Chemical Bonding

Introduction to Chemical Bonding

• Chemical Bond: The force that holds atoms or ions together in a molecule or compound.
• Atoms form bonds to achieve stability by attaining a full valence shell (octet rule).

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Types of Chemical Bonds

1. Ionic Bond (Electrovalent Bond):

   • Definition: Formed by the transfer of electrons from one atom to another, resulting in the formation of ions.
   • Key Features:
     • One atom loses electrons to become a cation, while another gains electrons to become an anion.
     • The electrostatic attraction between oppositely charged ions forms the bond.
   • Example:
     • Sodium chloride ($\text{NaCl}$): $\text{Na} \rightarrow \text{Na}^+ + e^-$ $\text{Cl} + e^- \rightarrow \text{Cl}^-$ $\text{Na}^+ + \text{Cl}^- \rightarrow \text{NaCl}$
   • Properties:
     • High melting and boiling points.
     • Conduct electricity in molten or aqueous states.
     • Soluble in polar solvents like water.

2. Covalent Bond:

   • Definition: Formed by the sharing of electrons between two atoms.
   • Key Features:
     • Shared electrons contribute to each atom's valence shell.
     • Common in non-metallic elements.
   • Example:
     • Water ($\text{H}_2\text{O}$): $\text{H} \cdot + \text{O} \cdot \cdot \rightarrow \text{H-O-H}$
   • Properties:
     • Low melting and boiling points.
     • Do not conduct electricity.
     • Insoluble in polar solvents but soluble in non-polar solvents.

3. Metallic Bond:

   • Definition: Formed by the attraction between a lattice of positive metal ions and a sea of delocalized electrons.
   • Key Features:
     • Electrons move freely within the metal lattice, providing conductivity.
   • Example:
     • Metals like copper, gold, and aluminum.
   • Properties:
     • Good electrical and thermal conductors.
     • Malleable and ductile.

4. Coordinate (Dative) Bond:

   • Definition: A type of covalent bond where both shared electrons come from the same atom.
   • Key Features:
     • Formed between a donor atom (provides the lone pair) and an acceptor atom.
   • Example:
     • Formation of ammonium ion ($\text{NH}_4^+$): $\text{NH}_3 + \text{H}^+ \rightarrow \text{[NH}_4\text{]}^+$
   • Properties:
     • Exhibits characteristics of covalent bonds.

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Bond Parameters

1. Bond Length:

   • The distance between the nuclei of two bonded atoms.
   • Shorter bond length indicates a stronger bond.
   • Example: $\text{C-C}$ (154 pm), $\text{C=C}$ (134 pm).

2. Bond Angle:

   • The angle between two bonds originating from the same atom.
   • Example: Water ($\text{H}_2\text{O}$): $104.5^\circ$.

3. Bond Energy:

   • The energy required to break one mole of bonds in a substance.
   • Higher bond energy indicates a stronger bond.
   • Example: $\text{H-H}$ (436 kJ/mol), $\text{C-H}$ (412 kJ/mol).

4. Bond Order:

   • The number of shared electron pairs between two atoms.
   • Example:
     • Single bond ($\text{C-C}$): Bond order = 1.
     • Double bond ($\text{C=C}$): Bond order = 2.

Theories of Chemical Bonding

1. Valence Bond Theory (VBT):
   • Overview: Explains the formation of covalent bonds based on the overlapping of atomic orbitals.
   • Key Features:
     1. A covalent bond forms when two half-filled atomic orbitals overlap.
     2. The strength of the bond depends on the extent of overlap.
   • Types of Overlapping:
     • Sigma ($\sigma$) Bond:
       • Formed by the head-on overlap of orbitals.
       • Stronger than a pi bond.
       • Example: $\text{H}_2$ molecule ($1s$ orbital overlap).
     • Pi ($\pi$) Bond:
       • Formed by the lateral overlap of orbitals.
       • Weaker than a sigma bond.
       • Example: $\text{C=C}$ bond in ethene.
   • Limitations:
     • Fails to explain the formation of multiple bonds.
     • Cannot account for the shapes of molecules.

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2. Molecular Orbital Theory (MOT):
   • Overview: Explains bonding by combining atomic orbitals to form molecular orbitals.
   • Key Features:
     1. Atomic orbitals combine to form bonding and antibonding molecular orbitals.
     2. Electrons are filled in molecular orbitals according to the Aufbau principle.
   • Bond Order: $\text{Bond Order} = \frac{\text{Number of electrons in bonding MOs} - \text{Number of electrons in antibonding MOs}}{2}$
   • Example:
     • For $\text{O}_2$:
       • Bonding MOs: 8 electrons.
       • Antibonding MOs: 4 electrons. $\text{Bond Order} = \frac{8 - 4}{2} = 2$
   • Magnetic Properties:
     • If unpaired electrons are present, the molecule is paramagnetic.
     • If all electrons are paired, the molecule is diamagnetic.
     • Example: $\text{O}_2$ is paramagnetic due to two unpaired electrons.

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3. Hybridization:
   • Definition: Mixing of atomic orbitals to form new hybrid orbitals of equivalent energy.
   • Types of Hybridization:
     1. sp Hybridization:
        • Linear geometry, bond angle: $180^\circ$.
        • Example: $\text{BeCl}_2$.
     2. sp^2 Hybridization:
        • Trigonal planar geometry, bond angle: $120^\circ$.
        • Example: $\text{BF}_3$.
     3. sp^3 Hybridization:
        • Tetrahedral geometry, bond angle: $109.5^\circ$.
        • Example: $\text{CH}_4$.
     4. sp^3d Hybridization:
        • Trigonal bipyramidal geometry, bond angles: $90^\circ$ and $120^\circ$.
        • Example: $\text{PCl}_5$.
     5. sp^3d^2 Hybridization:
        • Octahedral geometry, bond angle: $90^\circ$.
        • Example: $\text{SF}_6$.

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Polar and Non-Polar Bonds

1. Polar Covalent Bond:

   • Formed between atoms with different electronegativities.
   • The shared electrons are more attracted to the more electronegative atom, creating a dipole.
   • Example: $\text{HCl}$.

2. Non-Polar Covalent Bond:

   • Formed between atoms with identical or very similar electronegativities.
   • The shared electrons are evenly distributed.
   • Example: $\text{O}_2$, $\text{N}_2$.

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Dipole Moment

• Definition: A measure of the separation of charges in a molecule. $\mu = q \times d$ Where:
  • $\mu$ = Dipole moment,
  • $q$ = Magnitude of charge,
  • $d$ = Distance between charges.
• Unit: Debye (D).
• Applications:
  • Helps determine molecular polarity.
  • Example: $\text{HCl}$ has a dipole moment, while $\text{CO}_2$ does not due to its linear structure.

The VSEPR Theory (Valence Shell Electron Pair Repulsion)

• Overview: Explains the shapes of molecules based on the repulsion between electron pairs in the valence shell of the central atom.
• Key Principles:
  1. Electron pairs (bonding and lone pairs) arrange themselves to minimize repulsion.
  2. Lone pair-lone pair repulsion $>$ lone pair-bond pair repulsion $>$ bond pair-bond pair repulsion.
• Shapes of Molecules:
  • Linear: Bond angle = $180^\circ$ (e.g., $\text{CO}_2$).
  • Trigonal Planar: Bond angle = $120^\circ$ (e.g., $\text{BF}_3$).
  • Tetrahedral: Bond angle = $109.5^\circ$ (e.g., $\text{CH}_4$).
  • Trigonal Bipyramidal: Bond angles = $90^\circ$, $120^\circ$ (e.g., $\text{PCl}_5$).
  • Octahedral: Bond angle = $90^\circ$ (e.g., $\text{SF}_6$).

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Hydrogen Bonding

1. Definition:
   • A weak bond formed between a hydrogen atom covalently bonded to a highly electronegative atom (e.g., $\text{O}$, $\text{N}$, $\text{F}$) and another electronegative atom.
2. Types:
   • Intermolecular Hydrogen Bonding: Between molecules (e.g., in water, $\text{H}_2\text{O}$).
   • Intramolecular Hydrogen Bonding: Within the same molecule (e.g., in ortho-nitrophenol).
3. Effects of Hydrogen Bonding:
   • Increases boiling and melting points (e.g., water has a high boiling point).
   • Influences the structure of proteins and DNA.

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Resonance

1. Definition:
   • A phenomenon where a molecule or ion can be represented by two or more equivalent structures (resonance structures).
2. Examples:
   • Carbonate Ion ($\text{CO}_3^{2-}$):
     • Resonance structures: $\text{O=C-O}^- \leftrightarrow \text{O-C=O}^- \leftrightarrow \text{-C-O}_2^-$
   • Benzene ($\text{C}_6\text{H}_6$):
     • Alternating single and double bonds.
3. Resonance Energy:
   • The energy difference between the most stable resonance structure and the actual structure.
   • Example: Resonance stabilization in benzene.

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Molecular Polarity and Dipole Moment

1. Polarity:

   • Polar molecules have unequal distribution of charges due to differences in electronegativity.
   • Example: $\text{H}_2\text{O}$ (polar), $\text{CO}_2$ (non-polar due to symmetry).

2. Dipole Moment:

   • Measures the polarity of a molecule.
   • A molecule with a non-zero dipole moment is polar.
   • Example: $\text{HCl}$ has a dipole moment, while $\text{N}_2$ does not.

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Van der Waals Forces

1. Types:

   • London Dispersion Forces:
     • Weak forces caused by temporary dipoles in atoms or molecules.
     • Example: Noble gases like helium.
   • Dipole-Dipole Interactions:
     • Forces between polar molecules.
     • Example: $\text{HCl}$.
   • Dipole-Induced Dipole Forces:
     • Forces between a polar molecule and a non-polar molecule.
     • Example: $\text{H}_2\text{O}$ and oxygen gas.

2. Significance:

   • Influence boiling and melting points.
   • Example: Higher dispersion forces in larger molecules lead to higher boiling points.

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Metallic Bonding and Band Theory

1. Metallic Bonding:

   • Involves a lattice of positive metal ions immersed in a "sea" of delocalized electrons.
   • Explains properties of metals like conductivity and malleability.

2. Band Theory:

   • In solids, atomic orbitals combine to form energy bands.
   • Conduction Band: Electrons here can move freely.
   • Valence Band: Lower energy band filled with electrons.
   • Band Gap:
     • Conductors: Overlapping conduction and valence bands.
     • Insulators: Large band gap.
     • Semiconductors: Small band gap.

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Summary Table of Key Concepts

Concept	Key Idea	Example
Ionic Bond	Transfer of electrons	$\text{NaCl}$
Covalent Bond	Sharing of electrons	$\text{H}_2$, $\text{H}_2\text{O}$
Hydrogen Bonding	Weak bond between $\text{H}$ and electronegative atoms	Water, DNA
VSEPR Theory	Shape determined by electron pair repulsion	$\text{CH}_4$ (tetrahedral)
Metallic Bond	Positive ions in a sea of electrons	Copper, aluminum
Van der Waals Forces	Weak intermolecular forces	Noble gases, $\text{HCl}$