Atomic Foundations
of Matter
🧪 handwritten-style notes for quick revision 🧪
📑 Table of Contents
⚖️ Law of Conservation of Mass
Given by Antoine Lavoisier (1789) — the “Father of Modern Chemistry”.
He stated: “…in every operation an equal quantity of matter exists both before and after the operation.”
🧫 Quick Activity Recap
- Salt + Water (physical change): mass of solution = mass of salt + mass of water. No change in mass.
- Vinegar + Baking soda in open flask: gas (CO₂) escapes → final reading ≠ initial reading (looks like mass lost).
- Same reaction in a closed/tied balloon setup: gas is trapped → final reading = initial reading ✅
➡️ So mass is always conserved — the mismatch in the open flask was only because gas escaped, not because mass was destroyed.
🧭 Law of Constant Proportions
Also called Law of Definite Proportions or Proust’s Law — given by Joseph Louis Proust.
Example: Water from a river, borewell or ocean — when purified — always has Hydrogen : Oxygen = 1 : 8 by mass.
So 9 g of pure water always gives 1 g hydrogen + 8 g oxygen.
⚛️ Dalton’s Atomic Theory
Both laws above are explained by John Dalton’s postulates (1808):
- All matter is made of tiny particles called atoms.
- Atoms are indivisible — cannot be created or destroyed in a reaction.
- Atoms of the same element are identical in mass & properties.
- Atoms of different elements differ in mass & properties.
- Atoms combine in simple whole-number ratios to form compounds.
- The relative number & kind of atoms is constant in a given compound.
English scientist, taught in Manchester. Presented his atomic theory in 1808 — a turning point in the study of matter.
🔗 How Atoms Combine?
Atoms need 8 electrons in the valence shell (2 for K-shell) to be stable — i.e. a full octet.
To complete the octet, atoms combine by:
- Sharing of electrons → forms a Covalent Bond
- Transfer of electrons → forms an Ionic Bond
Noble gases like Helium already have a stable outer shell → they exist as single atoms.
🤝 Covalent Bond (Sharing)
Formed when atoms share electron pairs. The shared pair attracts both nuclei and holds the molecule together.
| Molecule | Electrons shared | Bond type | Formula |
|---|---|---|---|
| H₂ | 1 each | Single | H—H |
| Cl₂ | 1 each | Single | Cl—Cl |
| O₂ | 2 each | Double | O=O |
| HCl | 1 each | Single | H—Cl |
| H₂O | 1 e⁻ from each H to O | 2 single bonds | H—O—H |
✏️ Naming Covalent (Molecular) Compounds
- First element keeps its name; second element ends in -ide.
- Prefixes show atom count: mono-1, di-2, tri-3, tetra-4, penta-5, hexa-6.
- Mono- is skipped for the first element, used for the second.
- Drop the vowel before another vowel: monoxide, pentoxide (not monooxide).
H₂S → hydrogen sulfide (no prefix before H) | H₂O = water, NH₃ = ammonia (common names)
⚡ Ionic Bond (Transfer)
Metals (usually <4 valence e⁻) lose electrons → form cations (+). Non-metals (usually >4 valence e⁻) gain electrons → form anions (–).
Example — NaCl: Na loses 1 e⁻ → Na⁺; Cl gains 1 e⁻ → Cl⁻. Na⁺ + Cl⁻ → NaCl.
Ionic compounds form 3-D crystal lattices (not single molecules) — each Na⁺ is surrounded by 6 Cl⁻ and vice-versa.
🧾 Naming ionic compounds
Cation name first, then anion name (ending in -ide). Metal + non-metal → ionic compound.
| Common Cations | Common Anions |
|---|---|
| Na⁺, K⁺, Ag⁺, Li⁺ (valency 1) | Cl⁻, F⁻, Br⁻, I⁻, OH⁻, NO₃⁻ (valency 1) |
| Ca²⁺, Mg²⁺, Ba²⁺, Zn²⁺, Cu²⁺, Fe²⁺ (valency 2) | O²⁻, S²⁻, CO₃²⁻, SO₄²⁻ (valency 2) |
| Al³⁺, Fe³⁺ (valency 3) | PO₄³⁻ (valency 3) |
Polyatomic ions (e.g. NH₄⁺, OH⁻, NO₃⁻, CO₃²⁻, SO₄²⁻) are groups of atoms carrying a charge together — names usually don’t end in -ide.
✍️ Writing Chemical Formulae
Covalent compounds — criss-cross valencies
C (valency 4) + Cl (valency 1) → criss-cross → CCl₄
Ionic compounds — criss-cross charges
- Write cation symbol first, then anion symbol.
- Write charges below the symbols (not as superscripts).
- Criss-cross the numbers → these become subscripts.
- Simplify subscripts to the smallest whole-number ratio.
- Use brackets ( ) when 2+ polyatomic ions of the same kind are needed, e.g. Mg(OH)₂.
Ca²⁺ + CO₃²⁻ → CaCO₃ | Al³⁺ + OH⁻ → Al(OH)₃ | Al³⁺ + SO₄²⁻ → Al₂(SO₄)₃
🔬 Properties: Ionic vs Covalent
| Property | Ionic Compounds | Covalent Compounds |
|---|---|---|
| Example | NaCl, CuSO₄ | Camphor, Naphthalene, Sugar |
| Solubility | Soluble in water | Insoluble in water, soluble in kerosene/petrol |
| Conducts electricity (solid) | ❌ No (ions fixed in lattice) | ❌ No |
| Conducts electricity (in water) | ✅ Yes (ions free to move) | ❌ Usually no (e.g. sugar solution) |
| Melting/Boiling point | High | Low |
⚖️ Molecular Mass & Formula Unit Mass
Molecular Mass (for covalent compounds)
Sum of atomic masses of all atoms in one molecule.
CO₂ = (12×1) + (16×2) = 44 u
Formula Unit Mass (for ionic compounds)
Ionic compounds don’t form molecules (they form crystal lattices), so we use “formula unit” = simplest whole-number ratio of ions.
Ca(NO₃)₂ = (40×1) + {(14×1)+(16×3)}×2 = 164 u
🌟 At a Glance — Full Chapter Revision
- Mass is neither created nor destroyed in a chemical reaction → Law of Conservation of Mass.
- A compound always has the same elements in a fixed ratio by mass → Law of Definite/Constant Proportions.
- A molecule = neutral, independent, 2+ atom particle showing all properties of the substance.
- Atoms combine to become stable, held by a chemical bond.
- Covalent bond = sharing of electrons.
- Ionic bond = transfer of electrons → forms cations & anions.
- Covalent formula = number of atoms of each element.
- Ionic formula = simplest whole-number ratio of ions.
- Molecular mass = sum of atomic masses in a molecule.
- Formula unit mass = sum of atomic masses in one formula unit of an ionic compound.