Delocalization leads to a lowering of energy levels and narrowing of the spacings between levels; and the larger the delocalized system, the greater this effect. This can be seen in the series of aromatic molecules compared at the bottom of the opposite page. Benzene has six atoms in its delocalized p- electron system, and the spacings between the p-orbital energy levels are such that it absorbs energy at a set of wavelengths in the ultraviolet region, centered around 2550 Å. The visible wavelengths pass through the molecule untouched, so benzene is colorless to our eyes. So are naphthalene and anthracene, which have 10 and 14 atoms in the delocalized system, although the larger rings shift the absorption to longer wavelengths or lower energies: 3150 Å and 3800 Å. In contrast, delocalization in naphthacene is so extensive that the splitting between ir energy levels has narrowed to the point where blue light around 4800 Å is absorbed. With the blue light absorbed, the remaining visible wavelengths make naphthacene appear orange, the complement of blue. In pentacene, which has five rings, absorption is shifted down to even lower energies. Pentacene removes yellow light around 5800 Å and therefore appears indigo. This "eyeball spectroscopy" is surprisingly informative in revealing what aromatic molecules are doing. The visible spectrum is shown at the right, with colors recorded as a function of wavelength from the ultraviolet to the infrared. If any of these wavelengths is absorbed by a molecule, the remaining wavelengths give the molecule the complementary color. Removal of green wavelengths around 5300 Å makes a molecule appear purple. If the molecule absorbs red light at around 6800 Å, we will see it as blue-green. By looking at what is left of the visible spectrum after absorption, we can decide approximately what visible wavelengths the compound is absorbing.