Valence Bond Theory describes covalent bond formation in terms of the overlap of atomic orbitals on adjacent atoms. Two fundamental types of overlap occur in covalent bonds. A sigma (\sigma) bond forms by end-to-end (axial) overlap of orbitals: this can be s-s overlap (as in H-H), s-p overlap (as in H-F), or p-p overlap along the internuclear axis. A pi (\pi) bond forms by lateral (sideways) overlap of two parallel p orbitals that are perpendicular to the internuclear axis. The pi bond has two lobes of electron density, one above and one below the internuclear axis, unlike the cylindrically symmetric sigma bond. Option 'End-to-end overlap of two s orbitals' describes sigma bond formation in H_2, not a pi bond. Option 'End-to-end overlap of one s and one p orbital' describes sigma bond formation in molecules like HF or HCl. Option 'Lateral overlap of an s orbital and a p orbital' does not occur because s orbitals are spherically symmetric and cannot undergo directional lateral overlap with p orbitals. The concept of sigma and pi bonds through axial and lateral overlaps respectively is fundamental to NCERT valence bond theory and is widely tested in JEE. Plausibility: the weaker strength of pi bonds relative to sigma bonds (pi bonds are harder to break in isolation only when present as part of a double bond) is consistent with the less effective lateral overlap compared to axial overlap.