Bless you, Bob Lee
--
oozing on the muggy shore of the gulf coast
l...@fbtc.net
--- Begin Message ---Greetings beautiful people, Color in organic moleculesWhite light is actually a mixture of light of many different colors, as can be readily demonstrated with a prism or diffraction grating. Each of the so-called primary colors is itself a continuum of many wavelengths, each of which has a specific energy described by the equation; E = hc/lambda Where E is the energy of the light ( in calories when Planck`s constant, h, is expressed as 1.59 X 10-34 cal sec), c is the speed of light (3.00 X 10+10 cm sec-1), and lambda is the wavelength of the radiation (in centimeters). In most of the colored inorganic substances (minerals,metals), light of a specific wavelength is absorbed in a process described as electronic transition involving d orbitals ( or f orbitals, or hybrid orbitals with some d or f character). Few organic compounds however, contain atoms having d-orbital electrons, and the more common colored organic molecules contain only carbon and hydrogen or, in some cases, nitrogen, oxygen, sulfur, or the halogens. In most cases then, electron transitions associated with color of organic compounds must be described in terms of s, p, sp, sp2, sp3 orbitals or the sigma or pi bonds formed from such orbitals. There are two common types of colored organic molecules, and both types are found to contain large , delocalized pi bond systems, In addition, one type involves unshared electron pairs (so-called n electrons). A rigorous treatment of the interaction of light with such electron systems is beyond the scope of this posting, but the brief qualitative discussion in the following examples is useful in visualizing the general considerations involved. Before considering these, we need to introduce the concept of *anti-bonding* orbitals. In our discussions of atomic and molecular orbitals, we have employed a much simplified approach. Any more rigorous treatment requires the use of alebraic signs (+ or -) for the wave functions describing the orbitals. Combination of atomic orbitals (or orbital lobes) of like sign result in the bonding (molecular) orbital of the type we have employed in our bond descriptions. If on the other hand , combinations of unlike signs are considered, the result is a higher-energy system called an antibonding orbital. For our purposes, we may consider such cases to represent excited states in which the bond pair electrons have zero probability of being located exactly between the nuclei involved. Example 1 Azobenzene is an orange crystalline solid that can exist in either the cis or trans form, the latter being the more stable around room temperature. Light energy of about 10-19 cal ( at 4400A) is believed to cause one of the unshared electrons (n electrons) on nitrogen to enter the delocalized pi system of the molecule in a delocalized antibonding orbital. Example 2 Azulene is a beautiful deep blue solid. It has strong absorption in the red region of the visible spectrum, so its observed color is blue (complementary to red). Again, a relatively low-energy transition occurs as a result of the absorption of light, this time described as a delocalized pi-pi transition. Multiple Bonds For many molecules and polyatomic ions, experimental evidence indicates that bonding involves more than one pair of electrons shared by adjacent nuclei. Because all experience suggests that two electrons represent the maximum occupancy of an atomic or molecular orbital, we need some bonding model that can account for bouble bonds (two shared electron pairs) and triple bonds (three shared pairs). The model will indicate , as in the case of the single bond , a maximum bond pair electron density within the general internuclear region, but only one molecular orbital directly between the nuclei (along the nuclear axis). The directional characteristics of atomic p orbitals provide us with a useful approach to this model, by suggesting a side-to-side overlap of the p orbital rather than an end-to-end overlap. The side-to-side overlap is referred to as a pi bond, and any bond along the internuclear axis is called a sigma bond. Our model then suggests that a double bond consists of one sigma and one pi bond and a triple bond consists of one sigma and two pi bonds. Since a triple bond represents the maximum occupancy of the general internuclear region, no higher degree of multiple bonding can occur, that is , the combination of one sigma and two pi bonds occupies all of the available internuclear region. There are only a few cases in which overlap of simple atomic orbitals can provide a reasonable approximation of covalent bonding. In fact the representations of the C2 molecule and the N2 molecule are only partially consistant with the observed properties of these molecules. The extension of our atomic orbital overlap model to include hybrid atomic orbitals provide a much more general and satisfying description of covalent bonding. Bless you Bob Lee -- oozing on the muggy shore of the gulf coast l...@fbtc.net -- The silver-list is a moderated forum for discussion of colloidal silver. To join or quit silver-list or silver-digest send an e-mail message to: silver-list-requ...@eskimo.com -or- silver-digest-requ...@eskimo.com with the word subscribe or unsubscribe in the SUBJECT line. To post, address your message to: silver-list@eskimo.com List maintainer: Mike Devour <mdev...@id.net>
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