Chemical Bonding

Stability: Atoms combine to achieve stability by acquiring the electronic configuration of the nearest noble gas (Group 18).

Noble Gases: They are stable because their valence shells are complete (He has 2, others have 8 electrons).

Rules:

• Duplet Rule: Attaining 2 electrons in the valence shell (e.g., H, He).
• Octet Rule: Attaining 8 electrons in the valence shell.

Definition: The force of attraction that holds atoms together in a molecule or crystal.

Energetics:

• Attraction: Forces between nucleus of one atom and electrons of other. Lowers energy.
• Repulsion: Forces between nuclei or electrons of both atoms. Increases energy.
• Bond Formation: Occurs when attractive forces dominate repulsive forces, and potential energy is minimum.

Definition: A bond formed by the complete transfer of one or more electrons from one atom (metal) to another atom (non-metal).

Example - Sodium Chloride ($NaCl$):

• $Na (1s^2 2s^2 2p^6 3s^1) \to Na^+ (1s^2 2s^2 2p^6) + 1e^-$ (Loss of e, Octet complete)
• $Cl (3s^2 3p^5) + 1e^- \to Cl^- (3s^2 3p^6)$ (Gain of e, Octet complete)
• $Na^+ + Cl^- \to NaCl$ (Electrostatic attraction)

Definition: A bond formed by the mutual sharing of electrons between atoms (typically non-metals).

Types:

• Single Covalent Bond ($-$): One electron pair shared (e.g., $H-H$, $Cl-Cl$, $CH_4$).
• Double Covalent Bond ($=$): Two electron pairs shared (e.g., $O=O$, $C_2H_4$).
• Triple Covalent Bond ($\equiv$): Three electron pairs shared (e.g., $N\equiv N$, $C_2H_2$).

Definition: A type of covalent bond where the shared pair of electrons is donated by one atom only.

Components:

• Donor: Atom that provides electron pair (e.g., N in $NH_3$).
• Acceptor: Atom that accepts electron pair (e.g., $H^+$ or B in $BF_3$).

Arrow Representation: An arrow ($\to$) points from donor to acceptor.

Examples:

• $NH_3 + H^+ \to NH_4^+$ (Ammonium ion)
• $NH_3 + BF_3 \to H_3N \to BF_3$ (Adduct)

Non-Polar: Bond formed between two similar atoms (same difference in electronegativity = 0). Electrons are shared equally. E.g., $H_2$, $Cl_2$, $O_2$.

Polar: Bond formed between two different atoms (difference in electronegativity > 0). Electrons are attracted more towards the more electronegative atom (creating partial charges $\delta+$ and $\delta-$). E.g., $HCl$ ($H^{\delta+} - Cl^{\delta-}$), $H_2O$.

Definition: A bond formed by the attraction between positive metal ions and mobile (delocalized) electrons.

Theory: Electron Sea Model or Electron Gas Model.

Factors Affecting Strength:

• Number of valence electrons (more e = stronger bond).
• Size of metal ion (smaller size = stronger bond).
• Example: Mg has stronger metallic bond than Na (Mg gives 2e, Na gives 1e).

Definition: Weak forces of attraction present between molecules.

Types:

• Dipole-Dipole Forces: Attraction between positive end of one polar molecule and negative end of another (e.g., $HCl$).
• Hydrogen Bonding: Strong permanent dipole-dipole attraction between covalently bonded H atom and highly electronegative atom (N, O, F). E.g., Water ($H_2O$), HF, $NH_3$. Affects physical properties like boiling point.

Ionic Compounds: Crystalline solids, high MP/BP, conduct electricity in molten/solution form (due to free ions), soluble in polar solvents (Water).

Covalent Compounds: Gases/Liquids/Soft Solids, low MP/BP, bad conductors (no free ions), soluble in non-polar solvents (Benzene). Exceptions: Diamond (hard, high MP).

Metals: Malleable (sheets), Ductile (wires), Good conductors (free electrons), Metallic Luster, high MP/BP.