Chemical Bonding and Molecular Structure — Study Notes

1. Introduction

Atoms combine with one another to form molecules because the resulting species has lower energy and greater stability than the isolated atoms. The force that holds atoms together in a molecule or a crystal is called a chemical bond. This chapter explores why bonds form, the different types of bonds, and how the structure of molecules is predicted using simple and powerful models.

Bonding is ultimately an electrostatic phenomenon. Atoms achieve stable electronic configurations, usually an octet of electrons in their valence shell, by losing, gaining or sharing electrons. Kossel and Lewis independently proposed that this drive toward a stable octet is the underlying cause of chemical combination, leading to the Lewis octet rule. This insight gave chemists a simple language of dots and crosses for describing molecules.

Modern chemistry recognises a spectrum of bonds: the ionic bond, formed by complete electron transfer; the covalent bond, formed by electron sharing; and weaker interactions such as hydrogen bonds and van der Waals forces. The geometry of molecules, the polarity of bonds, and the hybridisation of atomic orbitals all determine the three-dimensional shape that gives a molecule its characteristic properties.

2. Kossel-Lewis Approach and the Octet Rule

Werner Kossel and G. N. Lewis proposed that atoms attain a stable configuration of eight electrons in their outermost shell, the octet, by forming bonds. The octet rule states that atoms combine to achieve an outer shell of eight electrons, like the noble gases. For small atoms such as hydrogen and helium, the stable configuration is a duplet of two electrons.

In the Lewis approach, the valence electrons of an atom are represented as dots around the symbol. For example, sodium has one dot, magnesium two, and chlorine seven. The Lewis dot structure of a molecule shows how the valence electrons are arranged among the atoms, whether shared or transferred. A bond formed by sharing a pair of electrons is called a covalent bond, and the shared pair constitutes a single bond.

The octet rule explains many simple molecules, but it has limitations. Molecules such as BeCl2 and BF3 have incomplete octets, whereas PCl5 and SF6 have expanded octets because the central atom uses d orbitals. Molecules like NO and NO2 have odd numbers of electrons. These exceptions make the octet rule a guideline rather than a strict law.

3. Ionic Bonding and Lattice Enthalpy

The ionic bond is formed by the complete transfer of one or more electrons from an electropositive atom to an electronegative atom. The resulting cations and anions are held together by strong electrostatic attraction. For example, sodium loses one electron to chlorine to form Na+ and Cl-, which attract each other in a giant three-dimensional lattice.

The driving force for ionic bond formation is the lowering of energy. Sodium has a low ionisation enthalpy and easily loses an electron, while chlorine has a highly negative electron gain enthalpy and readily accepts it. The strong electrostatic attraction releases a large amount of energy called lattice enthalpy, the energy released when one mole of an ionic compound is formed from its constituent ions in the gaseous state.

For an ionic bond to form, the lattice enthalpy must be large enough to compensate for the energy spent in ionisation. The melting and boiling points of ionic compounds are high because of the strong forces in the lattice. Ionic compounds are hard, brittle and crystalline, conduct electricity when molten or in aqueous solution, but not in the solid state, and have high melting points.

4. Covalent Bonding and Lewis Structures

A covalent bond forms when two atoms share one or more pairs of electrons. If one pair is shared, a single bond is formed; two pairs give a double bond and three pairs a triple bond. The number of shared pairs equals the bond order. In a covalent bond, the shared electron pair is attracted by the nuclei of both atoms, holding the molecule together.

Writing Lewis structures requires counting the total valence electrons, arranging atoms, and sharing electrons so each atom achieves an octet. For example, carbon dioxide has 16 valence electrons and forms two C=O double bonds. A coordinate or dative bond forms when one atom provides both electrons of the shared pair, as in the ammonium ion, where nitrogen donates a lone pair to the proton.

Covalent compounds are usually gases, liquids or low-melting solids because the intermolecular forces are weak. They are poor conductors of electricity and are generally insoluble in water but soluble in organic solvents.

5. Formal Charge and the VSEPR Theory

Formal charge is a bookkeeping device used to judge the quality of a Lewis structure. It is calculated as:

$$\text{Formal charge} = \text{valence electrons} - (\text{lone pair electrons} + \frac{1}{2} \times \text{bonding electrons})$$

The most plausible structure is the one with the lowest formal charges and with negative charges on the more electronegative atoms.

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular geometry. It states that electron pairs around the central atom arrange themselves as far apart as possible to minimise repulsion. Lone pairs repel more strongly than bond pairs, so the order of repulsion is lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.

The shapes derived from VSEPR include linear (two pairs, 180 degrees), trigonal planar (three pairs, 120 degrees), tetrahedral (four pairs, 109.5 degrees), trigonal bipyramidal (five pairs) and octahedral (six pairs). Ammonia is pyramidal rather than tetrahedral because one position is occupied by a lone pair, and water is bent with a bond angle of about 104.5 degrees.

6. Valence Bond Theory and Hybridisation

The valence bond theory states that a covalent bond forms when two atoms approach each other and a pair of electrons occupies the overlapping region of their atomic orbitals. Overlap can be head-on (sigma bond) or sideways (pi bond). A sigma bond is always formed first and is stronger, while a pi bond forms only when a sigma bond already exists.

Hybridisation is the mixing of atomic orbitals of nearly the same energy to form new equivalent orbitals called hybrid orbitals. In sp hybridisation, one s and one p orbital mix to form two linear orbitals at 180 degrees, as in BeCl2 and CO2. In sp2 hybridisation, one s and two p orbitals mix to form three trigonal-planar orbitals at 120 degrees, as in BF3 and C2H4. In sp3 hybridisation, one s and three p orbitals form four tetrahedral orbitals at 109.5 degrees, as in CH4, NH3 and H2O.

Sigma bonds are formed by sp3-sp3, sp2-sp2 or sp-sp orbital overlaps, while pi bonds result from sideways overlap of unhybridised p orbitals. A double bond consists of one sigma and one pi bond; a triple bond consists of one sigma and two pi bonds.

7. Molecular Orbital Theory

The molecular orbital (MO) theory describes electrons as occupying orbitals that belong to the entire molecule rather than to individual atoms. Atomic orbitals combine to form bonding and antibonding molecular orbitals. A bonding orbital has lower energy and higher stability, while an antibonding orbital has higher energy. Two atomic orbitals always produce two molecular orbitals.

The number of electrons in bonding and antibonding orbitals determines bond order:

$$\text{Bond order} = \frac{1}{2} \times (\text{number of electrons in bonding MOs} - \text{number of electrons in antibonding MOs})$$

A positive bond order indicates a stable molecule. H2 has a bond order of 1, He2 has a bond order of 0 and does not exist, and O2 has a bond order of 2 with two unpaired electrons in antibonding orbitals, explaining its paramagnetic behaviour. The molecular orbital approach successfully explains the paramagnetism of O2, which the valence bond theory cannot.

8. Polarity of Bonds and Dipole Moment

When atoms of different electronegativity share electrons, the shared pair is pulled toward the more electronegative atom, creating a polar covalent bond. The dipole moment, measured in debye (D), is the product of charge and distance of separation:

$$\mu = q \times d$$

The dipole moment of a molecule is the vector sum of the individual bond dipoles. In CO2, the two C=O bond dipoles cancel because the molecule is linear, giving zero net dipole moment. In water, the two O-H bond dipoles add vectorially to give a net dipole moment of 1.85 D because water is bent. Thus, molecular geometry determines whether a molecule is polar or non-polar.

Dipole moment is used to predict the shape of molecules and to distinguish between possible structures. For example, the existence of a zero dipole moment for CO2 proves that the molecule is linear rather than bent.

9. Hydrogen Bonding and van der Waals Forces

A hydrogen bond forms when a hydrogen atom covalently bonded to a highly electronegative atom (F, O or N) is attracted to a lone pair of electrons on another electronegative atom. It is represented as X-H...Y and is much weaker than a covalent bond but stronger than van der Waals forces. Hydrogen bonds can be intermolecular (between molecules) or intramolecular (within the same molecule).

Intermolecular hydrogen bonding in water explains its anomalously high boiling point, high surface tension and the fact that ice is less dense than liquid water. Intramolecular hydrogen bonding occurs in molecules such as o-nitrophenol. The strength of hydrogen bonding increases with the electronegativity of the bonded atom, in the order F > O > N.

van der Waals forces are weak intermolecular attractions that arise from temporary dipoles. Their strength increases with molecular size and surface area, which is why larger molecules boil at higher temperatures.

Quick Revision Tables

Table 1: Types of Bonds

Table 2: Hybridisation and Geometry

Table 3: Bond Order Examples

Mind Map

Important Diagrams (SVG)

Diagram 1: VSEPR Molecular Shapes

Diagram 2: Sigma and Pi Bonds in Ethene

Common Mistakes

1. Applying the octet rule to every molecule; BeCl2 and BF3 have incomplete octets, while PCl5 and SF6 have expanded octets.
2. Confusing bond order calculations in MO theory; remember the formula (Nb - Na)/2, where Nb and Na are bonding and antibonding electrons.
3. Treating CO2 as polar because C-O bonds are polar; the linear geometry cancels the bond dipoles, giving zero net dipole moment.
4. Writing sp3 hybridisation for ethene carbon; the carbon in C2H4 is sp2 hybridised because of the double bond.
5. Forgetting that a coordinate bond donates both electrons from one atom, as in NH4+ and H3O+.
6. Predicting the shape from the number of atoms instead of electron pairs; ammonia is pyramidal and water is bent because of lone pairs.
7. Assuming ionic compounds conduct electricity in the solid state; they conduct only when molten or dissolved.

Exam Tips

1. Memorise the geometry and bond angles for sp (180), sp2 (120) and sp3 (109.5) hybridisation; shapes questions are very frequent.
2. Remember that NH3 has a bond angle of 107 degrees and H2O 104.5 degrees because lone pairs compress the angle from the ideal 109.5 degrees.
3. Know that O2 is paramagnetic due to two unpaired electrons in antibonding MOs, a key experimental distinction from the valence bond theory.
4. Practise writing Lewis structures for common molecules such as CO2, NH3, H2O and CH4, paying attention to formal charge.
5. Learn that hydrogen bonding causes the anomalous properties of water, including its high boiling point and the density anomaly of ice.
6. For dipole moment questions, always consider the molecular shape, not just the presence of polar bonds.
7. Remember that sigma bonds are always present in single, double and triple bonds, while pi bonds exist only in multiple bonds.

Conclusion

Chemical bonding explains how and why atoms combine, and it determines the structure and properties of all substances. The octet rule, refined by Lewis structures and formal charge, provides a simple starting point, while the VSEPR theory predicts molecular geometry from electron-pair repulsion. Valence bond theory describes bonds in terms of orbital overlap and hybridisation, and molecular orbital theory accounts for properties such as the paramagnetism of oxygen. The polarity of bonds and molecules, along with weaker forces like hydrogen bonding, govern the physical behaviour of substances. Together these concepts form the theoretical backbone of chemistry, explaining everything from the shape of water to the reactivity of multiple bonds in organic molecules.