Covalent bonding is a fundamental concept in chemistry, essential for understanding the structure and behavior of molecules. This article explores the formation of covalent bonds in methanol (CH₃OH), ethylene (C₂H₄), oxygen (O₂), carbon dioxide (CO₂), and nitrogen (N₂). Tailored for the Cambridge IGCSE Chemistry - 0620 - Supplement syllabus, it delves into the intricacies of atomic sharing to form stable compounds, highlighting their significance in various chemical contexts.

Covalent bonds are chemical bonds formed by the sharing of electron pairs between atoms. Unlike ionic bonds, which involve the transfer of electrons, covalent bonding typically occurs between non-metal atoms with similar electronegativities. This sharing allows each atom to attain a stable electron configuration, often resembling that of noble gases. The general equation for a covalent bond formation can be represented as: $$ A + B \rightarrow A-B $$ where $A$ and $B$ are atoms sharing electrons. Lewis Structures: Lewis structures are diagrams that represent the bonding between atoms in a molecule and the lone pairs of electrons that may exist. For example, the Lewis structure of O₂ shows a double bond between the two oxygen atoms, each sharing two pairs of electrons: $$ O=O $$

2. Methanol (CH₃OH) Covalent Bonding

Methanol consists of one carbon atom bonded to three hydrogen atoms and one hydroxyl group (-OH). The carbon atom forms four covalent bonds: three with hydrogen and one with oxygen. Oxygen, in turn, shares two electrons with hydrogen, completing its valence shell. Lewis Structure of CH₃OH:

• Carbon (C): 4 valence electrons
• Hydrogen (H): 1 valence electron each
• Oxygen (O): 6 valence electrons

$$ H | H - C - O - H | H $$

3. Ethylene (C₂H₄) Covalent Bonding

Ethylene is an unsaturated hydrocarbon with a carbon-carbon double bond. Each carbon atom forms two single bonds with hydrogen atoms and one double bond with the other carbon atom. This double bond consists of one sigma ($\sigma$) bond and one pi ($\pi$) bond, allowing for the molecule’s planar structure and reactivity in addition reactions. Lewis Structure of C₂H₄: $$ H_2C=CH_2 $$

4. Oxygen (O₂) Covalent Bonding

Oxygen gas exists as diatomic molecules, where two oxygen atoms are connected by a double bond. Each oxygen atom shares two electrons, forming a stable O=O bond. This molecule is paramagnetic due to the presence of unpaired electrons in its molecular orbital configuration. Lewis Structure of O₂: $$ O=O $$

5. Carbon Dioxide (CO₂) Covalent Bonding

Carbon dioxide is a linear molecule where the carbon atom forms two double bonds with two oxygen atoms. Each double bond involves the sharing of two electron pairs, allowing carbon to achieve a complete octet and the molecule to be nonpolar despite the polar bonds due to its symmetrical geometry. Lewis Structure of CO₂: $$ O=C=O $$

6. Nitrogen (N₂) Covalent Bonding

Nitrogen gas consists of two nitrogen atoms connected by a triple bond. Each nitrogen atom shares three pairs of electrons, resulting in a strong N≡N bond. This triple bond makes N₂ a relatively inert molecule under normal conditions, contributing to its stability in the atmosphere. Lewis Structure of N₂: $$ N≡N $$

7. Bond Length and Bond Energy

The bond length is the distance between the nuclei of two bonded atoms. Multiple bonds (double, triple) are shorter and stronger than single bonds due to the increased sharing of electrons. Bond energy refers to the energy required to break a bond: $$ \text{Bond Energy} \propto \text{Bond Strength} $$ For example, the triple bond in N₂ has higher bond energy compared to the double bond in O₂.

8. Molecular Geometry and Hybridization

The shape of a molecule affects its physical and chemical properties. Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals suitable for bonding:

• sp³ Hybridization: Tetrahedral geometry, as seen in CH₃OH.
• sp² Hybridization: Trigonal planar geometry, as in C₂H₄ and CO₂.
• sp Hybridization: Linear geometry, observed in O₂ and N₂.

9. Electronegativity and Polar Covalent Bonds

Electronegativity is the tendency of an atom to attract electrons in a bond. Differences in electronegativity between bonded atoms can lead to polar covalent bonds, where electrons are unequally shared: $$ \text{Polar Bond} \rightarrow \delta^+ - \delta^- $$ For instance, in CH₃OH, the O-H bond is polar due to oxygen's higher electronegativity compared to hydrogen.

10. Resonance Structures

Some molecules can be represented by multiple valid Lewis structures called resonance structures, which depict delocalized electrons. However, resonance does not imply oscillation between structures but rather a hybrid that represents the correct electron distribution.

11. Formal Charge and Stability

Formal charge calculation helps in determining the most stable Lewis structure. The structure with formal charges closest to zero is generally more stable: $$ \text{Formal Charge} = \text{Valence Electrons} - (\text{Non-bonding Electrons} + \frac{\text{Bonding Electrons}}{2}) $$

Molecular Orbital (MO) Theory provides a more comprehensive understanding of bonding by considering the combination of atomic orbitals to form molecular orbitals that extend over the entire molecule. In diatomic species like O₂ and N₂, MO theory explains magnetic properties and bond order: $$ \text{Bond Order} = \frac{\text{Number of bonding electrons} - \text{Number of antibonding electrons}}{2} $$ For O₂:

• Bond Order = 2
• Resulting in a double bond
• Paramagnetic due to unpaired electrons

• Covalent bonds involve electron sharing between non-metal atoms to achieve stable electron configurations.
• Molecules like CH₃OH, C₂H₄, O₂, CO₂, and N₂ exhibit varying bond types and geometries based on their bonding structures.
• Advanced concepts such as molecular orbital theory and hybridization provide deeper insights into bond formation and molecular properties.
• Understanding bond polarity and molecular geometry is crucial for predicting physical and chemical behavior.
• Comparative analysis highlights the diversity in bonding, bond order, and molecule polarity across different compounds.