Chemical Bonding

Ionic Bond

Formation of ionic compounds primarily depends on following factors:

The ease of formation of positive and negative ions from neutral atoms.
The arrangement of positive and negative ions in the solid, i.e. the lattice of the crystalline compound.

Formation of a positive ion involves ionization, i.e. removal of electron from the neutral atom. On the other hand, formation of negative ion involves addition of electron to the neutral atom. So, ionization enthalpy is involved in formation of positive ion, while electron gain enthalpy is involved in formation of negative ion. Electron gain process may be endothermic or exothermic, but ionization is always endothermic. Hence, ionic bonds are formed more easily between elements with comparatively low ionization enthalpies and elements with comparatively high negative value of electron gain enthalpy.

Crystalline State: Ionic compounds in crystalline state consist of orderly three-dimensional arrangements of cations and anions held together by coulombic interaction energies. The crystalline structure of a particular compound is determined by the size of ions, their packing arrangements and other factors.

In ionic solids, the sum of electron gain enthalpy and ionization enthalpy may be positive. In spite of this, the crystal structure gets stabilized due to the energy released in the formation of crystal lattice. So, a qualitative measure of stability of an ionic compound is given by its enthalpy of lattice formation, and not simply by achieving octet of electrons.

Lattice Enthalpy: The energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions is called lattice enthalpy of the solid.

Example: Lattice enthalpy of NaCl is 788 kJ mol^-1. It means that 788 kJ of energy is required to separate one mole of NaCl into one mole of Na⁺ (g) and one mole of Cl^- (g) to an infinite distance.

The process of separation involves both the attractive forces between ions of opposite charges and the repulsive forces between ions of like charges. It is not possible to calculate lattice enthalpy directly from interaction of forces of attraction and repulsion only, because the solid crystal is a three-dimensional structure. We also need to include the factors associated with the crystal geometry.

Bond Parameters

Bond Length: The equilibrium distance between the nuclei of two bonded atoms in a molecule is called bond length.

Covalent Radius: Radius of an atom's core which is in contact with the core of an adjacent atom in a bonded situation is called covalent radius. Covalent radius is half of the distance between two similar atoms joined by a covalent bond in the same molecule.

The van der Waals Radius: The overall size of the atom (which includes its valence shell in a non-bonded situation) is called van der Waals radius. The van der Waals radius is half of the distance between two similar atoms in separate molecule in a solid.

Bond Angle: The angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion is called bond angle. Bond angle gives some idea about distribution of orbitals around the central atom in a molecule/complex ion. Hence, it helps in determining the shape of molecule/complex ion.

Bond Enthalpy: The energy required to break one mole of bonds of a particular type between two atoms in a gaseous state is called bond enthalpy. Example: The H-H bond enthalpy in hydrogen molecule is 435.8 kJ mol^-1

H₂ → H(g) + H(g); Δ_aH^v = 435.8 kJ mol^-1

Measurement of bond strength is more complicated in case of polyatomic molecules. For example; in case of H₂O molecule, the enthalpy needed to break the two O-H bonds is not the same.

H₂O(g) → H(g) + OH(g); Δ_aH₁^φ = 502 kJ mol^-1

OH(g) → H(g) + O(g); Δ_aH₂^φ = 427 kJ mol^-1

So, in polyatomic molecules, the term mean or average bond enthalpy is used. It is calculated by dividing the total bond dissociation enthalpy by the number of bonds broken. Following is the example of average bond enthalpy of water molecule.

Average bond enthalpy `=(502+427)/2=464.5` kJ mol^-1

Bond Order: In covalent bond, the number of bonds between two atoms in a molecule is called bond order. Bond order for single bond is 1, for double bond 2, and for triple bond it is 3. Isoelectric molecules and ions have identical bond orders. Bond enthalpy increases and bond length decreases with increase in bond order. Bond strength is least for single bond and most for triple bond. But single bond is more stable than double or triple bond. In fact, stability decreases with increasing order of bond.

Resonance Structure

It is often seen that a single Lewis structure is inadequate to represent a molecule in conformity with its experimentally determined parameters. In such cases, a number of structures with similar energy, positions of nuclei, bonding and non-bonding pairs of electrons are taken as the canonical structures. The hybrid of canonical structures accurately describes the molecule.

Example: For ozone (O₃), three structures are shown. Out of them, structures I and II are the canonical structures. Structure III shows the resonance hybrid and it represents the structure of O₃ more accurately. Resonance is represented by a double headed arrow.

Key Points of Resonance Structure:

Resonance stabilizes the molecule because energy of the resonance hybrid is less than the energy of any single canonical structure.
Resonance gives the average characteristic of the bond.
Canonical forms have no real existence.
The molecule does not exist for a certain fraction of time in one canonical form and for other fraction of time in other canonical form.
The resonance hybrid form is the single structure of molecule and it cannot be depicted by a single Lewis structure.