The second half of the double bond in ethylene arises from a combination of these last two p orbitals into a molecular orbital with z lobes of density of opposite sign above and below the plane of the molecule. The double bond is shorter than a single bond because the p. orbitals must come closer together before they overlap enough to bond. The orbital also forces the molecule to be planar. Twisting about a bond axis is harmless to a symmetrical s, bond, but breaks a p bond by pulling the p orbitals out of alignment. To twist one end of the ethylene molecule 90 relative to the other, one would have to supply energy equal to the difference between a C-C double bond and a single bond, or 147 - 83 = 64 kcal mole. The ideal H-C-H bond angle of 120 at each end of the ethylene molecule is decreased to 117 by electron-pair repulsion between the double bond and the two C-H single bonds. Double bonds are of great importance in biological molecules, both because they help make proteins and other molecules rigid and because of their unique ability to absorb light. We will come back to the structural rigidity aspects in the chapter on proteins, and to their light-absorbing properties in the postscript to this chapter.