Shapes of Organic Molecules; σ and π Bonds

· Organic molecules can be described as straight-chained, branched or cyclic.
· Molecular shape depends on the hybridisation of atoms involved in bonding: sp³, sp² or sp.
· Bond angles are determined by the arrangement of electron pairs/orbitals around the atom.
· For exam answers, always link hybridisation → shape → bond angle → σ/π bond arrangement.

Straight-Chained, Branched and Cyclic Organic Molecules

· Straight-chained molecules: carbon atoms are joined in one continuous chain, e.g. butane.
· Branched molecules: carbon chain has one or more side groups, e.g. methylpropane.
· Cyclic molecules: carbon atoms form a ring, e.g. cyclohexane.
· The same molecular formula can sometimes form different chain types, leading to structural isomerism.
· In displayed or skeletal formulae, check the carbon skeleton first before deciding whether a molecule is straight, branched or cyclic.

This diagram helps show how carbon atoms in a saturated straight-chain molecule use sp³ hybrid orbitals to form σ bonds. It also shows that rotation around a C–C single σ bond is possible. Source

sp³ Hybridised Atoms

· sp³ hybridisation occurs when one s orbital and three p orbitals mix to form four sp³ hybrid orbitals.
· Shape around an sp³ carbon = tetrahedral.
· Bond angle = 109.5°.
· Typical example: alkanes, e.g. methane, CH₄, and ethane, C₂H₆.
· Bonds made by sp³ carbon are σ bonds only.
· A single bond is always one σ bond.
· Example: in ethane, the C–C bond is a σ bond formed by end-on overlap of sp³ orbitals.

sp² Hybridised Atoms

· sp² hybridisation occurs when one s orbital and two p orbitals mix to form three sp² hybrid orbitals.
· Shape around an sp² carbon = trigonal planar.
· Bond angle = 120°.
· One unhybridised p orbital remains on each sp² carbon.
· In a C=C double bond, there is one σ bond and one π bond.
· The σ bond is formed by end-on overlap of sp² orbitals.
· The π bond is formed by sideways overlap of adjacent p orbitals above and below the plane of the molecule.
· Typical example: ethene, C₂H₄.
· Ethene is planar because each carbon is sp² hybridised and the π bond requires the p orbitals to remain parallel.

sp Hybridised Atoms

· sp hybridisation occurs when one s orbital and one p orbital mix to form two sp hybrid orbitals.
· Shape around an sp carbon = linear.
· Bond angle = 180°.
· Two unhybridised p orbitals remain on each sp carbon.
· In a C≡C triple bond, there is one σ bond and two π bonds.
· The σ bond is formed by end-on overlap of sp orbitals.
· The two π bonds are formed by sideways overlap of two separate pairs of p orbitals.
· Typical example: ethyne, C₂H₂.
· Molecules containing sp-hybridised atoms are often described as linear around the sp atom.

σ Bonds

· σ bond = covalent bond formed by direct end-on overlap of orbitals along the line between two nuclei.
· σ bonds can form from overlap such as s–s, s–p, p–p, sp³–sp³, sp²–sp² or sp–sp.
· Every single bond contains one σ bond.
· A double bond contains one σ bond + one π bond.
· A triple bond contains one σ bond + two π bonds.
· σ bonds are usually stronger than π bonds because end-on overlap is more effective than sideways overlap.

This diagram compares end-on overlap for σ bonds with sideways overlap for π bonds. It is useful for remembering that σ bonds lie along the internuclear axis, while π bonds exist above and below that axis. Source

π Bonds

· π bond = covalent bond formed by sideways overlap of adjacent p orbitals.
· π electron density is found above and below the σ bond framework.
· A π bond can only form when the adjacent p orbitals are parallel.
· π bonds restrict rotation because rotating would disrupt the sideways overlap of p orbitals.
· This explains why ethene is planar and why C=C bonds are important in later topics such as geometrical isomerism.
· π bonds are usually more reactive than σ bonds because the sideways overlap is less effective and easier to break.

This model shows the p orbitals involved in forming the π bond in ethene. It helps students visualise why the π bond lies above and below the plane of the molecule. Source

Planar Organic Molecules

· Planar means the relevant atoms lie in the same flat plane.
· Ethene is planar because both carbon atoms are sp² hybridised.
· Each carbon in ethene has trigonal planar geometry with bond angles close to 120°.
· The unhybridised p orbitals must remain parallel for the π bond to exist.
· Rotation around the C=C bond is restricted because it would break the π bond.
· In exam answers, use the phrase: “the molecule is planar because the carbon atoms are sp² hybridised and the p orbitals overlap sideways to form a π bond.”

Fast Exam Table: Hybridisation, Shape and Bonds

· sp³: 4 electron regions; tetrahedral; 109.5°; single bonds only; σ bonds only.
· sp²: 3 electron regions; trigonal planar; 120°; double bond possible; 1 σ + 1 π in C=C.
· sp: 2 electron regions; linear; 180°; triple bond possible; 1 σ + 2 π in C≡C.
· C–C single bond: 1 σ bond.
· C=C double bond: 1 σ bond + 1 π bond.
· C≡C triple bond: 1 σ bond + 2 π bonds.