1. The Law of Conservation of Mass

In a chemical reaction, matter can neither be created nor destroyed. This is known as the Law of Conservation of Mass, proposed by Antoine Lavoisier in 1789.

1.1 Antoine Lavoisier

Known as the Father of Modern Chemistry, he proposed that "...in every operation an equal quantity of matter exists both before and after the operation."

2. Law of Constant Proportions

Proposed by Joseph Proust, this law states that in a chemical compound, the elements are always present in definite proportions by mass. This is also known as the Law of Definite Proportions.

For example, if water is collected from various sources (rivers, borewells, oceans) and purified, it will always contain hydrogen and oxygen in a mass ratio of 1:8. This means if 9 g of water is decomposed, 1 g of hydrogen and 8 g of oxygen are always obtained.

3. Dalton’s Atomic Theory

John Dalton combined these early experimental observations to propose an atomic theory that logically explained the laws of chemical combination.

Postulates of Dalton's Atomic Theory:

1. All matter is made up of very tiny particles called atoms, which participate in chemical reactions.
2. Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction.
3. Atoms of a given element are identical in mass and chemical properties.
4. Atoms of different elements have different masses and chemical properties.
5. Atoms combine in the ratio of simple whole numbers to form compounds.
6. The relative number and kinds of atoms are constant in a given compound.

4. How Atoms Combine (Chemical Bonds)

Atoms of an element can combine to form molecules of that element (e.g., $H_2$), or they can combine with atoms of different elements to form molecules of a compound (e.g., $HCl$).

The force that holds atoms together to make the resulting arrangement more stable is called a chemical bond. Bonding generally takes place in two ways:

4.1 Covalent Bond (Sharing of Electrons)

A covalent bond is formed when atoms share a few or all of their valence electrons to complete their octet (or duplet).

• Single Covalent Bond: E.g., Hydrogen molecule ($H_2$). Two hydrogen atoms share one electron each, forming a single bond ($H-H$). Chlorine molecule ($Cl_2$) also forms a single bond ($Cl-Cl$).
• Double Covalent Bond: E.g., Oxygen molecule ($O_2$). Two oxygen atoms share two pairs of electrons, forming a double bond ($O=O$).

4.2 Ionic Bond (Transfer of Electrons)

If the valence shell of an atom has less than four electrons (typically metals), it generally donates them. Atoms with more than four valence electrons (typically non-metals) tend to gain them.

Formation of Sodium Chloride ($NaCl$):

• Sodium ($Na$) loses one electron to form a positively charged cation ($Na^+$).
• Chlorine ($Cl$) gains one electron to form a negatively charged anion ($Cl^-$).
• These oppositely charged ions are held together by a strong electrostatic force called an ionic bond.

5. Naming Ionic Compounds and Writing Formulae

5.1 Naming Ionic Compounds

The cation (metal) is written first, followed by the anion (non-metal). Simple anions end with -ide (e.g., chloride, oxide, sulfide). Groups of atoms acting as a single charged entity are called polyatomic ions (e.g., nitrate, sulfate, carbonate).

Table 9.1: Common Monoatomic and Polyatomic Ions

Name of ion	Formula	Valency
Sodium	$Na^+$	1
Calcium	$Ca^{2+}$	2
Aluminium	$Al^{3+}$	3
Chloride	$Cl^-$	1
Oxide	$O^{2-}$	2
Nitrate	$NO_3^-$	1
Carbonate	$CO_3^{2-}$	2
Sulfate	$SO_4^{2-}$	2
Hydroxide	$OH^-$	1
Ammonium	$NH_4^+$	1

5.2 The Criss-Cross Method

To write the chemical formula of a compound:

1. Write the symbols of the constituent elements side-by-side (cation first).
2. Write their valencies/charges directly below them.
3. Criss-cross the valencies to become the subscripts of the other element.
4. Simplify the ratio if possible (e.g., $Ca_2O_2$ becomes $CaO$).
5. Use brackets if there is more than one polyatomic ion (e.g., $Mg(OH)_2$).

6. Properties of Ionic and Covalent Compounds

• Solubility: Ionic compounds (like $NaCl$) are generally soluble in water but insoluble in organic solvents (like kerosene). Covalent compounds (like camphor, naphthalene) are often insoluble in water but dissolve in organic solvents.
• Electrical Conductivity: Ionic compounds conduct electricity when dissolved in water or in a molten state because their ions are free to move. They do not conduct in the solid state. Covalent compounds generally do not conduct electricity.
• Melting and Boiling Points: Ionic compounds have high melting/boiling points due to strong inter-ionic attractions. Covalent compounds generally have low melting/boiling points.

7. Molecular Mass and Formula Unit Mass

7.1 Molecular Mass (for Covalent Compounds)

The molecular mass is the sum of the atomic masses of all the atoms in a molecule.

Example: Water ($H_2O$)
Atomic mass of $H = 1\text{ u}$, $O = 16\text{ u}$.
Molecular mass = $(1\text{ u} \times 2) + (16\text{ u} \times 1) = 18\text{ u}$.

7.2 Formula Unit Mass (for Ionic Compounds)

Since ionic compounds form 3D crystal lattices instead of discrete molecules, we calculate the mass of their simplest ratio, known as the formula unit.

Example: Calcium nitrate, $Ca(NO_3)_2$
Atomic mass of $Ca = 40\text{ u}$, $N = 14\text{ u}$, $O = 16\text{ u}$.
Formula unit mass = $(40\text{ u} \times 1) + [ (14\text{ u} \times 1) + (16\text{ u} \times 3) ] \times 2$
$= 40\text{ u} + [14\text{ u} + 48\text{ u}] \times 2$
$= 40\text{ u} + [62\text{ u}] \times 2$
$= 40\text{ u} + 124\text{ u} = 164\text{ u}$.