Chemical Bonding: Ionic, Covalent, Coordinate Bonds

• Chemical bonding is the attractive force that holds atoms and ions together.
• It broadly classifies as ionic, covalent, and coordinated covalent bonds.
• Theories have been made to answer why atoms combine and why molecules have specific shapes.

Ionic Bond

• An ionic bond is a bond between ions that makes a stable compound.
• Generated by the attractive force between oppositely charged ions.
• Cations have a positive charge, such as Na+.
• Anions have a negative charge, such as Cl-.

Covalent Bond

• Formed by the sharing of electrons between two bonding atoms.
• In methane (CH4), all C-H bonds form by mutual sharing of electrons from both carbon and hydrogen atoms.

Coordinated Covalent Bond

• formed by the sharing of an electron pair from one of the bonding atoms.
• An example is the bond formation between NH3 and BF3, where nitrogen and boron bond.

Octet Rule

• Developed by Kössel and Lewis in 1916, relates to chemical bonding's electronic theory.
• Atoms combine by transferring valence electrons or sharing valence electrons to achieve an octet in their valence shells.
• For example, a covalent bond forms between two chlorine atoms.
• Lewis dot structures represent electrons with dots.
• Atoms sharing one electron pair become joined by a single covalent bond.
• Atoms sharing of two pairs of electrons create a double bond, demonstrated in CO2 and C2H4 molecules.
• Atoms sharing of three electron pairs a triple bond is formed like in the N2 and ethyne molecules.
• Lewis dot structure depicts how molecules and ions bond in terms of shared electron pairs and the octet rule.

Limitations of the Octet Rule

• Incomplete octet present on the central atom
• Odd number of electrons
• Expanded Octet

Valence Shell Electron Pair Repulsion (VSEPR) Theory

• Based on the repulsive interactions of the valence shell electron pairs.
• Proposed by Sidgwick and Powell in 1940.
• Nyholm and Gillespie further developed and redefined it in 1957.
• The Lewis concept fails to account for the shapes of molecules and it proposes that the shape of molecules depends on the number of valence shell electron pairs (bonded/nonbonded) around the central atom.
• Pairs of electrons in the valence shell repel each other due to negatively charged electron clouds.
• The electrons maximize distance.
• The valence shell is spherical, with electron pairs localizing at maximum distances on the surface.
• Multiple bonds are treated as a single electron pair, with multiple pairs treated as one 'super pair'.
• Where molecules have two or more resonance structures, the VSEPR model applies to any such structure.
• Repulsive interactions vary, with lone pair-lone pair repulsion greater than lone pair-bond pair, which is greater than bond pair-bond pair.

Main Postulates of VSEPR Theory

• Shape depends on number of valence shell electron pairs (bonded or nonbonded) around the central atom.
• Electron pairs in shell repel one another.
• Electron pairs occupy positions to minimize repulsion to maximize distance.
• Valence shell is a sphere; electron pairs localize on the surface at maximum distance.
• Multiple bonds are treated as single pairs; two or three pairs are a 'super pair'.

Geometric Prediction of Molecules with VSEPR Theory

• Molecules are either those in which the central atom has no lone pair, or those in which the central atom has more than one lone pair -AB2 type molecules with 180° angle have a linear arrangement, and examples such as BeCl2, and HgCl2.
• AB3 molecules with a 120° arrangement have a trigonal planar geometry, for example, BF3.
• AB4 types are tetrahedral shaped, such as CH4, and NH4+
• AB5 are trigonal bipyramidal types
• AB6 are Octahedral types
• AB2E type molecules have a bent shape, such as SO2O3
• AB3E shape
• AB3E2 shape
• AB4E shape
• AB4E2 shape

Special Cases

• Ammonia (NH3)*
• Expected to be trigonal planar geometry, it is found to be pyramidal due to the non-bonding lone pair of nitrogen.
• Its angle is reduced to 107° from 109.5° due to repulsion from the lone pair.
• Water (H2O)*
• Found to be V-shaped instead of linear.
• Shape is due to non-bonding lone pairs of oxygen.
• Its tetrahedral shape, which has two lone pairs and two bond pairs.
• The angle reduces further to 104.5° because of the higher repulsion between lone pairs over lone and bond pairs
• Sulfur tetrafluoride (SF4)*
• Shape is trigonal bipyramidal not tetrahedral.
• One lone pair and Four bond pairs (S-F).

Chlorinetryfluoride (CIF3)

• Chlorine (Cl) central atom
• Geometry of CIF3 is neither planar nor trigonal planar
• There are two lone pairs and three bond pairs. (Cl-F).

VSEPR Theory abilities

• Predict geometry of large molecules, especially p-block element compounds accurately.
• Successfully determine geometry with small energy.

VSEPR Theory downfalls

• The theoretical basis is unclear and subject to discussion.
• Theory fails to explain the molecule's energy and stability, or describe the magnetic property.

Valance Bond Theory

• Valence bond theory details how atomic orbitals create bonds and how those bonds affect a molecule's shape.
• It defines the directional properties

VBT

• Introduced by Heitler and London (1927) and further by Pauling et al.
• Based on atomic orbitals, hybridization, and overlapping of orbitals.
• Overlapping of atomic orbitals, which results in pairing of electrons.
• The extent of overlap determines the covalent bond's strength.
• Greater overlap results in a stronger bond.
• Covalent bonds is formed by pairing of electrons that have opposite spins in the valance shell

Directional Bond Properties

• Valence bond theory explains the shape, formation, and directional properties of bonds in molecules like CH4, NH3, and H2O, with overlap and hybridized atomic orbitals.
• Overlap may be positive, negative, or zero, based on the sign/phase or orientation of the orbital wave function.
• Orbitals forming a bond should have same sign/phase and spatial orientation with postive overlap.

Bonding

• When orbitals are negative or out of phase, this leads to negative overlapping
• In phase orbitals lead to zero overlapping as the overall energy remains the same
• Sigma bonds are the primary form of bonding and lead to end on collisions
• Pi bonds are weaker
• There are s-s, s-p, and p-p overlapping

Pi(π) Bonding

• The atomic orbitals overlap in a way that their axes are parallel and perpendicular to the internuclear axis forming sauce shaped clouds.
• Sigma bonds are stronger compared to pi bonds.

Hybridization

• In order to explain the characteristic geometrical shapes of polyatomic molecule molecules like CH4, NH3 and H2O etc, Pauling Introduced the concept of hybridization.

Hybridisation Process:

• Intermixing the orbitals of slightly different energies resulting in the formation of new sets of orbitals of equivalent energies and shape.
• The number of hybrid orbitals is equal to the number of atomic orbitals that get hybridised.
• The hybrid orbital are always equivalent in energy and shape.
• The type of hybridisation indicates the geometry of the molecules.

Conditions for Hybridization:

• Orbitals present in valence shell are hybridized
• Orbitals have almost equal energy.
• Promotion of electrons is not essential prior condition.

Types of Hybridization

• SP, SP2, SP3, dsp2, SP2d, and SP3d2.

Hybridization Examples

• Central atoms Be, B, C, P and S are used to create the respective SP, SP2, SP3, dsp2, and SP3d2 structures.

Limitations of VBT

• VBT does not discuss the energy of molecule, and also fails to explain the specific magnetic properties of oxygen.
• VBT does not explain the relative stability of molecules like He2+.