Chapter 9: Atomic Foundations of Matter
Discover the laws that govern matter, how atoms combine through chemical bonds, write chemical formulae, and explore the fascinating properties of ionic and covalent compounds.
📏 Constant Proportions
⚛️ Dalton’s Theory
🔗 Covalent Bond
⚡ Ionic Bond
🧮 Chemical Formulae
⚗️ Molecular Mass
Chapter Introduction: Matter & Its Properties
You already know that matter is made of atoms. But have you wondered — when hydrogen gas and oxygen gas combine to form water, why is water completely different from both gases?
Colourless gas. Highly combustible — catches fire easily. Lighter than air.
Colourless gas. Supports combustion — helps other things burn. Essential for breathing.
When H₂ and O₂ combine, they form water (H₂O) — a liquid that neither burns nor helps burning. It actually extinguishes fire! Yet the mass of water formed = sum of masses of H₂ and O₂ used. This leads us to the Law of Conservation of Mass.
Does mass change during physical changes? What about chemical changes? How do atoms combine? What types of bonds form? How do we write chemical formulae?
Law of Conservation of Mass (द्रव्यमान संरक्षण का नियम)
🔬 Experiments That Proved the Law
When common salt is dissolved in water, the mass of the solution = mass of water + mass of salt. No mass is lost or gained during a physical change.
- In Set-up 1 (open system): baking soda is added to vinegar — CO₂ gas escapes, so final mass < initial mass. Seems like mass is lost!
- In Set-up 2 (closed system with balloon): the CO₂ gas inflates the balloon and is captured. Now final mass = initial mass ✅
If products are not allowed to escape (closed system), total mass before reaction = total mass after reaction. The apparent “loss” of mass in open systems is because gas products escape into the air.
Law of Conservation of Mass: Mass of Reactants = Mass of Products
Matter can neither be created nor destroyed in a chemical reaction.
Lavoisier proposed this law. He conducted careful experiments where he weighed substances before and after reactions in sealed containers. He concluded: “In every operation, an equal quantity of matter exists both before and after the operation.”
📐 Solved Examples
Example 9.1 — Verify Conservation of Mass
Reactants: CaCO₃ = 4.0 g + HCl = 2.92 g
Total mass of reactants = 4.0 + 2.92 = 6.92 g
Products: CO₂ = 1.76 g + H₂O = 0.72 g + CaCl₂ = 4.44 g
Total mass of products = 1.76 + 0.72 + 4.44 = 6.92 g
Since mass of reactants = mass of products → Law of Conservation of Mass is obeyed ✅
Example 9.2 — Carbon + Oxygen → Carbon Dioxide
Given: 12 g C + 32 g O₂ → 44 g CO₂
Find: How much CO₂ is produced by 2.4 g of carbon?
1 g of C gives = 44/12 g of CO₂
2.4 g of C gives = (44/12) × 2.4 = 8.8 g CO₂
If a student burns ethanol in an open beaker and sees no residue, it does NOT mean mass is violated. The carbon and hydrogen atoms from ethanol form CO₂ and H₂O gases that escape into the air. Total mass is still conserved — you need a closed system to verify!
Law of Constant Proportions (निश्चित अनुपात का नियम)
After Lavoisier, French chemist Joseph Proust studied the composition of compounds and found something remarkable: a compound ALWAYS has the same elements in the SAME ratio by mass, no matter where it comes from!
Whether water comes from a river, a borewell, the ocean (after purification), or a lab — it always contains hydrogen and oxygen in the mass ratio of 1:8. If you take 9 g of pure water, you always get 1 g H and 8 g O upon decomposition.
In a given compound, the elements are always present in a fixed ratio by mass — regardless of source or method of preparation.
Ancient civilisations, including India, used red pigment from rocks called hingula (cinnabar in Latin). Heating cinnabar always gave mercury and sulfur in the same ratio — 86.22% and 13.78% by mass. Grinding mercury and sulfur in this exact ratio also reformed cinnabar! This is a perfect example of the Law of Constant Proportions.
📐 Solved Examples
Example 9.3 — Sodium Chloride (NaCl)
NaCl contains Na and Cl in mass ratio 23 : 35.5
Given: 46 g of sodium reacts completely. How much chlorine is needed?
Mass of chlorine = (35.5 ÷ 23) × 46
= 71 g of Chlorine needed
A compound always has fixed element ratios. NaCl always has Na and Cl in 23:35.5 ratio, no matter how it was made.
Air is a mixture — its composition varies (more or less O₂ in different places). Mixtures have no fixed ratio.
Proust was a French chemist famous for careful experimental work. He studied copper carbonate and showed it always contains copper, carbon, and oxygen in the same proportion by mass, regardless of how it was prepared or where it was found.
Dalton’s Atomic Theory (1808)
The two laws of Conservation of Mass and Constant Proportions formed the foundation of John Dalton’s Atomic Theory. Dalton used these laws to explain atomic behavior scientifically.
A postulate is a fundamental assumption accepted as truth without formal proof, from which further ideas are developed.
📋 Dalton’s Postulates
- All matter is made up of very tiny particles called atoms (परमाणु), which participate in chemical reactions.
- Atoms are indivisible particles — they cannot be created or destroyed in a chemical reaction.
- Atoms of a given element are identical in mass and chemical properties.
- Atoms of different elements have different masses and chemical properties.
- Atoms combine in the ratio of simple whole numbers to form compounds.
- The relative number and kinds of atoms are constant in a given compound.
In a chemical reaction, atoms are not destroyed or created. They simply rearrange. So total number of atoms before = after → total mass is conserved.
Atoms combine in fixed whole-number ratios. Since each atom has a definite mass, the same compound always has the same mass ratio of elements.
Later discoveries showed atoms ARE divisible (electrons, protons, neutrons). Atoms of the same element can have different masses (isotopes). But his theory was still revolutionary for its time and forms the foundation of chemistry!
In 1793, Dalton moved to Manchester to teach mathematics, physics, and chemistry. He spent most of his life teaching and researching there. In 1808, he presented his atomic theory — a turning point in the study of matter.
How Atoms Combine — Molecules & Chemical Bonds
A molecule is an electrically neutral entity consisting of more than one atom that is capable of independent existence and shows all the properties of that substance.
Examples: H₂ (two hydrogen atoms), Cl₂ (two chlorine atoms), H₂O (two hydrogen + one oxygen atom), HCl (one hydrogen + one chlorine atom).
Some elements like Helium (He) exist only as single atoms (monoatomic) because their outermost shell is already full (2 electrons). They don’t need to bond with any other atom!
🔗 Why Do Atoms Combine?
Atoms combine because they want to achieve a stable electronic configuration — 8 electrons in the outermost shell (2 for K-shell atoms like H and He). When atoms bond, the total energy of the system decreases, making it more stable.
Atoms share their valence electrons with another atom. Both atoms get to “count” the shared electrons. This forms a Covalent Bond.
One atom donates its valence electrons to another atom. Ions (charged particles) are formed. This forms an Ionic Bond.
Covalent Bond — Bonding by Sharing of Electrons
A covalent bond is formed by the sharing of one or more pairs of electrons between two atoms. Each atom contributes electrons to the shared pair, and the shared electrons are attracted to both nuclei.
🔵 Types of Covalent Bonds
One pair of electrons shared between two atoms. Example: H—H, H—Cl, Cl—Cl
Two pairs of electrons shared between two atoms. Example: O=O (oxygen molecule)
Three pairs of electrons shared. Example: N≡N (nitrogen molecule)
⚗️ Formation of Molecules of Elements
H has 1 electron, needs 1 more to complete K-shell (max = 2). Two H atoms each share 1 electron → H—H (single bond). Formula: H₂
Cl has 7 valence electrons, needs 1 more. Two Cl atoms each share 1 electron → Cl—Cl (single bond). Formula: Cl₂
+
H atom
→
H—H (H₂)
O
O atom
+
O
O atom
double bond
O=O
O₂ molecule
🧪 Formation of Compound Molecules
| Compound | Atoms Involved | Electrons Needed Each | Bond Type | Formula |
|---|---|---|---|---|
| Hydrogen Chloride | H (1e⁻) + Cl (7e⁻) | H needs 1; Cl needs 1 | Single (H—Cl) | HCl |
| Water | 2H + O (6e⁻) | O needs 2; each H needs 1 | 2 single bonds | H₂O |
| Ammonia | N (5e⁻) + 3H | N needs 3; each H needs 1 | 3 single bonds | NH₃ |
| Carbon dioxide | C (4e⁻) + 2O | C needs 4; each O needs 2 | 2 double bonds | CO₂ |
📛 Naming Covalent Compounds
mono(1), di(2), tri(3), tetra(4), penta(5), hexa(6), hepta(7), octa(8). First element keeps its name; second element ends in -ide. Mono- is usually omitted for the FIRST element.
| Formula | IUPAC Name | Trick to Remember |
|---|---|---|
| CO | Carbon monoxide | mono- used for second element O |
| CO₂ | Carbon dioxide | di = 2 oxygen atoms |
| CS₂ | Carbon disulfide | di = 2 sulfur atoms |
| PCl₃ | Phosphorus trichloride | tri = 3 chlorine atoms |
| SF₆ | Sulfur hexafluoride | hexa = 6 fluorine atoms |
| N₂O₄ | Dinitrogen tetroxide | both elements use prefix |
| H₂S | Hydrogen sulfide | No prefix before hydrogen! |
| H₂O | Water (common name) | Official: hydrogen monoxide |
| NH₃ | Ammonia (common name) | Official: nitrogen trihydride |
If a prefix ends in ‘a’ or ‘o’ AND the element name starts with a vowel → drop the last vowel. Example: penta + oxide = pentoxide (not pentaoxide), mono + oxide = monoxide (not monooxide).
Ionic Bond — Bonding by Transfer of Electrons
An ionic bond is the electrostatic force of attraction between oppositely charged ions (cation and anion) that holds them together. It forms when one atom transfers electrons to another atom.
🧂 Classic Example: Formation of NaCl (Common Salt)
Sodium (Z=11): Electronic config = 2, 8, 1. Has 1 valence electron — it donates this to get stable 2,8 configuration. After losing 1 electron: 11 protons, 10 electrons → Net charge = +1 → Forms Na⁺ cation
Chlorine (Z=17): Electronic config = 2, 8, 7. Has 7 valence electrons — it accepts 1 to get stable 2,8,8 configuration. After gaining 1 electron: 17 protons, 18 electrons → Net charge = −1 → Forms Cl⁻ anion
(2,8,1)
Sodium atom
→ e⁻ →
Cl
(2,8,7)
Chlorine atom
⟹
Na⁺
(2,8)
Cation (+)
+
Cl⁻
(2,8,8)
Anion (−)
→ NaCl
Cation (धनायन): Positively charged ion (lost electrons). Metals form cations. Example: Na⁺, Ca²⁺, Fe³⁺
Anion (ऋणायन): Negatively charged ion (gained electrons). Non-metals form anions. Example: Cl⁻, O²⁻, S²⁻
Together, cations and anions are called ions (आयन).
🏗️ Crystal Structure of Ionic Compounds
Ionic compounds do not form simple molecules. Instead, they form three-dimensional (3-D) crystal structures. In NaCl, each Na⁺ is surrounded by 6 Cl⁻ and each Cl⁻ is surrounded by 6 Na⁺, forming a regular repeating pattern called a crystal lattice.
📋 Naming Ionic Compounds
- Name the cation first, then the anion.
- Names of simple anions end with -ide (chloride, oxide, sulfide).
- Polyatomic ions generally do NOT end with -ide (sulfate, carbonate, nitrate).
- Examples: NaCl = Sodium chloride; CaO = Calcium oxide; Na₂S = Sodium sulfide
| Type | Name of Ion | Formula | Valency/Charge |
|---|---|---|---|
| Common Cations (+) | Sodium | Na⁺ | 1 |
| Potassium | K⁺ | 1 | |
| Calcium | Ca²⁺ | 2 | |
| Magnesium | Mg²⁺ | 2 | |
| Aluminium | Al³⁺ | 3 | |
| Iron (Ferrous) | Fe²⁺ | 2 | |
| Iron (Ferric) | Fe³⁺ | 3 | |
| Common Anions (−) | Chloride | Cl⁻ | 1 |
| Oxide | O²⁻ | 2 | |
| Sulfide | S²⁻ | 2 | |
| Hydroxide | OH⁻ | 1 | |
| Sulfate | SO₄²⁻ | 2 | |
| Polyatomic Ions | Carbonate | CO₃²⁻ | 2 |
| Nitrate | NO₃⁻ | 1 | |
| Hydrogencarbonate | HCO₃⁻ | 1 | |
| Ammonium | NH₄⁺ | 1 |
Writing Chemical Formulae — The Criss-Cross Method
There’s a quick method to write chemical formulae using the criss-cross method: write the symbols and their valencies, then swap (cross) the valencies as subscripts.
(1) If subscripts have a common factor, simplify (e.g., Mg₂O₂ → MgO). (2) Use brackets when more than one polyatomic ion is present: Mg(OH)₂, not MgOH₂. (3) Charges are NOT written in the final formula.
📝 Examples: Covalent Compounds
📝 Examples: Ionic Compounds
| Compound | Cation (charge) | Anion (charge) | Criss-cross | Final Formula |
|---|---|---|---|---|
| Calcium chloride | Ca (2+) | Cl (1−) | Ca¹Cl² | CaCl₂ |
| Aluminium oxide | Al (3+) | O (2−) | Al²O³ | Al₂O₃ |
| Magnesium oxide | Mg (2+) | O (2−) | Mg²O² → simplify | MgO |
| Calcium carbonate | Ca (2+) | CO₃ (2−) | Ca¹(CO₃)¹ → simplify | CaCO₃ |
| Magnesium hydroxide | Mg (2+) | OH (1−) | Mg¹(OH)² | Mg(OH)₂ |
| Aluminium sulfate | Al (3+) | SO₄ (2−) | Al²(SO₄)³ | Al₂(SO₄)₃ |
| Aluminium hydroxide | Al (3+) | OH (1−) | Al¹(OH)³ | Al(OH)₃ NOT AlOH₃ |
Use brackets ( ) ONLY when you have 2 or more polyatomic ions (ions with more than one atom like OH⁻, SO₄²⁻, CO₃²⁻). Single polyatomic ions don’t need brackets. Example: NaOH (no brackets needed — just one OH⁻), but Mg(OH)₂ (brackets needed — two OH⁻ ions).
Properties of Ionic vs Covalent Compounds
⚡ Ionic Compounds (NaCl, CuSO₄)
- Generally soluble in water ✅
- Generally insoluble in organic solvents (kerosene, petrol) ❌
- Do NOT conduct electricity in solid state (ions are fixed) ❌
- DO conduct electricity when dissolved in water (ions are free to move) ✅
- Also conduct in molten (liquid) state ✅
- High melting and boiling points (strong ionic bonds)
- Form crystalline solids
🤝 Covalent Compounds (Camphor, Naphthalene, Sugar)
- Most are insoluble in water (except some like sugar) ❌
- Generally soluble in organic solvents (kerosene, petrol) ✅
- Do NOT conduct electricity in any state (no ions formed) ❌
- Even if dissolved in water (like sugar) — no ions in solution → no conductivity ❌
- Low melting and boiling points (weak intermolecular forces)
- Can be liquids or gases at room temperature
NaCl (ionic) dissolves in water to give free Na⁺ and Cl⁻ ions. These ions can carry electric charge → conducts electricity. Sugar (covalent) dissolves but does NOT break into ions — just sugar molecules in water. No ions → no conductivity!
Ionic compounds DO NOT conduct electricity in the solid state even though they have ions. Why? Because in solid state, the ions are held in fixed positions in the crystal lattice and cannot move. Conductivity requires FREE-MOVING ions.
Molecular Mass & Formula Unit Mass
🔵 Molecular Mass (आण्विक द्रव्यमान) — for Covalent Compounds
Molecular mass = sum of atomic masses of ALL atoms present in ONE molecule of the compound. Unit: u (unified atomic mass unit)
H = 1 u; O = 16 u
MM = (1 × 2) + (16 × 1)
= 2 + 16
= 18 u
C = 12 u; O = 16 u
MM = (12 × 1) + (16 × 2)
= 12 + 32
= 44 u
🔶 Formula Unit Mass (सूत्र इकाई द्रव्यमान) — for Ionic Compounds
Ionic compounds don’t form molecules — they form 3-D crystals. The simplest whole number ratio of ions is called a formula unit. The mass of one formula unit = formula unit mass.
Na = 23 u; O = 16 u
FUM = (23 × 2) + (16 × 1)
= 46 + 16
= 62 u
Ca=40; N=14; O=16
FUM = 40 + [(14 + 48) × 2]
= 40 + 124
= 164 u
| Compound | Type | Atoms/Ions | Calculation | Mass |
|---|---|---|---|---|
| H₂O (water) | Covalent → Molecular Mass | 2H + 1O | (1×2)+(16×1) | 18 u |
| CO₂ | Covalent → Molecular Mass | 1C + 2O | (12×1)+(16×2) | 44 u |
| CH₄ (methane) | Covalent → Molecular Mass | 1C + 4H | (12×1)+(1×4) | 16 u |
| HNO₃ (nitric acid) | Covalent → Molecular Mass | 1H+1N+3O | 1+14+(16×3) | 63 u |
| NaCl (common salt) | Ionic → Formula Unit Mass | 1Na + 1Cl | 23+35.5 | 58.5 u |
| Na₂O | Ionic → Formula Unit Mass | 2Na + 1O | (23×2)+16 | 62 u |
| Mg(OH)₂ | Ionic → Formula Unit Mass | 1Mg+2O+2H | 24+(16+1)×2 | 58 u |
H=1, C=12, N=14, O=16, Na=23, Mg=24, Al=27, S=32, Cl=35.5, K=39, Ca=40, Fe=56, Cu=64, Zn=65
Law: Matter can neither be created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.
Experiment (Vinegar + Baking Soda):
Set-up: A conical flask with vinegar and a balloon filled with baking soda are weighed together. The balloon is fixed to the flask mouth. Initial mass is recorded. The baking soda is allowed to fall into vinegar — CO₂ is produced and captured in the balloon. The system is weighed again.
Observation: Initial mass = Final mass.
Conclusion: Mass is conserved in the chemical reaction. In an open system, if CO₂ escapes, apparent mass decreases — but total mass including the gas is still constant.
Sodium (Na, Z=11): Electronic config = 2, 8, 1. Has 1 valence electron. To achieve stable octet (2,8), it loses this electron → Na becomes Na⁺ cation (11 protons, 10 electrons, net charge +1).
Chlorine (Cl, Z=17): Electronic config = 2, 8, 7. Has 7 valence electrons. To achieve stable octet (2,8,8), it gains 1 electron → Cl becomes Cl⁻ anion (17 protons, 18 electrons, net charge −1).
Bond Formation: The oppositely charged Na⁺ and Cl⁻ are held together by electrostatic force of attraction → this is an ionic bond (आयनिक बंध). It forms because Na (less than 4 valence electrons) donates its electron to Cl (more than 4 valence electrons), allowing both to achieve stability.
(ii) Calcium carbonate: Ca²⁺ and CO₃²⁻ → same valency: Ca¹(CO₃)¹ → simplify → CaCO₃
(iii) Ferric oxide: Fe³⁺ (ferric = Fe³⁺) and O²⁻ → Criss-cross: Fe²O³ → Fe₂O₃
(iv) Magnesium hydroxide: Mg²⁺ and OH⁻ → Criss-cross: Mg¹(OH)² → Mg(OH)₂
(i) Formation: Ionic compounds form by transfer of electrons from one atom to another, forming ions (cations and anions). Covalent compounds form by sharing of electrons between atoms, forming molecules.
(ii) Solubility: Ionic compounds are generally soluble in water but insoluble in organic solvents like kerosene and petrol. Covalent compounds are generally insoluble in water but soluble in organic solvents.
(iii) Electrical conductivity: Ionic compounds do not conduct electricity in solid state (ions are fixed) but conduct in aqueous solution or molten state (ions are free to move). Covalent compounds generally do not conduct electricity in any state as they do not form ions.
(ii) CH₄: = 12 + (1×4) = 12 + 4 = 16 u
(iii) KCl: = 39 + 35.5 = 74.5 u
(iv) Ca(NO₃)₂: = 40 + [(14 + 16×3) × 2] = 40 + [62 × 2] = 40 + 124 = 164 u
Remember 2 key laws: Conservation of Mass (Lavoisier, 1789) and Constant Proportions (Proust). Learn Dalton’s 6 postulates — especially that atoms combine in simple whole numbers. Covalent = sharing = molecules. Ionic = transfer = ions + crystal lattice. For formulae: criss-cross the valencies, simplify if needed, use brackets for polyatomic ions. Ionic compounds conduct in solution but NOT in solid state. All the best! 🚀
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