Structure and Bonding of Organic Molecules
Organic chemistry is the study of carbon‑containing compounds—their structures, properties, reactions, and synthesis.
Carbon’s ability to form strong covalent bonds with itself and with common heteroatoms (H, O, N, S, halogens) underlies the staggering diversity of organic molecules found in nature and the laboratory.
This lesson covers:
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1. electronic structure of the main organic elements
2. types of covalent bonds (σ and π) and their energies
3. hybridisation and three‑dimensional geometry
4. resonance and electron delocalisation
Carbon, nitrogen, oxygen, and the halogens build almost all of the molecules we study. Their s and p orbitals determine how they bond:
s orbital – spherical, holds two electrons.
p orbital – dumb‑bell shaped, three orientations (px, py, pz).
Electronic Structure of Hydrogen
Hydrogen has one electron in the 1s orbital and no electrons in the outer shell. In order to achieve a stable configuration, hydrogen needs to share electrons with other atoms. Hydrogen can form one covalent bond, which allows it to bond with other elements like carbon, oxygen, and nitrogen.
- Configuration: 1s¹
- Valence electrons: 1
- Forms one σ bond to complete a duet (2 e⁻).
Electronic Structure of Boron
Boron has five electrons, with two in the 1s orbital, three in the outer shell. In order to achieve a stable octet configuration, boron needs to share electrons with other atoms. Boron can form three covalent bonds, which allows it to bond with other elements like carbon, hydrogen, and fluorine.
- Configuration: 1s² 2s² 2p¹
- Valence electrons: 3
- Typically forms three σ bonds (e.g., BH₃). Boron is electron‑deficient and often violates the octet rule.
Electronic Structure of Carbon
Carbon has six electrons, with two in the 1s orbital and four in the outer shell. In order to achieve a stable octet configuration, carbon needs to share electrons with other atoms. Carbon can form four covalent bonds, which allows it to bond with other carbon atoms and with other elements like hydrogen, oxygen, nitrogen, and sulfur.
- Configuration: 1s² 2s² 2p²
- Valence electrons: 4 → can form four σ bonds via hybridisation (sp³, sp², sp).
Note: hybrid orbitals, not an excited “2s¹ 2p³” state, explain the four‑bond capacity.
Electronic Structure of Nitrogen
Nitrogen has seven electrons, with two in the 1s orbital and five in the outer shell. In order to achieve a stable octet configuration, nitrogen needs to share electrons with other atoms. Nitrogen can form three covalent bonds, which allows it to bond with other elements like carbon, hydrogen, and oxygen.
- Configuration: 1s² 2s² 2p³
- Valence electrons: 5 → forms three σ bonds plus one lone pair (e.g., NH₃).
Electronic Structure of Oxygen
Oxygen has eight electrons, with two in the 1s orbital and six in the outer shell. In order to achieve a stable octet configuration, oxygen needs to share electrons with other atoms. Oxygen can form two covalent bonds, which allows it to bond with other elements like carbon, hydrogen, nitrogen, and sulfur.
- Configuration: 1s² 2s² 2p⁴
- Valence electrons: 6 → forms two σ bonds plus two lone pairs (e.g., H₂O).
Types of Bonds in Organic Molecules
There are three types of covalent bonds in organic molecules: single bonds, double bonds, and triple bonds. Single bonds are the most common and are formed when two atoms share one pair of electrons. Double bonds are formed when two atoms share two pairs of electrons, and triple bonds are formed when two atoms share three pairs of electrons.
The strength of covalent bonds depends on the bond length and the bond order. Bond length is the distance between the nuclei of two bonded atoms, and bond order is the number of electron pairs shared between two atoms. As the bond length decreases and the bond order increases, the bond strength increases.
| Bond type | Construction | Bond order | Typical C–C energy (kJ mol⁻¹) | Rotation? |
|---|---|---|---|---|
| Single (σ) | head‑on overlap (s–s, s–p, or hybrid–hybrid) | 1 | ~350 | free |
| Double (σ + π) | one σ + one π (side‑on p–p) | 2 | ~610 | restricted |
| Triple (σ + 2 π) | one σ + two π | 3 | ~835 | linear, locked |
Three-Dimensional Structure of Organic Molecules
The three-dimensional structure of organic molecules is determined by the hybridization of the carbon atom and the arrangement of the atoms around it. Hybridization refers to the mixing of atomic orbitals to form new hybrid orbitals that have different shapes and energies. The most common hybridization states in organic chemistry are sp3, sp2, and sp.
In an sp3 hybridized carbon, the carbon atom forms four single bonds to other atoms, and the bonds are arranged in a tetrahedral geometry. In an sp2 hybridized carbon, the carbon atom forms two single bonds and one double bond, and the bonds are arranged in a trigonal planar geometry. In an sp hybridized carbon, the carbon atom forms one single and one triple bond, and the bonds are arranged in a linear geometry.
sp³ Hybridisation
Four equivalent sp³ orbitals adopt a tetrahedral geometry (109.5°). Sigma bonds form with each hydrogen in CH₄.
sp² Hybridisation
Three sp² orbitals lie in a trigonal‑planar arrangement (120°). The unhybridised p orbital participates in a pi (π) bond, yielding the C=C double bond in ethene.
sp Hybridisation
Two sp orbitals adopt a linear geometry (180°). The remaining two p orbitals form two π bonds, giving the C≡C triple bond of acetylene rigidity and high bond energy.
Hybridisation and Molecular Geometry
Hybrid orbitals are mathematical combinations of one s and up to three p orbitals.
The steric number (SN) = σ bonds + lone pairs on the central atom.
| Hybrid set | SN | Geometry (AXE) | Ideal angle | Example |
|---|---|---|---|---|
| sp³ | 4 | AX₄ (tetrahedral) | 109.5° | CH₄ |
| sp² | 3 | AX₃ (trigonal planar) | 120° | CH₂=CH₂ |
| sp | 2 | AX₂ (linear) | 180° | HC≡CH |
An sp‑hybridised carbon can bear a triple bond or two cumulative double bonds.
Three‑Dimensional Consequences
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Free rotation about σ bonds leads to conformational isomers (e.g., staggered vs gauche ethane).
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Restricted rotation about π bonds yields cis/trans (E/Z) stereoisomers.
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Stereocentres (sp³ atoms bonded to four different substituents) create enantiomers, mirror‑image molecules with identical physical data except optical activity and chiral interactions.
The arrangement of the atoms around the carbon atom also affects the reactivity of the molecule. For example, molecules with stereocenters, which are atoms that are bonded to four different substituents, can exist in two different three-dimensional configurations called enantiomers. Enantiomers have identical physical properties but interact differently with other chiral molecules and with polarized light.
Resonance and Electron Delocalisation
Certain molecules require multiple Lewis structures. These are resonance structures; the real molecule is a hybrid and made up of many states at the same time. Check out our Resonance Solver to help visualize the different transition states and different resonance structures! Or read further with our Resonance Structure Lesson
Some molecules require multiple valid Lewis structures (e.g., acetate, benzene).
Rules:
- Move only π electrons or lone pairs, never nuclei.
- Keep total charge constant.
- The weighted hybrid favours structures with full octets and minimal formal charge.
Summary
Quick Run Down:
σ and π bonds tells you which bonds can rotate.
Hybridisation predicts geometry and acidity trends (sp < sp² < sp³ pKa values).
Resonance explains charge distribution in intermediates.
Test Your Knowledge:
- How many electrons does a carbon atom have?
- What is the hybridization state of a carbon atom that forms two single bonds and one double bond?