Class 9 Science: Atomic Foundations of Matter — Chapter Question Set

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created; destroyed

Antoine Lavoisier; 1789

Proust's; Definite Proportions

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mass; chemical

sharing

transfer

cations; anions

ionic

1; 2; 3

cation; anion

solid; water; molten

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cation (Na⁺); positive (+1)

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The Law of Conservation of Mass states that matter can neither be created nor destroyed in a chemical reaction; the total mass of the reactants always equals the total mass of the products. For example, when 4.0 g of CaCO₃ reacts with 2.92 g of HCl, the products formed have a total mass of exactly 6.92 g, showing that mass is conserved.

The Law of Constant Proportions states that a chemical compound always contains the same elements combined in a fixed ratio by mass, regardless of its source or method of preparation. For example, water always contains hydrogen and oxygen in the mass ratio 1:8, whether it comes from a river, a well or a laboratory.

When baking soda reacts with vinegar in an open container, carbon dioxide gas is produced as one of the products. This gas escapes into the surrounding air, so the measured mass of the substances left in the container decreases, even though the total mass of all matter is actually conserved.

According to Dalton's atomic theory, (i) all matter is made up of tiny, indivisible particles called atoms, and (ii) atoms of a given element are identical in mass and properties, while atoms of different elements differ from one another. Also, atoms combine in simple whole-number ratios to form compounds — any two of these postulates answer the question.

A molecule is a neutral group of two or more atoms that are chemically bonded together and capable of independent existence, showing all the properties of the substance. A molecule of an element is H₂ (hydrogen), and a molecule of a compound is H₂O (water).

A covalent bond is formed when two atoms share one or more pairs of valence electrons between them, allowing each atom to attain a stable electron arrangement. For example, two hydrogen atoms each share one electron with the other, forming a single covalent bond in the hydrogen molecule (H₂).

A single covalent bond is formed when two atoms share one pair of electrons, as in the hydrogen molecule (H–H in H₂). A double covalent bond is formed when two atoms share two pairs of electrons, as in the oxygen molecule (O=O in O₂), and is stronger than a single bond.

An ionic bond is the electrostatic force of attraction between oppositely charged ions, formed when one atom transfers electrons to another. For example, sodium loses one electron to become Na⁺ and chlorine gains that electron to become Cl⁻; these ions attract each other to form sodium chloride (NaCl).

A cation is a positively charged ion formed when an atom loses one or more electrons, such as Na⁺ formed from a sodium atom. An anion is a negatively charged ion formed when an atom gains one or more electrons, such as Cl⁻ formed from a chlorine atom.

In the solid state, the ions of an ionic compound are held in fixed positions within the crystal lattice by strong electrostatic forces and cannot move freely. Since there are no mobile charge carriers to carry current, solid ionic compounds do not conduct electricity.

Water is a polar molecule that helps separate and stabilise the positive and negative ions of ionic compounds, so they dissolve easily. Most covalent compounds are non-polar and do not interact favourably with polar water molecules, so they remain insoluble in water.

Ionic compounds generally have high melting and boiling points due to strong electrostatic forces between ions, and they conduct electricity when dissolved in water or molten. Covalent compounds usually have low melting and boiling points because their intermolecular forces are weak, and they generally do not conduct electricity.

Molecular mass is the mass of a molecule, equal to the sum of the atomic masses of all the atoms present in it, expressed in atomic mass units (u). For CO₂, molecular mass = 12 + (16 × 2) = 12 + 32 = 44 u.

Formula unit mass is the sum of the atomic masses of all the ions present in the simplest whole-number ratio represented by the formula of an ionic compound. For Na₂O, formula unit mass = (23 × 2) + 16 = 46 + 16 = 62 u.

In the criss-cross method, the numerical value of the charge on each ion becomes the subscript of the other ion. For Ca²⁺ and Cl⁻, the charge 2 on calcium is transferred to chlorine and the charge 1 on chlorine is transferred to calcium, giving the formula CaCl₂.

(i) CO₂ is named carbon dioxide, as one carbon atom shares electrons with two oxygen atoms. (ii) PCl₃ is named phosphorus trichloride, with one phosphorus atom bonded to three chlorine atoms. (iii) SF₆ is named sulfur hexafluoride, with one sulfur atom bonded to six fluorine atoms.

(i) Sodium hydrogencarbonate has the formula NaHCO₃, made up of Na⁺ and HCO₃⁻ ions. (ii) Ferric chloride has the formula FeCl₃, obtained by the criss-cross method from Fe³⁺ and three Cl⁻ ions.

Since 12 g of carbon produces 44 g of CO₂, the mass of CO₂ formed from 2.4 g of carbon is found by simple proportion: (44/12) × 2.4 = 8.8 g. Therefore, 8.8 g of CO₂ is formed from 2.4 g of carbon.

Since NaCl contains sodium and chlorine in the fixed mass ratio 23:35.5, the mass of chlorine needed for 46 g of sodium is (35.5/23) × 46 = 71 g. Hence, 71 g of chlorine is required to react completely with 46 g of sodium.

A monoatomic ion consists of a single atom that has lost or gained electrons, such as Na⁺ or Cl⁻. A polyatomic ion is a group of two or more bonded atoms that together carry a net charge, such as OH⁻ or SO₄²⁻.

Introduction: The vinegar-baking soda experiment shows that mass is conserved when no gas escapes.
Key Points:
1. Set-up 1 (open) — Baking soda and vinegar react in an open flask; CO₂ gas escapes and mass appears to decrease.
2. Set-up 2 (closed) — A balloon seals the flask; gas is trapped and the total mass remains unchanged after reaction.
3. Conclusion — Mass is conserved in chemical reactions — the apparent loss in set-up 1 was only escaping gas.
Conclusion: This confirms Lavoisier's Law of Conservation of Mass when a closed system is used.

Introduction: The Law of Constant Proportions states that a compound's elements always combine in a fixed mass ratio.
Key Points:
1. Statement — In any compound, the elements always combine in a fixed ratio by mass, regardless of source.
2. Water example — Water from any source (river, ocean, borewell) always contains H:O = 1:8 by mass.
3. Implication — Decomposing 9 g of water always gives 1 g of hydrogen and 8 g of oxygen.
Conclusion: This fixed ratio, observed by Joseph Proust, holds for all compounds and for no mixtures.

Introduction: Dalton's atomic theory explains both conservation of mass and constant proportions.
Key Points:
1. Key postulates — Atoms are indivisible, cannot be created or destroyed, and combine in whole-number ratios.
2. Conservation of Mass — Since atoms are neither created nor destroyed, the total mass of reactants equals that of products.
3. Constant Proportions — Atoms combine in fixed whole-number ratios, so a compound always has the same mass proportion of elements.
Conclusion: Thus Dalton's theory provided the logical foundation for both fundamental laws of chemical combination.

Introduction: Covalent bonding by electron sharing forms the molecules of gaseous elements.
Key Points:
1. Hydrogen molecule — Each H atom has 1 electron in K-shell; sharing one electron each completes the duplet, forming H–H (H₂).
2. Chlorine molecule — Each Cl atom has 7 valence electrons; sharing one electron each completes the octet, forming Cl–Cl (Cl₂).
3. Bond type — Both are single covalent bonds, depicted by a single line between the two atom symbols.
Conclusion: The shared pair of electrons is attracted by both nuclei, making each molecule stable.

Introduction: NaCl forms when sodium transfers its valence electron to chlorine, creating oppositely charged ions.
Key Points:
1. Sodium's action — Na (2,8,1) loses its one valence electron to achieve a stable octet, becoming Na⁺.
2. Chlorine's action — Cl (2,8,7) gains that electron to complete its octet, becoming Cl⁻.
3. Bond formation — The electrostatic attraction between Na⁺ and Cl⁻ forms the ionic bond in NaCl.
Conclusion: This electron-transfer mechanism is the basis of all ionic compounds formed between metals and non-metals.

Introduction: Ionic and covalent compounds differ sharply in their physical properties.
Key Points:
1. Melting/boiling points — Ionic compounds have high values; covalent compounds usually have low values.
2. Solubility — Ionic compounds dissolve in water; most covalent compounds dissolve in kerosene/petrol, not water.
3. Conductivity — Ionic compounds conduct when dissolved or molten; covalent compounds generally do not conduct.
Conclusion: These differences arise from the strong electrostatic forces in ionic lattices versus weaker covalent interactions.

Introduction: Covalent compounds are named using prefixes that indicate the number of atoms of each element.
Key Points:
1. Rules — First element keeps its name; second ends in -ide; prefixes mono-, di-, tri-, tetra- indicate atom counts.
2. Mono- omission — 'Mono-' is omitted for the first element but used for the second: CO is carbon monoxide, not monocarbon monoxide.
3. Examples — CO₂ = carbon dioxide; PCl₃ = phosphorus trichloride; SF₆ = sulfur hexafluoride.
Conclusion: This system gives every covalent compound a unique and systematic name.

Introduction: The criss-cross method uses ion charges to determine the subscripts in ionic formulae.
Key Points:
1. CaCl₂ — Ca²⁺ and Cl⁻: cross 2 to Cl and 1 to Ca → Ca₁Cl₂ = CaCl₂ (1 omitted).
2. Al₂O₃ — Al³⁺ and O²⁻: cross 3 to O and 2 to Al → Al₂O₃.
3. Simplification — If subscripts share a common factor (e.g. Mg₂O₂), divide to get MgO.
Conclusion: The criss-cross method ensures the positive and negative charges always balance in the formula.

Introduction: Molecular mass is found by adding the atomic masses of all atoms in the molecule.
Key Points:
1. HNO₃ — Molecular mass = 1 + 14 + (16 × 3) = 1 + 14 + 48 = 63 u.
2. CH₄ — Molecular mass = 12 + (1 × 4) = 12 + 4 = 16 u.
3. Key idea — Count every atom in the formula, multiply by its atomic mass, and add up the products.
Conclusion: The molecular masses of nitric acid and methane are 63 u and 16 u respectively.

Introduction: Formula unit mass is the sum of atomic masses in the simplest ionic formula unit.
Key Points:
1. KCl — Formula unit mass = 39 + 35.5 = 74.5 u.
2. Mg(OH)₂ — Formula unit mass = 24 + 2 × (16 + 1) = 24 + 2 × 17 = 24 + 34 = 58 u.
3. Bracket meaning — The subscript 2 outside the bracket applies to every atom inside it (2 O and 2 H).
Conclusion: The formula unit masses of KCl and Mg(OH)₂ are 74.5 u and 58 u respectively.

Introduction: The total masses of reactants and products must match if mass is conserved.
Key Points:
1. Reactants total — 4.0 g + 2.92 g = 6.92 g of reactants.
2. Products total — 1.76 + 0.72 + 4.44 = 6.92 g of products.
3. Comparison — Mass of reactants = mass of products = 6.92 g, so the law is obeyed.
Conclusion: This calculation confirms that the Law of Conservation of Mass holds for the reaction.

Introduction: Electronic configurations determine whether atoms transfer or share electrons when bonding.
Key Points:
1. Element A — Configuration 2,8,1 means A has 1 valence electron it readily loses, forming a +1 cation.
2. Element B — Configuration 2,6 means B needs 2 electrons to complete its octet, forming a −2 anion.
3. Bond and formula — A transfers one electron to B (ionic bond); two A atoms are needed for each B atom → formula A₂B.
Conclusion: The compound A₂B is ionic, formed by electron transfer from metal A to non-metal B.

Introduction: Oxygen (six valence electrons) forms a diatomic molecule by sharing, and bonds with metals by electron transfer.
Key Points:
1. Diatomic X₂ — X (six valence electrons) shares two electrons with another X, forming a double bond (X=X, like O₂).
2. Bond type in X₂ — This is a double covalent bond since two pairs of electrons are shared between the atoms.
3. Compound with Y — Y (two valence electrons) donates both to X; one Y provides 2 electrons to one X → formula YX (like MgO).
Conclusion: X₂ is a covalent molecule with a double bond; with metal Y it forms an ionic compound YX.

Introduction: Checking whether the total mass of reactants equals the total mass of products verifies the law.
Key Points:
1. Reactants total — 5.3 g + 6.0 g = 11.3 g of reactants.
2. Products total — 2.2 g + 0.9 g + 8.2 g = 11.3 g of products.
3. Conclusion — Since reactant mass = product mass = 11.3 g, the Law of Conservation of Mass is verified.
Conclusion: The reaction confirms that mass is conserved when all products (including gases) are accounted for.

Introduction: Calcium (a metal) transfers electrons to oxygen (a non-metal), forming an ionic bond.
Key Points:
1. Calcium's action — Ca (2,8,8,2) loses its two valence electrons to achieve a stable octet, forming Ca²⁺.
2. Oxygen's action — O (2,6) gains those two electrons to complete its octet, forming O²⁻.
3. Formula and name — One Ca²⁺ balances one O²⁻; after criss-cross and simplification, the formula is CaO, named calcium oxide.
Conclusion: Calcium oxide is an ionic compound formed by complete transfer of two electrons from calcium to oxygen.

(a) In the open flask, the carbon dioxide gas produced during the reaction escaped into the air, so the measured mass of the remaining substances decreased.
(b) In the sealed set-up, the mass remained the same before and after mixing, because the balloon trapped the gas and nothing left the system.
(c) The word equation is: Vinegar + Baking soda → Carbon dioxide + Other substances.
(d) This experiment verifies the Law of Conservation of Mass, which states that matter can neither be created nor destroyed in a chemical reaction.

(a) The mass of oxygen = (60/40) × 20 = 30 g, since oxygen always makes up 60% of the compound's mass.
(b) The Law of Constant Proportions (Definite Proportions) allows this determination, as the compound always contains its elements in a fixed ratio by mass.
(c) No, a mixture would not follow a fixed ratio, because the elements in a mixture can be combined in any proportion — only a chemical compound has a constant composition.
(d) The mass of oxygen = (60/40) × 30 = 45 g, maintaining the same fixed ratio of 40% sulfur to 60% oxygen.

(a) The mass of oxygen combining with 9 g of carbon = (4/3) × 9 = 12 g, following the fixed C:O mass ratio of 3:4.
(b) By the Law of Conservation of Mass, the total mass of CO formed = 3 g + 4 g = 7 g.
(c) The Law of Constant Proportions is used, as carbon monoxide always contains C and O in the fixed mass ratio 3:4.
(d) Yes, it is consistent, because the ratio 6:8 simplifies to 3:4, which is the same fixed ratio required by the law.

(a) Sodium loses its single valence electron from the M-shell and becomes Na⁺, achieving a stable octet in its L-shell.
(b) Chlorine gains that electron to complete its M-shell octet, becoming the chloride ion Cl⁻.
(c) An ionic bond is formed, because the oppositely charged Na⁺ and Cl⁻ ions are held together by strong electrostatic attraction.
(d) The formula of the compound is NaCl, and its name is sodium chloride.

(a) By the criss-cross method, the charge 3 on Al³⁺ becomes the subscript of SO₄ and the charge 2 on SO₄²⁻ becomes the subscript of Al, giving Al₂(SO₄)₃.
(b) Brackets are used around SO₄ because it is a polyatomic ion, and the subscript 3 outside the bracket applies to the whole sulfate group.
(c) The formula unit contains 2 aluminium ions (Al³⁺) and 3 sulfate ions (SO₄²⁻).
(d) Formula unit mass = (27 × 2) + 3 × (32 + 16 × 4) = 54 + 3 × 96 = 54 + 288 = 342 u.

(a) Sodium chloride and copper sulfate are ionic, as they dissolve in water; camphor, sugar and naphthalene are covalent — though sugar is an exception that dissolves in water despite being covalent.
(b) NaCl and CuSO₄ solutions conduct electricity because they furnish free ions; sugar solution does not conduct.
(c) Sugar is a covalent compound that dissolves as neutral molecules without forming ions, so there are no charge carriers to carry current.
(d) Yes, ionic compounds conduct electricity when molten, because the ions become free to move and can carry charge.

(a) Element B is more reactive because it needs only one electron to complete its octet, whereas A needs three electrons to attain a stable configuration.
(b) Since both A and B are non-metals, they form covalent bonds by sharing electrons rather than transferring them.
(c) A shares three electrons, one with each of three B atoms, giving the formula AB₃ (similar to PCl₃).
(d) No, the compound would not conduct electricity when dissolved in water, because covalent compounds generally do not produce ions in solution.

(a) An ionic bond is present in the compound, as shown by its conduction in the dissolved and molten states.
(b) In the solid state, the ions are held in fixed positions in the crystal lattice and cannot move, so there are no mobile charge carriers.
(c) When dissolved in water or molten, the ions are freed from the lattice and can move freely, allowing them to carry electric charge.
(d) Sodium chloride (NaCl) is one such compound; copper sulfate (CuSO₄) is another example.

(a) The configuration 2,8,2 corresponds to atomic number 12, so metal M is magnesium (Mg), with electrons arranged as K=2, L=8, M=2.
(b) Magnesium loses two electrons to form Mg²⁺ and oxygen gains them to form O²⁻; the criss-cross method gives Mg₂O₂, which simplifies to MgO.
(c) An ionic bond is present, formed by the transfer of electrons from the metal (magnesium) to the non-metal (oxygen).
(d) Yes, since MgO is slightly soluble in water, the dissolved ions conduct electricity, though only weakly.

(a) The bulb glows in the sodium chloride solution, because NaCl dissociates into free Na⁺ and Cl⁻ ions that carry current.
(b) The bulb does not glow in the sugar solution because sugar dissolves as neutral molecules, producing no free ions, so no current flows.
(c) Solid NaCl does not conduct electricity, because its ions are fixed in the crystal lattice and cannot move.
(d) Sodium chloride is an ionic compound, while sugar is a covalent compound.

Introduction: The two fundamental laws of chemical combination are verified by simple experiments and explained by Dalton's theory.
Key Points:
1. Law of Conservation of Mass — Matter can neither be created nor destroyed in a chemical reaction (Lavoisier, 1789).
2. Verification — Vinegar + baking soda in a sealed balloon flask shows no mass change; open flask loses mass due to escaping CO₂.
3. Law of Constant Proportions — A compound always has its elements in a fixed mass ratio, e.g. water = H:O = 1:8 (Proust).
4. Dalton's support — Dalton's postulate that atoms are indivisible and combine in whole-number ratios explains both laws — atoms are merely rearranged, not created or destroyed.
Conclusion: Thus Dalton's atomic theory provides the logical explanation for both conservation of mass and constant proportions.

Introduction: Covalent bonding by electron sharing forms molecules of both elements and compounds.
Key Points:
1. H₂ — Two H atoms share one electron each to complete the duplet, forming a single bond H–H.
2. Cl₂ — Two Cl atoms share one electron each to complete the octet, forming a single bond Cl–Cl.
3. O₂ — Two O atoms share two electrons each, forming a double bond O=O.
4. HCl — H shares one electron with Cl; both achieve stability, forming a single bond H–Cl.
5. H₂O — O shares one electron with each of two H atoms, forming two single bonds and the molecule H–O–H (H₂O).
Conclusion: In every covalent molecule, shared electron pairs are attracted by both nuclei, making the molecule stable.

Introduction: Ionic bonding occurs when a metal transfers electrons to a non-metal, creating oppositely charged ions.
Key Points:
1. NaCl formation — Na (2,8,1) loses 1 electron → Na⁺; Cl (2,8,7) gains it → Cl⁻; electrostatic attraction forms NaCl.
2. MgCl₂ formation — Mg (2,8,2) loses 2 electrons → Mg²⁺; two Cl atoms each gain 1 electron → 2 Cl⁻; formula MgCl₂.
3. Cations — Positively charged ions formed by losing electrons (e.g. Na⁺, Mg²⁺).
4. Anions — Negatively charged ions formed by gaining electrons (e.g. Cl⁻, O²⁻).
5. Ions — Cations and anions are collectively called ions; the electrostatic attraction between them is the ionic bond.
Conclusion: Ionic compounds form when metals donate electrons to non-metals, producing ions held by electrostatic attraction.

Introduction: Covalent and ionic compounds follow different rules for naming and formula writing.
Key Points:
1. Covalent naming — Use prefixes (mono-, di-, tri-) for atom counts; second element ends in -ide, e.g. CO₂ = carbon dioxide.
2. Covalent formulae — Cross over valencies: H(1) + S(2) → H₂S; C(4) + Cl(1) → CCl₄.
3. Ionic naming — Cation first, then anion ending in -ide, e.g. NaCl = sodium chloride; CaO = calcium oxide.
4. Ionic formulae — Cross over charge numbers: Ca²⁺ + Cl⁻ → CaCl₂; Al³⁺ + O²⁻ → Al₂O₃.
5. Simplification — Divide subscripts by common factors: Mg₂O₂ → MgO; brackets used for multiple polyatomic ions: Al(OH)₃.
Conclusion: These systematic rules ensure every compound has a unique, unambiguous name and formula.

Introduction: Ionic and covalent compounds differ in physical properties due to their fundamentally different bonding.
Key Points:
1. Melting/boiling points — Ionic: high (strong inter-ionic attraction); Covalent: low (weaker intermolecular forces).
2. Solubility — Ionic: soluble in water; Covalent: generally insoluble in water but soluble in organic solvents.
3. Conductivity — Ionic: conduct when dissolved or molten (free ions); Covalent: generally do not conduct (no free ions).
4. Crystal structure — Ionic compounds form 3-D crystal lattices (e.g. NaCl: each Na⁺ surrounded by 6 Cl⁻); covalent compounds exist as discrete molecules.
5. Reason for differences — The strong electrostatic forces in ionic lattices versus the weaker interactions between neutral covalent molecules explain all the property differences.
Conclusion: The type of chemical bond — ionic or covalent — determines all the characteristic physical properties of a compound.

Introduction: The criss-cross method uses the charges on ions to systematically build ionic formulae.
Key Points:
1. Method — Write cation and anion symbols with charges below, then cross the charge numbers as subscripts.
2. CaCl₂ and Al₂O₃ — Ca²⁺/Cl⁻ → CaCl₂; Al³⁺/O²⁻ → Al₂O₃.
3. Mg(OH)₂ and CaCO₃ — Mg²⁺/OH⁻ → Mg(OH)₂ (brackets for 2 OH groups); Ca²⁺/CO₃²⁻ → CaCO₃ (subscripts 2 and 2 simplify to 1:1).
4. Al₂(SO₄)₃ and Fe(OH)₃ — Al³⁺/SO₄²⁻ → Al₂(SO₄)₃; Fe³⁺/OH⁻ → Fe(OH)₃.
5. Brackets rule — Brackets are used when two or more polyatomic ions of the same type appear in the formula.
Conclusion: The criss-cross method always produces a neutral formula where total positive and negative charges balance.

Introduction: Electronic configurations determine how elements bond and what properties the resulting compound has.
Key Points:
1. Element A — Configuration 2,8,1 — A tends to LOSE its 1 valence electron to complete the octet; it forms cation A⁺.
2. Element B — Configuration 2,6 — B tends to GAIN 2 electrons to complete its octet; it forms anion B²⁻.
3. Bond type — Since A is a metal (loses electrons) and B is a non-metal (gains electrons), an ionic bond forms.
4. Formula — Two A⁺ ions are needed for each B²⁻ ion → formula A₂B.
5. Properties — A₂B is an ionic compound: soluble in water, conducts electricity when dissolved or molten, and has a high melting point.
Conclusion: From the configurations alone, we can predict the bond type, formula, and physical properties of the compound.

Introduction: Molecular mass applies to covalent molecules; formula unit mass applies to ionic formula units.
Key Points:
1. Molecular masses — H₂O = 2(1)+16 = 18 u; CO₂ = 12+2(16) = 44 u; HNO₃ = 1+14+3(16) = 63 u; CH₄ = 12+4(1) = 16 u.
2. Formula unit mass — Na₂O — Na₂O = 2(23)+16 = 62 u.
3. Formula unit mass — Ca(NO₃)₂ — Ca(NO₃)₂ = 40 + 2×(14+3×16) = 40 + 2×(62) = 40 + 124 = 164 u.
4. Formula unit mass — Mg(OH)₂ and KCl — Mg(OH)₂ = 24 + 2×(16+1) = 24 + 34 = 58 u; KCl = 39 + 35.5 = 74.5 u.
5. Key distinction — Molecular mass is for covalent molecules; formula unit mass is for ionic compounds which do not form discrete molecules.
Conclusion: Both are calculated by adding the atomic masses of all atoms, but they apply to different types of compounds.

Introduction: The two laws of chemical combination provided the experimental basis for Dalton's atomic theory.
Key Points:
1. Conservation of Mass — Lavoisier showed total mass is unchanged in reactions — explained by atoms being indivisible and uncreatable.
2. Constant Proportions — Proust showed compounds have fixed element ratios — explained by atoms combining in simple whole-number ratios.
3. Key postulates — All matter is made of atoms; atoms are indivisible; atoms of an element are identical; atoms of different elements differ; atoms combine in whole-number ratios to form compounds.
4. Significance — Dalton's theory was the first scientific explanation of chemical behaviour and became the foundation of modern chemistry.
5. Limitation — Later discoveries (electrons, protons, neutrons, radioactivity) showed that atoms are divisible, refining the theory.
Conclusion: Dalton's theory logically unified the two laws and established the atomic view of matter that chemistry still builds on today.

Introduction: Each ionic compound forms when a metal transfers valence electrons to a non-metal, creating ions that achieve stable octets.
Key Points:
1. NaCl — Na (2,8,1) → Na⁺ (2,8); Cl (2,8,7) → Cl⁻ (2,8,8); one Na⁺ + one Cl⁻.
2. MgO — Mg (2,8,2) → Mg²⁺ (2,8); O (2,6) → O²⁻ (2,8); one Mg²⁺ + one O²⁻.
3. Al₂O₃ — Al (2,8,3) → Al³⁺ (2,8); O (2,6) → O²⁻ (2,8); two Al³⁺ + three O²⁻.
4. Na₂S and CaCl₂ — Na (2,8,1) → Na⁺; S (2,8,6) → S²⁻; two Na⁺ + one S²⁻ → Na₂S. Ca (2,8,8,2) → Ca²⁺; Cl (2,8,7) → Cl⁻; one Ca²⁺ + two Cl⁻ → CaCl₂.
5. Common principle — In each case, the metal loses and the non-metal gains just enough electrons for both to reach a stable octet.
Conclusion: The number of electrons transferred always ensures both ions achieve stable noble-gas-like electronic configurations.








NCERT Textbook (2025–26 edition) | Student: Shipan Das

Covers: Ch 10 Sound Waves: Characteristics and Applications | Ch 11 Reproduction: How Life Continues | Ch 12 Patterns in Life: Diversity and Classification | Ch 13 Earth as a System: Energy, Matter, and Life

Each chapter contains: Section A (55 MCQs, 1 mark) | Section B (15 Fill in the Blanks) | Section C (15 True/False) | Section D (2 Match the Column sets) | Section E (8 Assertion–Reason) | Section F (20 Two-Mark questions) | Section G (15 Three-Mark questions with structured answers) | Section H (10 Case-Based four-mark questions) | Section I (10 Five-Mark questions with structured answers)

Approximately half the questions are competency-based (application, analysis, numericals, case scenarios, diagram-based).