hydroxide | OH | -1 |
| Carbonate | CO₃ | -2 |
| Sulphate | SO₄ | -2 |
| Nitrate | NO₃ | -1 |
| Phosphate | PO₄ | -3 |
The Mole Concept
Avogadro's Number (Nₐ): 6.022 × 10²³ is called Avogadro's constant or Avogadro's number, named after Italian scientist Amedeo Avogadro.
Molar Mass
Examples:
- Molar mass of water (H₂O) = 18 g/mol
- Molar mass of carbon dioxide (CO₂) = 44 g/mol
- Molar mass of sodium chloride (NaCl) = 58.5 g/mol
Important Formulas
Relationship between Mass, Moles, and Particles
Key Relationships:
- 1 mole of any substance = Its molar mass in grams
- 1 mole of any substance = 6.022 × 10²³ particles
- Mass in grams = Number of moles × Molar mass
Examples of Calculations
Example 1: Calculate the number of moles in 32g of oxygen (O₂).
Molar mass of O₂ = 2 × 16 = 32 g/mol
Number of moles = 32g / 32 g/mol = 1 mole
Example 2: How many molecules are present in 18g of water?
Molar mass of H₂O = 18 g/mol
Number of moles = 18g / 18 g/mol = 1 mole
Number of molecules = 1 × 6.022 × 10²³ = 6.022 × 10²³ molecules
Important Points to Remember
- Atoms are the building blocks of matter
- Molecules are formed when atoms combine
- Chemical formulas represent the composition of molecules
- The mole is a convenient unit for counting particles
- 1 mole = 6.022 × 10²³ particles = Molar mass in grams
- Atomic mass is measured in atomic mass units (u)
- Molecular mass is the sum of atomic masses
Multiple Choice Questions (MCQ)
Subjective Questions
Practice these questions to strengthen your understanding. Write your answers in the space provided.