CBSE Class 9 Science Chapter 3: Atoms And Molecules Notes

Key Points to Remember

• Law of Conservation of Mass states that mass can neither be created nor destroyed in a chemical reaction.
• Law of Constant Proportions states that a pure chemical compound always contains the same elements combined in a fixed proportion by mass.
• Dalton's Atomic Theory proposed that matter consists of indivisible atoms, atoms of the same element are identical, and atoms combine in simple whole-number ratios.
• An atom is the smallest particle of an element that may or may not exist independently but takes part in a chemical reaction.
• A molecule is the smallest particle of an element or a compound that can exist independently and shows all the properties of that substance.
• Ions are charged particles formed by the gain or loss of electrons. Cations are positively charged, Anions are negatively charged.
• Valency is the combining capacity of an element, determining how many other atoms it can bond with.
• The atomic mass unit (amu), now 'u' (unified mass), is defined as exactly 1/12th the mass of one atom of Carbon-12.
• Molar Mass is the mass of one mole of a substance, numerically equal to its atomic or molecular mass in grams.
• One mole of any substance contains Avogadro's number ($6.022 \times 10^{23}$) of particles (atoms, molecules, or ions).

Key Definitions

Atom

The smallest indivisible particle of an element that retains its chemical identity during a chemical reaction.

Molecule

A group of two or more atoms chemically bonded together, existing as an independent unit.

Ion

An atom or group of atoms that has gained or lost one or more electrons, acquiring an electrical charge (positive for cation, negative for anion).

Valency

The combining power of an element or a radical, represented by the number of hydrogen atoms with which it can combine or replace.

Atomic Mass Unit (amu or u)

A standard unit for expressing atomic and molecular masses. One 'u' is defined as exactly one-twelfth of the mass of an unbound atom of carbon-12.

Molecular Mass

The sum of the atomic masses of all the atoms in a molecule, expressed in unified atomic mass units (u).

Formula Unit Mass

The sum of the atomic masses of all atoms in a formula unit of an ionic compound, expressed in unified atomic mass units (u).

Molar Mass

The mass of one mole of a substance, expressed in grams per mole (g/mol). Numerically equal to atomic or molecular mass.

Mole

The amount of substance that contains as many elementary entities (atoms, molecules, ions, etc.) as there are atoms in exactly 12 grams of carbon-12. This number is Avogadro's constant.

Avogadro's Number ($N_A$)

The number of particles (atoms, molecules, ions) in one mole of any substance, approximately $6.022 \times 10^{23}$.

Laws of Chemical Combination and Dalton's Atomic Theory

Understanding how elements combine to form compounds is fundamental to chemistry. Two crucial laws govern these combinations, which later formed the basis for Dalton's Atomic Theory.

Law of Conservation of Mass

Proposed by Antoine Lavoisier, this law states that mass can neither be created nor destroyed in a chemical reaction. This means that during any chemical change, the total mass of the reactants must be equal to the total mass of the products. For example, if 10g of hydrogen reacts completely with 80g of oxygen, it will form exactly 90g of water. There is no loss or gain of mass in an isolated system. This law highlights that atoms merely rearrange during a chemical reaction, they are not destroyed or created.

Law of Constant Proportions (or Law of Definite Proportions)

Formulated by Joseph Proust, this law states that in a pure chemical substance, the elements are always present in definite proportions by mass. Regardless of the source or method of preparation, a given compound always contains the same elements in the same fixed mass ratio. For instance, water ($H_2O$) always consists of hydrogen and oxygen in a mass ratio of 1:8, irrespective of whether it's river water, tap water, or laboratory-prepared water. This consistency in composition is a defining characteristic of pure chemical compounds.

Dalton's Atomic Theory

John Dalton, building upon these laws, proposed his atomic theory in 1808. His postulates laid the groundwork for modern chemistry:

1. Matter consists of indivisible atoms. (Though later proven incorrect as atoms are composed of subatomic particles, this was a revolutionary idea at the time).
2. Atoms of a given element are identical in all respects, i.e., they have the same mass, size, and chemical properties. (Isotopes later showed variations in mass within an element).
3. Atoms of different elements have different masses, sizes, and chemical properties.
4. Atoms combine in the ratio of small whole numbers to form compounds. This explains the Law of Constant Proportions.
5. Atoms can neither be created nor destroyed in a chemical reaction. This explains the Law of Conservation of Mass. Chemical reactions involve the rearrangement of atoms.

While some postulates were later refined with advancements in atomic structure, Dalton's theory was pivotal in explaining chemical reactions and the nature of matter, paving the way for quantitative chemistry.

Steps to Write Chemical Formulae

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Calculating Molecular Mass and Molar Mass

• {"title":"Example 1: Molecular Mass of Water ($H_2O$)","description":"Given atomic masses: H = 1 u, O = 16 u.\n\nMolecular mass of $H_2O$ = (2 \\times Atomic mass of H) + (1 \\times Atomic mass of O)\n= (2 \\times 1 u) + (1 \\times 16 u)\n= 2 u + 16 u\n= 18 u\n\nMolar Mass of $H_2O$ = 18 g/mol"}
• {"title":"Example 2: Molecular Mass of Sulphuric Acid ($H_2SO_4$)","description":"Given atomic masses: H = 1 u, S = 32 u, O = 16 u.\n\nMolecular mass of $H_2SO_4$ = (2 \\times Atomic mass of H) + (1 \\times Atomic mass of S) + (4 \\times Atomic mass of O)\n= (2 \\times 1 u) + (1 \\times 32 u) + (4 \\times 16 u)\n= 2 u + 32 u + 64 u\n= 98 u\n\nMolar Mass of $H_2SO_4$ = 98 g/mol"}
• {"title":"Example 3: Formula Unit Mass of Sodium Chloride (NaCl)","description":"Given atomic masses: Na = 23 u, Cl = 35.5 u.\n\nFormula unit mass of NaCl = (1 \\times Atomic mass of Na) + (1 \\times Atomic mass of Cl)\n= (1 \\times 23 u) + (1 \\times 35.5 u)\n= 23 u + 35.5 u\n= 58.5 u\n\nMolar Mass of NaCl = 58.5 g/mol"}

Exam Tips for Atoms and Molecules

1. Memorize common element symbols and valencies: This is crucial for quickly writing chemical formulae. Make flashcards using YoLearn.ai's tools for elements like Na, K, Al, Ca, Mg, Zn, Fe, Cu, Ag, Pb, O, S, N, Cl, Br, I, C, H. Also, learn common polyatomic ions like nitrate ($NO_3^-$), sulphate ($SO_4^{2-}$), carbonate ($CO_3^{2-}$), ammonium ($NH_4^+$).
2. Distinguish between Molecular Mass and Molar Mass: Molecular mass is in 'u' (unified mass unit), while Molar Mass is in 'g/mol'. They are numerically identical, but the units are different and important for conceptual clarity.
3. Master the Mole Concept Formulae: Know the relationships:

• Number of moles = Given mass / Molar mass
• Number of moles = Given number of particles / Avogadro's Number
• Mass = Number of moles \times Molar mass
• Number of particles = Number of moles \times Avogadro's Number

1. Practice chemical formula writing: This is a frequently tested skill. Understand the criss-cross method thoroughly.
2. Read questions carefully: Pay attention to whether the question asks for mass in 'u', 'grams', or 'number of moles/atoms/molecules'.
3. Don't confuse atoms, molecules, and ions: An atom is a single unit, a molecule is a group of atoms covalently bonded, and an ion is a charged atom or molecule.

Practice Questions with Solutions

• Q: State the Law of Constant Proportions. A: The Law of Constant Proportions states that a pure chemical compound always contains the same elements combined together in a fixed proportion by mass, irrespective of its source or method of preparation.
• Q: What is the valency of the sulphate ion ($SO_4$)? A: The valency of the sulphate ion ($SO_4^{2-}$) is 2.
• Q: Calculate the molecular mass of ammonia ($NH_3$). (Atomic masses: N=14 u, H=1 u) A: Molecular mass of $NH_3$ = (1 x 14 u) + (3 x 1 u) = 14 u + 3 u = 17 u.
• Q: How many atoms are present in one molecule of phosphorus ($P_4$)? A: One molecule of phosphorus ($P_4$) contains 4 atoms of phosphorus.

Frequently Asked Questions