These atomic foundations of matter class 9 notes compress the full chapter into a revision-ready reference: the Law of Conservation of Mass, the Law of Constant Proportions, Dalton’s postulates, covalent and ionic bonding, naming and formula writing, compound properties, and mass calculations.
Revise in this order: mass laws → Dalton’s explanation → bonds → formulae → numerics. This page is not a solutions bank — NCERT exercise solutions and MCQs live in separate resources.
Chapter 9 at a Glance: From Mass Laws to Chemical Bonds
This chapter builds one idea on another. The teaching order matters more than the textbook’s section order:
- Weighing experiments → Law of Conservation of Mass (mass is never lost in a reaction).
- Fixed ratios by mass → Law of Constant Proportions (compounds always combine in the same mass ratio).
- Dalton’s Atomic Theory explains both laws with atoms that rearrange but never vanish.
- Molecules form in two ways — covalent bonding (sharing electrons) and ionic bonding (transferring electrons).
- Naming compounds and the criss-cross method for writing formulae.
- Properties of ionic vs covalent compounds (solubility, conductivity, melting point).
- Calculating molecular mass and formula unit mass.
Before and after: In Chapter 8, Journey Inside the Atom, you learnt the octet rule, valence shell and electronic configuration. Chapter 9 applies those ideas to explain why atoms bond, how compounds form, and how their masses are calculated.
Key Terms Explained: Definitions Table with Examples
Here is the chapter’s vocabulary in student-friendly words (NCERT, pp. 169–180).
| Term | Meaning | Example |
|---|---|---|
| Molecule | An electrically neutral group of more than one atom that exists independently and shows the substance’s properties | H₂, HCl |
| Ion | An atom or group of atoms carrying a net electric charge | Na⁺, Cl⁻ |
| Cation | A positively charged ion, formed by losing electrons | Na⁺, Ca²⁺ |
| Anion | A negatively charged ion, formed by gaining electrons | Cl⁻, O²⁻ |
| Chemical bond | The force holding atoms together; bonding lowers the total energy of the system | Bond in NaCl |
| Covalent bond | A bond formed by sharing a pair of electrons between atoms | H—H in H₂ |
| Single bond | One shared electron pair between two atoms | H—H |
| Double bond | Two shared electron pairs between two atoms | O=O in O₂ |
| Ionic bond | Electrostatic attraction between oppositely charged ions | Na⁺ and Cl⁻ in NaCl |
| Valency | The combining capacity of an atom, equal to its charge or number of bonds it can form | Na = 1, O = 2 |
| Polyatomic ion | An ion made of atoms of two or more elements | NH₄⁺, SO₄²⁻, CO₃²⁻ |
| Formula unit | The simplest whole-number ratio of ions in an ionic compound | NaCl, not Na₂Cl₂ |
| Molecular mass | Sum of atomic masses of all atoms in a molecule | H₂O = 18 u |
| Formula unit mass | Sum of atomic masses of all atoms in one formula unit | Na₂O = 62 u |
Law of Conservation of Mass: Open and Closed Systems
Lavoisier proposed in 1789 that matter can neither be created nor destroyed in a chemical reaction (NCERT, p. 166). The experiments prove it in three setups.
Physical change (Activity 9.1): Dissolve salt in water — the solution’s mass equals the sum of the separate masses. No change in a physical change.
Chemical change, open flask (set-up 1): Vinegar + baking soda (sodium hydrogencarbonate) produces carbon dioxide. In an open flask the gas escapes, so the final reading drops. This looks like a mass loss but is only the gas leaving.
Chemical change, sealed balloon (set-up 2): Fix the balloon to the flask mouth so CO₂ inflates it instead of escaping. The final reading matches the initial reading — mass is conserved.
Group verification (Activity 9.3): Sodium sulfate + barium chloride forms a white precipitate of barium sulfate. No gas forms, so both flasks stay on the balance and the reading is unchanged.
Worked Example: Finding the Missing Reactant Mass
Method: Law of Conservation of Mass.
Step 1: 6.0 g of magnesium burns in a closed crucible to give 9.9 g of magnesium oxide.
Oxygen must have combined with the magnesium.
- Step 1: Mass of oxygen = mass of product − mass of magnesium = 9.9 g − 6.0 g = 3.9 g.
- Step 2: Verify: reactants (6.0 g Mg + 3.9 g O = 9.9 g) equal product (9.9 g MgO).
✓ Final answer: 3.9 g of oxygen combined with the magnesium.
NCERT’s Example 9.1 (calcium carbonate + hydrochloric acid) and Example 9.2 (carbon + oxygen) show the same logic with their own numbers — rework them yourself to practise.
Law of Constant Proportions: Why Water Is Always 1:8
Joseph Proust stated that a compound always contains the same elements in the same ratio by mass, whatever its source or method of preparation (NCERT, p. 167). Called the Law of Definite Proportions or Proust’s Law.
Anchor example: Water from any source — river, borewell, ocean — decomposes to hydrogen and oxygen in the mass ratio 1:8. So 9 g of purified water gives 1 g H and 8 g O.
Supporting evidence:
- Cinnabar: mercury 86.22% : sulfur 13.78% by mass.
- Sodium chloride: sodium : chlorine = 23 : 35.5 by mass.
Compounds, not mixtures: The law holds for compounds because atoms combine in fixed ratios. Mixtures can vary — this is why the law fails for them (Pause and Ponder Q5).
Worked Example: How Much Oxygen Combines with Hydrogen?
Method: Constant Proportions ratio.
Step 1: Water’s mass ratio is H:O = 1:8.
Every 1 g of hydrogen needs 8 g of oxygen.
- Step 1: For 4.5 g of hydrogen: oxygen needed = 8 × 4.5 g = 36 g.
- Step 2: Water formed = 4.5 g + 36 g = 40.5 g.
Final answer: 36 g of oxygen combines with 4.5 g of hydrogen to form 40.5 g of water.
Dalton’s Atomic Theory: Six Postulates and What They Explain
John Dalton presented his atomic theory in 1808 (NCERT, p. 168). Its postulates:
- All matter is made of tiny particles called atoms, which take part in chemical reactions.
- Atoms are indivisible — they cannot be created or destroyed in a 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 simple whole-number ratios to form compounds.
- The relative number and kinds of atoms are constant in a given compound.
How the postulates explain the laws:
- Postulate 2 (indestructible atoms) → Law of Conservation of Mass: atoms merely rearrange, so mass is conserved.
- Postulates 5 and 6 (fixed combining ratios) → Law of Constant Proportions: the same ratio always appears.
Misconception autopsy: Later experiments (Chapter 8) revealed electrons, protons and neutrons, proving atoms are divisible. Dalton’s “indivisible atoms” postulate was refined — but the two laws themselves were never broken. The octet rule and electronic configurations you studied in Chapter 8 build directly on this.
Covalent Bonds: Sharing Electrons to Form Molecules
A molecule is an electrically neutral entity of more than one atom that exists independently and shows the substance’s properties (NCERT, p. 169). Atoms with a partially filled valence shell share electrons to become stable — this sharing forms a covalent bond.
H₂ (hydrogen): H has one electron in the K-shell, which holds two. Two H atoms share one electron each → single bond, written H—H.
Cl₂ (chlorine): Each Cl has 7 valence electrons and needs 1. Two share one pair → single bond, Cl—Cl.
O₂ (oxygen): O has 6 valence electrons and needs 2. Two O atoms share two pairs → double bond, written O=O.
HCl (hydrogen chloride): H needs 1 electron (duplet), Cl needs 1 (octet). Both share one → single bond, H—Cl. This is a covalent compound.
H₂O (water): O needs 2 electrons; each H needs 1. Two H atoms each share one electron with the single O atom → H₂O, showing two H atoms and one O atom.
Why helium forms no molecule: helium’s only shell (K) is already full with 2 electrons, so it is stable alone. It exists as single atoms (NCERT, p. 169).
Ionic Bonds: Transferring Electrons to Form Ions
When one atom transfers electrons to another, the result is an ionic bond — the electrostatic attraction between oppositely charged ions (NCERT, p. 173).
Na → Na⁺: Na (2,8,1) loses one electron. It now has 11 protons and 10 electrons, so it carries a +1 charge.
Cl → Cl⁻: Cl (2,8,7) gains one electron to complete its octet, carrying a −1 charge.
\[ \mathrm{Na} + \mathrm{Cl} \rightarrow \mathrm{Na}^{+} + \mathrm{Cl}^{-} \rightarrow \mathrm{NaCl} \]
Crystal lattice: Ionic compounds form 3-D crystals. In NaCl, each Na⁺ is surrounded by six Cl⁻ and each Cl⁻ by six Na⁺ in a repeating pattern. Ions are arranged as points in a crystal lattice (NCERT, p. 174).
Two-electron example: Sulfur has 6 valence electrons and needs 2 to complete its octet. Gaining two electrons gives S²⁻.
Common Ions and Their Valencies
Condensed from NCERT Table 9.1 (pp. 174–175).
| Type | Ions (valency) |
|---|---|
| Monoatomic cations | Na⁺ (1), K⁺ (1), Li⁺ (1), Ag⁺ (1), Ca²⁺ (2), Ba²⁺ (2), Mg²⁺ (2), Zn²⁺ (2), Fe²⁺ (2), Fe³⁺ (3), Cu⁺ (1), Cu²⁺ (2), Al³⁺ (3) |
| Monoatomic anions | F⁻ (1), Cl⁻ (1), Br⁻ (1), I⁻ (1), O²⁻ (2), S²⁻ (2) |
| Polyatomic ions | OH⁻ (1), NO₃⁻ (1), HCO₃⁻ (1), NH₄⁺ (1), CO₃²⁻ (2), SO₄²⁻ (2) |
Naming Compounds: Prefixes, -ide Endings and Common Names
Covalent compounds (NCERT, pp. 172–173): first element keeps its name, second element ends in -ide. Prefixes show atom counts.
- mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6).
- mono- is omitted for the first element — CO₂ is carbon dioxide, not monocarbon dioxide.
- mono- is kept for the second element — CO is carbon monoxide, not carbon monooxide.
- Drop ‘o’ or ‘a’ before a vowel: monoxide, pentoxide. Keep ‘i’: dioxide, trioxide.
- No prefix is added to hydrogen: H₂S is hydrogen sulfide, not dihydrogen sulfide.
- Common names win: H₂O is water (not hydrogen monoxide), NH₃ is ammonia (not nitrogen trihydride).
Ionic compounds (NCERT, p. 174): cation name first, then anion. Simple anions end in -ide; polyatomic ions keep their names (sulfate, nitrate, carbonate, hydroxide, ammonium).
Memory device for prefixes: “My Dear Teacher, Please Pass Honey” — M-D-T-P-P-H = mono, di, tri, tetra, penta, hexa.
Memory hook for the three laws: “Lavoisier Locks Mass, Proust Pins Proportions, Dalton Draws Atoms.”
Writing Chemical Formulae: The Criss-Cross Method
Two methods, depending on the bond type (NCERT, pp. 175–177).
Method 1 — Covalent compounds
- Write the symbols of the elements.
- Write the valencies (from Table 9.1).
- Crossover the valencies as subscripts.
- A valency of 1 is not written.
Examples: HCl (1,1 → no subscript), H₂S (H₂S₁ → H₂S), CCl₄ (C₁Cl₄ → CCl₄).
Method 2 — Ionic compounds
- Write the cation symbol first, then the anion.
- Write the charges under the symbols, not as superscripts.
- Crossover only the numbers as subscripts.
- Simplify subscripts by their common factor (Mg₂O₂ → MgO; Ca₂(CO₃)₂ → CaCO₃).
- Use brackets when two or more polyatomic ions appear: Al(OH)₃, Al₂(SO₄)₃ — never AlOH₃.
- Charges are not shown in the final formula.
Worked Example: Ferric Sulfate
Method: Criss-cross with brackets.
- Step 1: Cation Fe³⁺ first, anion SO₄²⁻ second.
- Step 2: Crossover charges: 3 from iron goes to sulfate, 2 from sulfate goes to iron.
- Step 3: Fe₂(SO₄)₃ — the bracket groups three sulfate ions.
Final answer: Fe₂(SO₄)₃.
Ionic vs Covalent Compounds: Solubility, Conductivity, Melting Point
Activity 9.4 (NCERT, pp. 178–179) compares camphor, sodium chloride, copper sulfate, sugar and naphthalene.
| Property | Ionic compounds (NaCl, CuSO₄) | Covalent compounds (camphor, naphthalene, sugar) |
|---|---|---|
| Bond formation | Transfer of electrons | Sharing of electrons |
| Melting / boiling point | High, due to strong inter-ionic attractions | Low |
| Solubility in water | Generally soluble | Insoluble (sugar is an exception — soluble) |
| Solubility in kerosene / petrol | Insoluble | Soluble |
| Conductivity in solid state | No — ions fixed in lattice | No — no free ions |
| Conductivity in aqueous solution | Yes — ions free to move | No — no ions formed (sugar stays neutral) |
| Conductivity in molten state | Yes (predict: ions free to move when melted) | No |
Why the difference? In ionic solids, ions are locked in a lattice and cannot move, so no conduction. Once dissolved or melted, ions are free — they conduct. Covalent compounds have no ions at all, so even dissolved sugar (neutral molecules) does not conduct.
Real-life application: Salt water conducts electricity because NaCl splits into Na⁺ and Cl⁻ ions; sugar water does not because sugar stays as neutral molecules. In your body, dissolved electrolytes — sodium and potassium ions — carry electrical signals along nerves; this is why electrolyte balance matters for muscle and nerve function.
Analogy: A covalent bond is like two friends sharing a pen — both use it, neither owns it alone. An ionic bond is like one friend handing over a coin: one gains (becomes negative), one loses (becomes positive), and the electrostatic pull between them holds the pair together.
Pause and Ponder Q19 asks which bond a solid has if it conducts only when dissolved in water. The answer is an ionic bond — the ions need freedom of movement to carry current.
Molecular Mass and Formula Unit Mass: Formulas with Worked Examples
Molecular mass = sum of (atomic mass × number of atoms) for every element in a molecule, expressed in atomic mass units (u) (NCERT, p. 180).
Formula unit mass applies to ionic compounds: the sum of atomic masses in one formula unit (simplest whole-number ratio of ions), because ionic compounds form 3-D lattices, not molecules (NCERT, p. 180).
Known anchors from NCERT: water = 18 u (Example 9.4), CO₂ = 44 u (Example 9.5), sodium oxide = 62 u (Example 9.6), calcium nitrate = 164 u (Example 9.7).
Worked Example: Formula Unit Mass of Potassium Sulfate
Method: Sum of atomic masses, multiplied per atom.
Step 1: Formula K₂SO₄.
Atomic masses: K = 39 u, S = 32 u, O = 16 u.
- Step 1: Potassium: 39 u × 2 = 78 u.
- Step 2: Sulfur: 32 u × 1 = 32 u.
- Step 3: Oxygen: 16 u × 4 = 64 u.
- Step 4: Total = 78 + 32 + 64 = 174 u.
Final answer: Formula unit mass of K₂SO₄ = 174 u.
Practise on your own: KCl (Pause and Ponder Q23) and Mg(OH)₂ (Q24) use the same full-multiplication method.
Common Mistakes to Avoid in This Chapter
| Students write X | Correct is Y | How to check |
|---|---|---|
| Open beaker mass drop means the law is violated | The CO₂ gas escaped; the closed balloon setup keeps mass constant | Repeat with a sealed balloon — readings match |
| AlOH₃ | Al(OH)₃ — three hydroxide ions need brackets | Count OH groups: one O and three H is wrong |
| Mg₂O₂ | MgO — divide subscripts by common factor 2 | Simplest ratio: 1:1, not 2:2 |
| Molecular mass of NaCl | Formula unit mass — NaCl forms a crystal lattice, not molecules | Ionic compound → use formula unit mass |
| Monocarbon dioxide | Carbon dioxide — mono- is omitted for the first element | First element: no prefix |
| Chloride sodium | Sodium chloride — cation is always named first | Metal (positive) first, non-metal (negative) second |
Exam Notes: How Questions on This Chapter Are Asked
Observed patterns from the chapter’s own exercises (Revise, Reflect, Refine, pp. 181–183).
- Assertion-reason items built on the laws — e.g. 2 g hydrogen + 16 g oxygen → 18 g water (Q14), and copper sulfate’s conductivity in molten vs solid state (Q14).
- Writing formulae from given ion pairs — Q6 (Ca²⁺/Br⁻, Al³⁺/CO₃²⁻, K⁺/SO₃²⁻, NH₄⁺/Cl⁻). The crossover-then-simplify step earns the mark.
- Numericals on molecular and formula unit mass — Q8 (ammonium nitrate, phosphoric acid, sodium hydrogencarbonate). Show the full multiplication with units.
- Explain-the-result questions on the vinegar-baking soda balloon experiment, open vs closed system — Q11.
- Species identification from proton/neutron/electron counts, deciding neutral vs cation vs anion — Q12, Q15.
- Visual items — choosing the correct diagram of Cl⁻ (Q7).
Full worked answers to these exercises are in the separate NCERT solutions resource — this page is for concept revision only.
Atomic Foundations of Matter Class 9 Notes: One-Page Recap
| Law / Concept | One-line statement | Who / key numbers |
|---|---|---|
| Conservation of Mass | Mass is neither created nor destroyed in a reaction | Lavoisier, 1789 |
| Constant Proportions | Elements in a compound combine in a fixed mass ratio | Proust; water 1:8; NaCl 23:35.5 |
| Dalton’s postulates | Atoms are tiny, indivisible, identical per element, combine in whole-number ratios | Dalton, 1808 |
| Covalent bond | Sharing of electron pairs | H—H single, O=O double |
| Ionic bond | Transfer of electrons → electrostatic attraction | Na⁺ + Cl⁻ → NaCl |
| Ionic vs covalent | Ionic: high melting, conducts when molten/dissolved. Covalent: low melting, no conduction | NaCl vs camphor |
| Molecular vs formula unit mass | Sum of atomic masses of all atoms in the molecule / formula unit | H₂O = 18 u; K₂SO₄ = 174 u |
For more revision material, browse the Class 9 Science notes hub and the full CBSE notes library. You can verify every figure and page reference in the official NCERT Chapter 9 PDF.
Reference: NCERT Class 9 Science textbook, chapter Atomic Foundations of Matter.
Frequently Asked Questions About Atomic Foundations of Matter
Why does the mass reading drop when baking soda reacts with vinegar in an open flask?
The reaction produces carbon dioxide gas. In an open flask the gas escapes into the air, so the balance reads less. The law is not violated — mass is conserved, but the measurement excludes the escaped gas. Sealing a balloon over the flask traps the CO₂, and the final reading matches the initial one (NCERT, p. 165).
What is the difference between molecular mass and formula unit mass?
Molecular mass is the sum of atomic masses in a covalent molecule. Formula unit mass is the same sum for one formula unit of an ionic compound, which exists as a 3-D crystal lattice rather than a molecule. Water uses molecular mass (18 u); sodium chloride uses formula unit mass (NCERT, p. 180).
How can I quickly decide whether a compound is ionic or covalent?
Check the participating elements. Metals (which lose electrons) with non-metals form ionic bonds — e.g. NaCl, CaCl₂. Non-metals with non-metals share electrons and form covalent bonds — e.g. H₂O, CO₂, HCl. If the compound contains polyatomic ions like SO₄²⁻ or CO₃²⁻ with a metal, it is ionic (NCERT, p. 174).
Why does salt solution conduct electricity but sugar solution does not?
Salt (NaCl) dissociates into free-moving Na⁺ and Cl⁻ ions in water, which carry current. Sugar dissolves as neutral molecules — no ions are produced, so no current flows. Ionic compounds conduct only when ions are free to move, which happens when dissolved or molten (NCERT, p. 179).
How do I use the law of constant proportions in a numerical?
First write the fixed mass ratio from the compound’s identity — water is H:O = 1:8. Then scale: for 4.5 g of hydrogen, oxygen needed = 8 × 4.5 = 36 g, and water formed = 4.5 + 36 = 40.5 g. The ratio stays constant regardless of the sample size (NCERT, p. 167).
Why does helium exist as single atoms while hydrogen forms H₂ molecules?
Helium’s K-shell is already full with 2 electrons, so it is stable alone and needs no bonding. Hydrogen has only 1 electron in its K-shell, which holds 2 — it needs one more to complete its duplet, so two H atoms share one electron each to form H₂ (NCERT, p. 169).
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