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

9.1 Law of Conservation of Mass

Law of Conservation of Mass: Formulated by Antoine Lavoisier in 1789, it states that matter can neither be created nor destroyed in a chemical reaction.
Physical Changes: Demonstrated by dissolving salt in water, where the final solution mass is exactly equal to the sum of the individual masses of salt and water.
Chemical Changes: Verified via reactions like vinegar reacting with baking soda in a closed system to prevent gaseous carbon dioxide (CO₂) from escaping, showing total reactant mass equals total product mass.

9.2 Law of Constant Proportions

Law of Constant Proportions: Proposed by Joseph Proust, stating that in any chemical compound, elements always combine in a fixed, definite ratio by mass, regardless of source or method of preparation.
Water Composition: Purified water (H₂O) from any source consistently contains hydrogen and oxygen in a fixed mass ratio of 1:8.
Cinnabar (Hingula): Heating cinnabar always yields mercury and sulfur in a fixed mass ratio of approximately 86.22% to 13.78%.

9.3 Dalton's Atomic Theory

John Dalton: Introduced a turning point in chemistry in 1808 by proposing postulates that explain chemical combinations and mass conservation.
Atom Indivisibility: All matter is composed of extremely tiny, indivisible particles called atoms, which cannot be created or destroyed.
Elemental Identity: Atoms of a given element are identical in mass and chemical properties, whereas atoms of different elements differ in mass and chemical properties.
Compound Formation: Atoms combine in ratios of simple whole numbers to form stable compounds.

9.4 How Atoms Combine

Covalent Bond: A chemical bond formed through the mutual sharing of valence electron pairs between atoms to achieve a stable octet or duplet.
Single Bonds: Formed by sharing one pair of electrons, such as in hydrogen (H₂) and hydrogen chloride (HCl).
Double and Triple Bonds: Formed by sharing multiple pairs, exemplified by double-bonded oxygen (O₂) and triple-bonded nitrogen (N₂).
Covalent Nomenclature: Named using Greek prefixes (mono-, di-, tri-, tetra-) to specify atom counts, ending the second element with "-ide" (e.g., carbon disulfide (CS₂)).

Ionic Bond: The electrostatic force of attraction that holds oppositely charged ions together, formed when metals transfer electrons to non-metals.
Cation: A positively charged ion formed when an atom loses valence electrons (e.g., sodium cation (Na⁺)).
Anion: A negatively charged ion formed when an atom gains electrons (e.g., chloride anion (Cl^-)).
Crystal Lattice: Ionic compounds organize into a repeating 3-D pattern of oppositely charged ions instead of remaining as discrete single units.
Polyatomic Ions: Multi-atom groups carrying a net charge, such as hydroxide (OH^-), sulfate (SO₄₂-), and ammonium (NH₄+).

9.5 Writing Chemical Formulae

Covalent Formulae: Written by listing symbols of elements, stating their valencies, and crossing them over to write as subscripts (e.g., carbon tetrachloride (CCl₄)).
Ionic Formulae: The cation symbol is written first, followed by the anion. Charges are crossed over as subscripts and simplified to the lowest whole-number ratio.
Brackets in Formulae: Required when a formula contains two or more identical polyatomic ions, such as in aluminium hydroxide (Al(OH)₃).

9.6 Properties of Compounds

Ionic Compounds: Highly soluble in water, insoluble in organic solvents (like petrol or kerosene), and conduct electricity in aqueous or molten states due to mobile ions.
Covalent Compounds: Generally insoluble in water, soluble in organic solvents, and non-conductive of electricity due to the absence of free ions.
Thermal Stability: Ionic compounds have high melting and boiling points due to strong electrostatic attraction, while covalent compounds have lower transition points.

9.7 & 9.8 Quantitative Chemical Mass

Molecular Mass: Calculated for covalent compounds by summing the atomic masses of all atoms present in a single molecule (e.g., carbon dioxide (CO₂) has a mass of 44 u).
Formula Unit Mass: Used for ionic compounds because they exist in crystalline lattice networks rather than distinct molecules; calculated using the simplest formula unit ratio (e.g., sodium oxide (Na₂O) has a mass of 62 u).