Structure, Bonding & Reactivity (AQA AS Chemistry): Revision Note

Shapes & Bond Angles in Organic Molecules

• Carbon has four electrons in its outer shell (electronic configuration: 1s² 2s² 2p²)

• To achieve a full outer shell, carbon forms four covalent bonds by sharing its four outer electrons with other atoms

• Covalent bonds are formed by the overlap of atomic orbitals

  • There are two types of covalent bonds: sigma (σ) bonds and pi (π) bonds

• A sigma bond is formed by the direct overlap of orbitals along the line between two nuclei

• In a carbon–carbon double bond, each carbon forms three sigma bonds using three of its outer electrons

  • The remaining electron occupies a p orbital

• When the unhybridised p orbitals on two carbon atoms overlap sideways, a pi bond is formed

  • The π bond contains two electrons and lies above and below the plane of the atoms

• The three regions of electron density around each carbon repel equally, resulting in a trigonal planar shape with bond angles of approximately 120°

The bonding in ethene

• In ethene (C₂H₄), each carbon atom uses three of its four outer electrons to form sigma (σ) bonds

• Two σ bonds are formed between each carbon atom and two hydrogen atoms

• One σ bond is formed between the two carbon atoms

• The remaining electron on each carbon atom occupies an unhybridised p orbital

  • These p orbitals overlap sideways to form a pi (π) bond

• As a result, the carbon–carbon bond in ethene is a double bond, consisting of one σ bond and one π bond

Molecular Orbitals

Orbital overlap in covalent bonds

• A single covalent bond is formed when two atoms, usually non-metals, share a pair of electrons

• Each atom contributes one unpaired electron from an atomic orbital

  • When the bond forms, these atomic orbitals overlap to produce a shared region of electron density containing two electrons

  • This region of electron density can be described as a bonding molecular orbital

• The greater the overlap between the atomic orbitals, the stronger the bond formed

• Sigma (σ) bonds are formed by end-to-end (axial) overlap of orbitals along the line between the two nuclei

• Pi (π) bonds are formed by the sideways overlap of parallel p orbitals, creating regions of electron density above and below the sigma bond

σ bonds

• Sigma (σ) bonds are formed from the end-to-end overlap of atomic orbitals

• s orbitals overlap this way, as well as p orbitals

• In a sigma (σ) bond, the electron density is symmetrical about the line joining the nuclei of the two bonded atoms

• The shared pair of electrons is located directly between the nuclei

• The bond is formed due to the electrostatic attraction between the negatively charged shared electrons and the positively charged nuclei of both atoms

 Hydrogen

• The hydrogen atom has only one s orbital

• The s orbitals of the two hydrogen atoms will overlap to form a σ bond

π bonds

• Pi (π) bonds are formed by the sideways overlap of parallel p orbitals on adjacent atoms

• The overlapping p orbitals create regions of electron density above and below the plane of the sigma (σ) bond.

• This sideways overlap allows effective orbital interaction and forms a single π bond

• A π bond is often represented as two electron clouds, one above and one below the plane of the nuclei.

  • Together, these two regions of electron density contain one shared pair of electrons, forming one π bond