Study Notes on Chemical Bonding

Introduction

• Fundamental Fact: Except for inert gases, no element exists as independent atoms under ordinary conditions.

Chemical Bonds

• Types of Bonds:
    - Ionic bond
    - Covalent bond
    - Co-ordinate bond
    - Metallic bond
    - Hydrogen bond
    - Van der Waals forces (intermolecular).

Classification of Chemical Bonds

• Strong Bonds:
    - Interatomic bonds: such as ionic, covalent, and metallic.

• Weak Bonds:
    - Mainly intermolecular forces such as hydrogen bonds and van der Waals forces.

Tendency to Acquire Minimum Energy

• When atoms approach, their nuclei and electrons repel each other.

• As attraction increases, the energy of the system decreases, leading to greater stability.

Chemical Bonding Explained

• Chemical Bond Definition:
    - A force that acts between two or more atoms to hold them together as a stable molecule.
    - Involves the redistribution of electrons among atoms, which is accompanied by a decrease in energy.
    - Thus, as energy decreases, stability increases, and the strength of the chemical bond increases.
    - Therefore, molecules are more stable than individual atoms.

Causes of Chemical Combination

• Attraction of Nuclei:
    - When two atoms approach, the nucleus of one atom attracts the electrons of another atom.
    - If the net result is attraction, the total energy of the system decreases, leading to bond formation.

• Bond Formation:
    - Represents an exothermic process (\Delta H < 0).

Potential Energy vs. Internuclear Distance Curve

Ionic or Electrovalent Bond

• Definition:
    - A chemical bond formed between two oppositely charged species as a result of the complete transfer of electrons from one species to another.
    - Electrostatic force of attraction between cation and anion results in the formation of an ionic bond.

Formation Process

• Electropositive (metal) atom loses electrons, while an electronegative (non-metal) atom gains electrons.

• Example:
    - Group 1 and Group 17 elements form ionic compounds effectively.
    - Electrovalency: The total number of electrons lost or gained during the formation of an ionic bond.
      - Example values:
        - Electrovalency of $Mg = 2$
        - Electrovalency of $Ca = 2$
        - Electrovalency of $O = 2$
        - Electrovalency of $Cl = 1$.

Bond Representation

• Compounds Formula Representation:
    1. Write the symbols of ions with the positive ion on the left and the negative ion on the right: $A^+ B^-$.
    2. Write their electro valency above each symbol: $A^x B^y$.
    3. Apply the crisscross rule for the formula: $A^y B^x$.
      - Example: For Calcium Chloride: $Ca^{2+} Cl^{-} <br>ightarrow CaCl_2$.

Lattice Energy (L.E.)

• Definition:
    - The amount of energy released when one mole of a crystal lattice is formed from its gaseous ion constituents or the energy required to dissociate one mole of lattice into gaseous ions.

• Stability Relation:
    - Higher lattice energy means greater stability and strength of the ionic compound.

• Factors Affecting Lattice Energy:
    1. Magnitude of Charge:
       - Lattice energy is directly proportional to the product of charge on ions: $L.E. \propto q_1 q_2$.
       - Example assumptions: NaCl < MgCl_2 < AlCl_3, where charge increases.   2. Size of Cations:      - Example assumption: LiCl > NaCl > KCl > RbCl > CsCl as cation size increases.

Conditions for Forming Ionic Bonds

1. Ionization Energy (I.E.):
      - The energy required to remove an electron from an isolated gaseous atom; lower I.E. means more tendency to form cations.

2. Electron Affinity (E.A.):
      - Energy released when an electron is added to a gaseous atom; higher E.A. means greater tendency to form anions.

3. Lattice Energy (L.E.):
      - Higher lattice energy correlates with stronger ionic bonds.

4. Electronegativity Difference:
      - Influences the nature and strength of the ionic bond.

Covalent Bond

• Definition:
    - Formed by mutual sharing of electrons between two atoms.
    - Characteristics: Shared electron pairs should possess opposite spins, localized between the concerned atoms.

• Types of Bonds Based on Electrons Shared:
     1. Single bond: (2 electrons shared) represented as $H-H$.
     2. Double bond: (4 electrons shared) represented as $O=O$.
     3. Triple bond: (6 electrons shared) represented as $N ext{ triple bond } N$.

• Covalency:
    - The number of covalent bonds formed by an atom in a molecule (e.g., Covalency of Hydrogen in $H_2$ is 1, Oxygen in $O_2$ is 2, Nitrogen in $N_2$ is 3).

Types of Electron Pairs

• Bond Pair: A pair of electrons involved in bonding.

• Lone Pair: Non-bonding pairs of electrons.

• Example:
    - Bond pair: 2
    - Lone pair: 4.

Co-ordinate Bond (Dative Bond)

1. Definition: A covalent bond where both electrons come from one atom, known as a coordinate bond.

2. Donor and Acceptor:
      - The atom providing the lone electron pair is the donor (Lewis base).
      - The atom accepting the electron pair is the acceptor (Lewis acid).
      - Necessary conditions for a coordinate bond:
        1. The octet of the donor atom should be complete and possess at least one lone pair.
        2. The acceptor atom should have a vacant orbital for accommodation of lone pairs.

3. Types:
      - Sigma (σ) coordinate bond: Formed by head-on overlap of atomic orbitals, with one atom donating both electrons.
      - Pi (π) coordinate bond: Formed through side-by-side overlap of atomic orbitals.

Examples of Coordinate Bonds

• $NH_3.BF_3$.
   

KÖssel-Lewis Approach to Chemical Bonding

• In 1916, Kössel and Lewis independently provided an explanation of chemical bonding through electrons.

• Lewis Structure Concept:
    - Atom represented with a positively charged 'Kernel' (nucleus + inner electrons) surrounded by outer shell locations for a maximum of eight valence electrons, called the octet.

• Atoms achieve stability through forming bonds, either through donation or sharing of electrons.

Octet Rule

• Atoms can combine either by:
    1. Transfer of valence electrons:
       - Gaining or losing electrons to achieve an octet.
    2. Sharing of valence electrons.

Lewis Dot Structures

Drawing Lewis Structures

1. Identify the central atom, usually less electronegative or present in smaller quantities.

2. Surrounding atoms are bonded directly to the central atom to achieve octet, adjusting bonds as necessary.

3. Represent lone pairs on both the central atom and surrounding atoms, ensuring octet completion.

Formal Charge Assignment

• Definition: Helps assign charge to individual atoms in a molecule, aiding in structure understanding.

• Formula for Calculation:
  $Q_F = N_A - N_L.P. - rac{1}{2}N_B.P.$
     Where:
     - $N_A$: Number of valence electrons in the free atom.
     - $N_L.P.$ : Number of lone pair electrons.
     - $N_B.P.$ : Number of bonding pair electrons.

Limitations of Octet Rule

1. Incomplete Octet: Molecules not fulfilling the eight electrons in outer shell (e.g., $BF_3$, $AlCl_3$).

2. Expansion of Octet: Atoms having more than eight outer electrons (e.g., $PCl_5$, $SF_6$).

3. Odd Electron Molecules: Molecules with unpaired electrons (e.g., $NO$).

Valence Bond Theory (VBT)

• Presented by Heitler and London, extended by Pauling and Slater.

• Key Points:
    - Covalent bond formation involves overlapping of half-filled orbitals.
    - Directional character of covalent bonds is due to overlaps.
    - The strength of covalent bonds is proportional to the extent of overlapping.
    - Overlapping types and its influence on bond strength involve:
      1. Nature of orbitals (overlap extent).
      2. Strength of bonds affected by orbital dimensions.
   

Concept of Variable Covalency

• Definition: Elements with empty orbitals in the outer shell show variable valencies through excited states of their electrons.

• Examples:
    - Nitrogen: Covalency 3 (exists as $NCl_3$), no excited states due to absence of empty orbitals.
    - Phosphorus: Covalency can be variable based on excited states (exists in $PCl_5$).

Types and Overlapping of Covalent Bonds

• Types:
    1. Sigma (σ) bonds: End-to-end overlap.
    2. Pi (π) bonds: Sidewise overlap.
    3. Delta (δ) bonds: Four lobes of d-orbitals overlap.

Bond Parameters

• Bond Order: Number of bonds between two atoms.

• Bond Length: Distance between nuclei of bonded atoms, influenced by electronegativity.

• Bond Energy: Energy required to break a bond, influenced by bond length and order.

Dipole Moment

• Definition: Product of charge and distance, used to express polarity.

• Calculated: $\mu = q \times d$.

• Significance: Helps in determining molecular polarity and % ionic character based on dipole moment values.

Hydrogen Bonds

• Definition: Attractive forces between hydrogen atoms covalently bonded to electronegative atoms and electronegative atoms of different molecules.

• Important Conditions:
    1. Hydrogen must be bonded to highly electronegative elements (N, O, F).
    2. Size of electronegative element must be small.

• Overall Importance: Affects boiling/melting points, viscosity, and solubility in compounds with hydrogen bonds.

Van Der Waals Forces

• Definition: Weak intermolecular forces not involving chemical bonds, significant for physical properties.

• Types:
    1. Dipole-dipole attraction.
    2. Dipole-induced dipole attraction.
    3. Instantaneous dipole-induced dipole forces.

Applications of Dipole Moment

1. Polar molecule determination.

2. % Ionic character in compounds.

3. Organic compound behavior.

Conclusion

• Understanding chemical bonding involves classifying bonds, recognizing factors that influence bond formation, and predicting molecular geometry and behavior based on these principles.

• The interactions outlined provide a foundation for predicting chemical properties and reactions across various contexts in chemistry.