01 Structure and Bonding in Organic Molecules
Learn how bonding, molecular representations, geometry, polarity, and resonance work together to explain organic structure and reactivity.
Atoms and covalent bonds
Organic molecules consist of atoms joined by covalent bonds. Their connectivity, electron distribution, and three-dimensional shape help explain properties such as polarity and stability, as well as where reactions may occur.
A covalent bond is a pair of electrons shared between atoms. A single bond contains one shared pair, a double bond contains two, and a triple bond contains three. In the valence-bond model, a single bond is a formed by end-to-end orbital overlap. A double bond contains one and one , formed by sideways overlap of p orbitals; a triple bond contains one and two pi bonds.
Takeaway: Bond type describes the number of shared electron pairs, while sigma and pi describe kinds of orbital overlap.
and molecular representations
are the outer-shell electrons that participate in bonding. Carbon has four and commonly forms four bonds; hydrogen commonly forms one, oxygen two, and nitrogen three. These are useful patterns, but evaluating a structure also requires checking each atom’s bonds, lone pairs, and formal charge.
Molecular representations emphasize different features:
A molecular formula, such as
C₂H₆O, gives the number of each kind of atom but not their connectivity.A Lewis structure shows bonds and lone pairs, making it useful for tracking electrons and formal charge.
A condensed formula, such as
CH₃CH₂OH, writes groups compactly.A skeletal, or line-angle, formula uses lines for carbon–carbon bonds. Each line end and vertex represents a carbon; carbon-bound hydrogens are usually omitted. Heteroatoms and hydrogens attached to them are shown.
In three-dimensional drawings, solid wedges point toward the viewer, hashed wedges point away, and ordinary lines lie in the plane of the page.
Choose the representation that makes the feature you need to inspect easiest to see. A drawing models a molecule; it is not the molecule itself.
and molecular shape
is a valence-bond model for describing the arrangement of orbitals around an atom. Count electron-density regions around the atom: each single, double, or triple bond counts as one region, as does each lone pair. The count suggests a common and approximate geometry:
Four regions:
sp³, typically tetrahedral, with an angle of about ; for example, the carbon in methane,CH₄.Three regions:
sp², typically trigonal planar, with angles of about ; for example, each carbon in ethene,H₂C=CH₂.Two regions:
sp, typically linear, with an angle of ; for example, each carbon in ethyne,HC≡CH.
In sp² carbon, three hybrid orbitals form sigma bonds in a plane, leaving an unhybridized p orbital that can form a . In sp carbon, two hybrid orbitals point in opposite directions, leaving two p orbitals available for pi bonding. These labels are explanatory models, not physical components that can be directly seen in a molecule.
Takeaway: Electron-density regions help predict local geometry, and the orbital model connects that geometry to sigma and pi bonding.
and polarity
A bond is polar when its atoms attract shared electrons unequally. Differences in commonly cause this unequal sharing. The more electron-attracting atom carries a partial negative charge, written δ−, and the other carries a partial positive charge, written δ+. In a carbonyl bond, C=O, oxygen attracts electron density, leaving the carbonyl carbon partially positive.
Molecular polarity depends on bond polarity and three-dimensional shape. Bond dipoles add as vectors and may reinforce or cancel. The two polar carbon–oxygen bonds in linear carbon dioxide point in opposite directions and cancel, so the molecule has no net dipole. In formaldehyde, H₂C=O, the dipoles do not cancel, so the molecule is polar.
Takeaway: To determine molecular polarity, consider both the direction of individual bond dipoles and the molecule’s overall shape.
Resonance and electron delocalization
Some molecules and ions cannot be adequately represented by one Lewis structure. are alternative electron-bookkeeping drawings with the same positions of atoms but different placements of pi electrons, lone pairs, or formal charges. They are linked by a resonance arrow, not an equilibrium arrow. The actual molecule is a : its electrons are delocalized, and it does not switch back and forth between the drawings.
For example, acetate has two equivalent :
CH₃–C(=O)–O⁻ ↔ CH₃–C(–O⁻)=O
The drawings show that the negative charge is shared across both oxygen atoms. In the actual ion, the two carbon–oxygen bonds are equivalent. When drawing , keep the atoms in place and move only electrons; preserve the total charge and reasonable valence.
Delocalizing charge over several atoms often stabilizes a species. This stabilization can affect acidity and reaction pathways. For example, acetate is more stabilized than a structure with the same negative charge confined to one oxygen.
From structure to reactivity
Structure helps identify regions where electron density is concentrated, bonds that may change, and factors that can stabilize a product or intermediate.
Polar bonds create reactive sites. In a carbonyl, the carbon is partially positive and can be attacked by an electron-pair donor called a nucleophile.
Pi bonds provide accessible electron density. Their electrons are more exposed than sigma-bond electrons and can participate in reactions such as addition across a carbon–carbon double bond.
Resonance can stabilize charge. Delocalization may change the stability of species and influence reaction pathways.
Geometry affects electron delocalization. In a planar conjugated system, p orbitals can align and overlap. Twisting the system out of alignment can reduce overlap and change stability and reactivity.
These structural patterns guide predictions but do not determine every reaction on their own. Solvent, reagents, and reaction conditions also matter.
Takeaway: Bonding, electron distribution, and shape work together to explain likely reactive sites and the relative stability of chemical species.