Chemistry Chapter 3: Atoms and Molecules

Laws of Chemical Combination

Law of Conservation of Mass

This law states that in a physical or chemical reaction, the total mass of the products is equal to the total mass of the reactants. Matter can neither be created nor destroyed.

Total mass of reactants = Total mass of products

[Diagram showing balanced chemical equation with mass conservation]

Example: Decomposition of Calcium Carbonate

CaCO₃ ⟶ CaO + CO₂

100g ⟶ 56g + 44g

Law of Definite Proportions

A chemical compound always contains the same elements combined together in the same proportion by mass.

Compound	Elements	Ratio by Mass
Water (H₂O)	Hydrogen and Oxygen	1:8
Ammonia (NH₃)	Nitrogen and Hydrogen	14:3
Carbon dioxide (CO₂)	Carbon and Oxygen	3:8

Dalton's Atomic Theory

The theoretical proof for the validity of different laws of chemical combination was given by John Dalton.

Important Postulates:

All matter is made up of very tiny particles called atoms.
Atoms cannot be divided (atoms are indivisible).
Atoms can combine to form compounds in fixed ratios.
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 small whole numbers to form compounds.
The relative number and kinds of atoms are constant in a given compound.

Memory Tip:

"All Atoms Are Identical, Different Elements Differ" - helps remember key postulates of Dalton's theory.

Atoms and Molecules

Atom

An atom is the smallest unit of any substance that cannot be divided and is made up of electrons, protons, and neutrons. Atoms are extremely small, with atomic radius measured in nanometers (1 nm = 10^-9 m).

[Diagram of atomic structure showing nucleus and electron shells]

Molecule

A molecule is a group of two or more atoms that are chemically bonded together. It is the smallest particle of an element or compound that is capable of independent existence and retains all the properties of the substance.

Quick Fact:

Molecules can be made of same type of atoms (elements) or different types (compounds).

Atomicity

Atomicity is defined as the number of atoms present in one molecule of an element.

Type	Description	Examples
Monoatomic	Single atom molecules	Noble gases (He, Ne, Ar)
Diatomic	Two atom molecules	O₂, N₂, Cl₂
Triatomic	Three atom molecules	O₃, CO₂
Polyatomic	Many atom molecules	S₈, P₄, SO₄^2-

Ions and Valency

Ions

An ion is a charged particle formed by loss or gain of electrons.

Cation: Positively charged particle formed by loss of electron(s). Example: Na⁺, Mg²⁺
Anion: Negatively charged particle formed by gain of electron(s). Example: Cl^-, O^2-

[Diagram showing formation of cations and anions]

Polyatomic Ions

A group of atoms carrying a charge is called a polyatomic ion.

Ion	Name
NH₄⁺	Ammonium ion
SO₄^2-	Sulfate ion
NO₃^-	Nitrate ion
CO₃^2-	Carbonate ion

Valency

Valency is defined as the combining power of an element with another atom to form chemical bonds. It is equal to the number of electrons an atom can gain, lose, or share to achieve the nearest noble gas configuration.

Valency Memory Tip:

"Happy Henry Likes Beer But Can't Obtain Food" - helps remember common elements with valency 1 (H, He, Li, Be, B, C, O, F).

Element	Valency
Sodium (Na)	1
Magnesium (Mg)	2
Aluminum (Al)	3
Oxygen (O)	2

Writing Chemical Formulas

Rules for Writing Formulas

Rule 1: Write the symbols of elements with their valencies and cross them to get the formula.

Al³⁺ O^2- → Al₂O₃

Rule 2: When the valency is 1, it is not written.

Na⁺ OH^- → NaOH

Rule 3: When valencies are numerically equal, they are not written.

Ca²⁺ O^2- → CaO

Rule 4: For multiple polyatomic ions, use parentheses.

Al³⁺ SO₄^2- → Al₂(SO₄)₃

Rule 5: Always simplify the formula to lowest terms (except for covalent compounds).

Sn⁴⁺ O^2- → SnO₂ (not Sn₂O₄)

Examples

Compound	Formula
Calcium Sulphate	CaSO₄
Iron Sulphate	FeSO₄
Ammonium Nitrate	NH₄NO₃
Silver Iodide	AgI

Mole Concept

Definition

The mole is the amount of substance that contains as many elementary entities as there are atoms in exactly 12 grams of carbon-12 (C¹²).

Avogadro's Number:

N_A = 6.022 × 10²³ particles/mol (Remember as "6.022 part of Avogadro's party!")

1 mole = 6.022 × 10²³ particles (atoms, molecules, ions)
At STP (Standard Temperature and Pressure):
- 1 mole of gas = 22.4 liters
- 1 mole of gas = 22400 mL

Molar Mass

The mass of 1 mole of any substance is known as its molar mass (in g/mol).

Substance	Molar Mass (g/mol)
Nitrogen (N₂)	28
Water (H₂O)	18
Carbon dioxide (CO₂)	44

Conversions

Number of moles (n) = Mass (W) / Molar mass (M)

Problem	Solution
Convert 35g of aluminum to moles	n = 35g / 27g/mol = 1.3 moles
Convert 56g of calcium to moles	n = 56g / 40g/mol = 1.4 moles
Find mass of 0.8 moles of SiO₂	W = n × M = 0.8 × 60 = 48g
Find moles in 55.5L of CO₂ at STP	n = Volume / 22.4 = 55.5 / 22.4 ≈ 2.48 moles