Hybridisation, conjugation and delocalisation in organic molecules
Organic structure has been drawn with lines that all look alike. Orbital mixing sets geometry. When p orbitals line up, electrons can be shared over several atoms, affecting bond length, strength and reactivity.
What a learner can do afterwards
- Assigns sp2 hybridisation to ethene carbon from stated sigma-bond and lone-pair counts, and relates sp3 and sp carbon models to ideal tetrahedral angles and linear geometry
- Explains why neighbouring p-orbital alignment favours coplanarity in a conjugated region and why delocalisation restricts amide C-N rotation
- Ranks otherwise comparable anion contributors with full octets and equal charge separation by negative-charge location, and rejects a contributor that moves a nucleus
- Uses delocalisation to explain a measured bond length that lies between single and double
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Hybridisation is a local bonding model. Combining one s orbital with three p orbitals gives four sp3 hybrids; two p orbitals give three sp2 hybrids; one p orbital gives two sp hybrids. Their ideal geometries are tetrahedral (109.5 degrees), trigonal planar (120 degrees) and linear (180 degrees). The sp2 model leaves one unhybridised p orbital; the sp model leaves two.
Without conjugation, count sigma bonds and local lone pairs: four suggests sp3, three sp2, two sp. A double bond contains one sigma and one pi bond; a triple bond has one sigma and two pi bonds. Each counts once in the sigma count. Alcohol oxygen has two sigma bonds and two lone pairs, so is usually sp3. Amine nitrogen is usually sp3 with three sigma bonds and one local lone pair. Conjugating lone pairs need a separate check.
Conjugation needs continuous overlap between neighbouring p orbitals. Alternating bonds, suitable lone pairs or empty p orbitals can take part. Coplanarity often improves overlap. Twisting reduces it without breaking a sigma bond; the whole molecule need not be flat.
Resonance contributors keep nuclei, electron count and total charge fixed. They describe one electronic structure. The molecule does not switch between contributors. For common C, N and O anions, compare valid octets, then charge separation. With these comparable, negative charge on oxygen usually gives a larger contribution than on carbon.
An amide nitrogen's lone pair occupies a p orbital overlapping with the carbonyl. Its three sigma bonds use roughly planar sp2 hybrids. Carbonyl carbon and oxygen are also described as sp2. Partial C-N double-bond character makes twisting cost energy, so rotation is restricted.
Carboxylate ions have two equivalent C-O bonds, intermediate in length between comparable single and double bonds. Delocalisation gives both partial double-bond character. Intermediate lengths support this model, but length alone is not proof. The value need not be exactly halfway.
Check sigma bonds, lone pairs and conjugation before assigning hybrids. Delocalisation can change bond lengths and restrict rotation.
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Counting bonds without losing the lone pairs
Start a hybridisation assignment by identifying the atom and its sigma bonds. A bond to one neighbour contributes one sigma bond even when it is drawn as a double or triple bond. The additional bonds are pi bonds.
For a localised bonding description, add the atom's lone pairs to its sigma-bond count. Four directions suggest sp3, three suggest sp2 and two suggest sp. The corresponding ideal arrangements are tetrahedral, trigonal planar and linear. Lone pairs occupy directions too, so orbital geometry and molecular shape are not interchangeable. An alcohol oxygen usually has an sp3 description, although only two of its four directions point towards bonded atoms.
Before accepting the count, look for conjugation. At an amide nitrogen, the lone pair participates in overlap with the carbonyl through a p orbital. The three sigma bonds are described using sp2 hybrids. Counting that lone pair as another local hybrid would miss this relationship. Hybridisation is a useful bonding model; the surrounding groups can make measured angles differ from ideal values.
What a resonance contributor represents
Several Lewis drawings can describe the same molecule or ion when different electron placements are possible within one nuclear framework. These drawings are resonance contributors. They must preserve the atoms, total electron count and overall charge. The molecule does not alternate between the drawings.
To compare contributors for ordinary carbon, nitrogen and oxygen compounds, first check electron counts and octets. A drawing that gives carbon more than eight valence-shell electrons is not valid in this model. Among valid contributors, full octets and less charge separation generally favour a greater contribution. When these factors are comparable, negative charge is usually favoured on a more electronegative atom.
A carboxylate group illustrates equivalent contributors. One drawing puts the C=O bond at one oxygen; another puts it at the other oxygen. The actual ion has equivalent C-O bonds with partial double-bond character. The contributors help account for those bonds without assigning one oxygen a permanently single bond.
Overlap, rotation and bond length
Sigma overlap lies along the line joining bonded nuclei. Pi overlap is sideways, commonly between p orbitals, and places electron density on either side of that line. A double bond contains one sigma bond and one pi bond. A triple bond contains one sigma bond and two pi bonds.
Turning about a sigma-bond axis can preserve sigma overlap. The same turn can disrupt neighbouring p-orbital overlap. This helps explain the energy barrier to rotation at a double bond and restricted rotation in conjugated groups. Other interactions can also oppose rotation, so a single bond does not guarantee unrestricted movement.
Bond lengths provide another check on a structural model. Compare bonds between the same elements in similar chemical surroundings. Partial multiple-bond character can give a length between typical single- and double-bond values. The result need not be their arithmetic midpoint. Use the measurement alongside connectivity, symmetry and other evidence instead of treating one intermediate length as proof of delocalisation.
Why conjugation can affect colour
A conjugated region has a continuous set of neighbouring orbitals that can share pi electron density. Alternating single and double bonds are a familiar example, but suitably aligned lone pairs and empty p orbitals can also participate.
Light can supply the energy needed for an electronic transition. In many related molecules, extending conjugation reduces the energy of the relevant transition and shifts absorption towards longer wavelengths. This is a comparison trend, not a rule that determines every molecule's colour from its number of double bonds.
Beta carotene has an extended conjugated region that absorbs some visible wavelengths. The remaining visible light gives it an orange appearance. Many shorter conjugated systems absorb ultraviolet light instead, so conjugation does not always produce a visible colour. The arrangement of the orbitals and the rest of the molecular structure also affect absorption.
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