The terms are symmetry designations, $\ce{\uppi\! molecular orbitals}$ being antisymmetric with respect to a defining plane containing at least one atom (e.g. the molecular plane of ethene) and $\ce{\upsigma\! molecular orbitals}$ symmetric with respect to the same plane. In practice the terms are used both in this rigorous sense (for orbitals encompassing the entire molecule) and also for localized two-centre orbitals or bonds, and it is necessary to make a clear distinction between the two usages. In the case of two-centre bonds, a $\ce{\uppi\!\mbox{-}bond}$ has a nodal plane that includes the internuclear bond axis, whereas a $\ce{\upsigma\!\mbox{-}bond}$ has no such nodal plane. (A $\ce{\updelta\!\mbox{-}bond}$ in organometallic or inorganic molecular species has two nodes.) Radicals are classified by analogy into $\ce{\upsigma\!\mbox{-}}$ and $\ce{\uppi\!\mbox{-}radicals}$. Such two-centre orbitals may take part in molecular orbitals of $\ce{\upsigma\!\mbox{-}}$ or $\ce{\uppi\!\mbox{-}symmetry}$. For example, the methyl group in propene contains three $\ce{C–H}$ bonds, each of which is of local $\ce{\upsigma\!\mbox{-}symmetry}$ (i.e. without a nodal plane including the internuclear axis), but these three '$\ce{\upsigma\!\mbox{-}bonds}$' can in turn be combined to form a set of group orbitals one of which has $\ce{\uppi\!\mbox{-}symmetry}$ with respect to the principal molecular plane and can accordingly interact with the two-centre orbital of $\ce{\uppi\!\mbox{-}symmetry}$ ($\ce{\uppi\!\mbox{-}bond}$) of the double-bonded carbon atoms, to form a molecular orbital of $\ce{\uppi\!\mbox{-}symmetry}$. Such an interaction between the $\ce{CH3}$ group and the double bond is an example of what is called hyperconjugation. This cannot rigorously be described as '$\ce{\upsigma\!\mbox{-}\uppi\!}$ conjugation' since $\ce{\upsigma\!}$ and $\ce{\uppi\!}$ here refer to different defining planes, and interaction between orbitals of different symmetries (with respect to the same defining plane) is forbidden.