5.1 Basic Concept of Formation of Compounds
Atoms of elements combine chemically to achieve a stable duplet (2 valence electrons) or octet (8 valence electrons) electron arrangement, similar to noble gases in Group 18.
- Ionic Bonding: Involves the transfer of electrons between metal and non-metal atoms.
- Covalent Bonding: Involves the sharing of valence electrons between non-metal atoms.
- Other Chemical Bonds: Hydrogen bond, dative bond, and metallic bond.
5.2 Ionic Bond
An ionic bond is formed by electrostatic forces of attraction between oppositely charged ions created through electron transfer.
Formation of Ionic Bonds
- Metal atoms lose valence electrons to form positively charged ions (cations):
$$\text{M} \rightarrow \text{M}^{n+} + n e^-$$
- Non-metal atoms gain valence electrons to form negatively charged ions (anions):
$$\text{X} + n e^- \rightarrow \text{X}^{n-}$$
- Example - Sodium Chloride ($\text{NaCl}$):
- Sodium atom ($\text{Na}$, electron arrangement $2.8.1$) donates $1$ valence electron to achieve a stable octet arrangement ($2.8$), forming a $\text{Na}^+$ ion.
- Chlorine atom ($\text{Cl}$, electron arrangement $2.8.7$) receives $1$ electron to achieve a stable octet arrangement ($2.8.8$), forming a $\text{Cl}^-$ ion.
- Strong electrostatic forces of attraction hold the $\text{Na}^+$ and $\text{Cl}^-$ ions together in a giant ionic lattice structure.
5.3 Covalent Bond
A covalent bond is formed when non-metal atoms share one or more pairs of valence electrons to achieve stable duplet or octet electron arrangements.
Types of Covalent Bonds
- Single Bond ($-$): Shares $1$ pair of electrons (e.g., $\text{H}_2$, $\text{HCl}$, $\text{Cl}_2$, $\text{H}_2\text{O}$, $\text{CH}_4$).
- Double Bond ($=$): Shares $2$ pairs of electrons (e.g., $\text{O}_2$, $\text{CO}_2$).
- Triple Bond ($\equiv$): Shares $3$ pairs of electrons (e.g., $\text{N}_2$).
5.4 Hydrogen Bond
A hydrogen bond is an attraction force between a hydrogen atom bonded to a highly electronegative atom (such as Nitrogen, Oxygen, or Fluorine) and a lone pair of electrons on another electronegative atom ($\text{N}$, $\text{O}$, or $\text{F}$).
Effects of Hydrogen Bonding
- Higher Boiling Points: Molecules with hydrogen bonds (e.g., $\text{H}_2\text{O}$, $\text{NH}_3$, $\text{HF}$) have significantly higher boiling points than expected because additional thermal energy is needed to overcome the strong hydrogen bonds alongside weak van der Waals forces.
- Solubility in Water: Compounds like ethanol ($\text{C}_2\text{H}_5\text{OH}$) dissolve in water because they can form hydrogen bonds with water molecules.
- Moistened Hair / Paper Behavior: Water molecules form hydrogen bonds with protein chains (keratin in hair or cellulose in paper), causing fibers to stick together when wet.
5.5 Dative Bond (Coordinate Bond)
A dative bond is a special type of covalent bond where the shared pair of electrons originates from only one of the bonding atoms.
Examples of Dative Bond Formation
- Ammonium Ion ($\text{NH}_4^+$): Formed when a hydrogen ion ($\text{H}^+$) with an empty shell bonds to the nitrogen atom in an ammonia ($\text{NH}_3$) molecule using the nitrogen's lone pair of electrons.
$$\text{NH}_3 + \text{H}^+ \rightarrow \text{NH}_4^+$$
- Hydroxonium Ion ($\text{H}_3\text{O}^+$): Formed when a hydrogen ion ($\text{H}^+$) bonds with a water molecule ($\text{H}_2\text{O}$) via one of Oxygen's lone pairs.
$$\text{H}_2\text{O} + \text{H}^+ \rightarrow \text{H}_3\text{O}^+$$
5.6 Metallic Bond
A metallic bond is the electrostatic attraction force between a sea of delocalized electrons and positively charged metal ions.
- Delocalized Electrons: Valence electrons in metal atoms are free to move throughout the entire crystal lattice structure rather than remaining tied to a specific atom.
- Electrical & Thermal Conductivity: Metals conduct electricity and heat efficiently because the delocalized electrons are mobile and act as charge/energy carriers.
5.7 Properties of Ionic and Covalent Compounds
Comparison Table
| Property |
Ionic Compounds |
Simple Covalent Compounds |
| Melting & Boiling Points |
High (strong ionic bonds require large amounts of heat to break). |
Low (weak intermolecular van der Waals forces require little heat to overcome). |
| Electrical Conductivity |
Conducts electricity in molten or aqueous state (free-moving ions present); does not conduct in solid state. |
Does not conduct electricity in any state (exists as neutral molecules, no free-moving ions or electrons). |
| Solubility |
Usually soluble in water; insoluble in organic solvents. |
Usually soluble in organic solvents; insoluble in water (except those forming hydrogen bonds). |
Giant Covalent Structures
Structures such as diamond and silicon dioxide ($\text{SiO}_2$) have extremely high melting points and hardness because all atoms are held together by continuous, strong covalent bonds throughout a 3D network.