Atoms and Molecules: The Building Blocks of Matter

📅 Date: 11 December 2025 (THURSDAY)

Have you ever wondered what happens if you keep cutting a piece of matter into smaller and smaller pieces? Today, we dive into the invisible world of Atoms and Molecules. This chapter bridges ancient philosophy with modern chemistry.

1. The History: Philosophy meets Science

Long before microscopes, philosophers predicted the existence of atoms.

• Maharishi Kanad (500 BC): Named the smallest, indivisible particle "Parmanu".
• Democritus (Greek): Called them "Atoms" (meaning indivisible).

2. Laws of Chemical Combination

By the 18th century, Lavoisier and Proust established two rules that govern how matter behaves.

A. Law of Conservation of Mass

Definition: Mass can neither be created nor destroyed in a chemical reaction.

B. Law of Constant Proportions

Definition: In a chemical substance, elements are always present in definite proportions by mass.

• Example (Water): Whether you take water from a tap, a river, or a lab, the ratio of Hydrogen to Oxygen is always 1:8 by mass.

3. Dalton's Atomic Theory

John Dalton provided the theory that explained why the above laws work. His key postulates:

1. All matter is made of tiny particles called atoms.
2. Atoms are indivisible (cannot be created or destroyed). (Explains Conservation of Mass).
3. Atoms of a given element are identical in mass and properties.
4. Atoms combine in the ratio of small whole numbers to form compounds. (Explains Constant Proportions).

4. What is an Atom?

Atoms are the building blocks of matter. They are incredibly small (measured in nanometers).

Symbols of Elements

IUPAC approves names and symbols. The rule is simple: The first letter is Uppercase, the second is Lowercase.

• Aluminium: Al (Not AL)
• Cobalt: Co (Not CO - that would be Carbon Monoxide!)
• Sodium: Na (From Latin 'Natrium')
• Potassium: K (From Latin 'Kalium')

Atomic Mass (The Watermelon Analogy)

Since atoms are too light to weigh directly, we use Relative Atomic Mass.
Standard: Carbon-12 isotope.
Definition: One atomic mass unit (u) is a mass unit equal to exactly 1/12th the mass of one atom of carbon-12.

5. Molecules and Ions

• Molecule: A group of two or more atoms chemically bonded. (e.g., $O_2$, $H_2O$).
• Atomicity: The number of atoms in a molecule.
  • Argon (Ar) = Monoatomic
  • Oxygen ($O_2$) = Diatomic
  • Ozone ($O_3$) = Triatomic
  • Phosphorus ($P_4$) = Tetra-atomic
• Ion: A charged particle.
  • Cation: Positively charged (e.g., $Na^+$).
  • Anion: Negatively charged (e.g., $Cl^-$).

6. Writing Chemical Formulae (The Criss-Cross Method)

To write a formula, you need the Valency (combining capacity) of the elements.

Example: Magnesium Chloride

1. Symbols: Mg | Cl
2. Valencies: +2 | -1
3. Criss-Cross: Mg gets 1, Cl gets 2.
4. Formula: $MgCl_2$

7. The Mole Concept (Don't Panic!)

This is the most crucial part for numericals.

What is a Mole?

Just as 1 Dozen = 12 items,
1 Mole = $6.022 \times 10^{23}$ particles.
(This number is called the Avogadro Constant).

The Magic Link: Mass and Moles

The mass of 1 mole of a substance is its atomic/molecular mass in grams.

Example: Water ($H_2O$)

• Molecular Mass = $(2 \times 1) + 16 = 18 u$.
• Molar Mass = 18 grams.
• This means: 18g of water contains $6.022 \times 10^{23}$ molecules of water.

Formulas to Remember:

• $Number \ of \ Moles (n) = \frac{Given \ Mass}{Molar \ Mass}$
• $Number \ of \ Moles (n) = \frac{Given \ Number \ of \ Particles}{Avogadro \ Number}$

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Study Tip: Practice writing formulas for Aluminum Sulphate ($Al_2(SO_4)_3$) and calculating the molar mass of Calcium Carbonate ($CaCO_3$). These are exam favorites!