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Organic molecules can adopt different three-dimensional arrangements (conformations) without breaking bonds. Each conformation has a different energy due to steric interactions (van der Waals repulsion, torsional strain). The lowest energy conformation is most stable and predominates at equilibrium.
Build molecular models and rotate single bonds to observe different arrangements. Calculate relative energies by identifying eclipsed vs staggered interactions. Draw energy diagrams showing conformation vs rotation angle.
Conformations are NOT the same as isomers—they interconvert rapidly at room temperature. The energy differences are small compared to bond-breaking energies. Not all atoms eclipse equally (geminal vs vicinal interactions matter differently).
You already know from alkane structure that rotation around C–C single bonds produces different spatial arrangements called conformations, and that staggered conformations are lower in energy than eclipsed ones. Conformational analysis takes this further by quantifying the energy costs of specific interactions, giving you a toolkit to predict which conformation predominates for any molecule and by how much.
The two main sources of strain are torsional strain and steric strain. Torsional strain arises from the repulsion between bonding electron pairs on adjacent carbons when they are forced into an eclipsed arrangement — even when the atoms involved are small hydrogens, this costs about 4 kJ/mol per eclipsing H–H interaction. Steric strain adds an additional penalty when bulky groups are forced close together. In butane, for instance, the eclipsed conformation where two methyl groups overlap costs significantly more than an H–H eclipse because the larger methyl groups have greater van der Waals repulsion. By assigning approximate energy values to each type of eclipsing interaction (H–H ≈ 4 kJ/mol, H–CH₃ ≈ 6 kJ/mol, CH₃–CH₃ ≈ 11 kJ/mol), you can estimate the relative energy of any conformation.
To analyze a molecule systematically, draw it as a Newman projection along each rotatable C–C bond, then rotate in 60° increments to survey all six key conformations (three staggered, three eclipsed). At each position, identify which groups are eclipsing or gauche and sum the strain energy contributions. Plot these values on an energy diagram with dihedral angle on the x-axis and relative energy on the y-axis. The result is the characteristic oscillating curve: energy minima at staggered conformations and maxima at eclipsed conformations, with the deepest minimum at the anti arrangement and the highest maximum where the largest groups eclipse.
The energy differences between conformations are small — typically 4–20 kJ/mol — compared to bond energies of 350+ kJ/mol. This means conformations interconvert millions of times per second at room temperature and cannot be isolated individually. However, the Boltzmann distribution tells you that lower-energy conformations are more populated. A 6 kJ/mol difference corresponds roughly to an 80:20 population ratio at room temperature. This quantitative thinking becomes critical when you move to cycloalkanes, where ring constraints lock certain conformational relationships in place and strain energies determine ring stability, chair preferences, and the axial-equatorial behavior of substituents on cyclohexane.