Feeling atoms

When an atom or molecule approaches a surface, it feels an attractive force. The interaction potential between the atom or molecule and the surface, which depends on the distance between the molecule and the surface and on the lateral position above the surface, detemiines the strength of this force. The incoming molecule feels this potential, and upon adsorption becomes trapped near the minimum m the well. Often the molecule has to overcome an activation barrier, before adsorption can occur. 

The characteristic feature of valence bond theory is that it pictures a covalent bond between two atoms in terms of an m phase overlap of a half filled orbital of one atom with a half filled orbital of the other illustrated for the case of H2 m Figure 2 3 Two hydrogen atoms each containing an electron m a Is orbital combine so that their orbitals overlap to give a new orbital associated with both of them In phase orbital overlap (con structive interference) increases the probability of finding an electron m the region between the two nuclei where it feels the attractive force of both of them... 

Did we predict the number of atoms required to complete additional layers around the metal-coated C(jo correctly Figure 6 shows a spectrum of Qo covered with the largest amount of Ca experimentally possible (note the logarithmic scale). Aside from the edges of A = 32 and a = 104 which we have already discussed, there are additional clear edges at a = 236 and A = 448 (completion of a third layer was also observed at QoSr23g). Note that these values are identical to the ones just predicted above for the completion of the third and fourth layer of metal atoms. We, therefore, feel confident that the alkaline earth metals studied do, in fact, form the distinct layers around a central C50 molecule with the structures depicted in Fig. 5. 

The fluorine atom oh the left feels eight electrons... 

Now we can say why the chemical bond forms between two fluorine atoms. First, the electron affinity of a fluorine atom makes it energetically favorable to acquire one more electron. Two fluorine atoms can realize a part of this energy stability by sharing electrons. All chemical bonds form because one or more electrons are placed so as to feel electrostatic attraction to two or more positive nuclei simultaneously. 

Thus we can expect a stable molecular species, LiF. The term stable again means that energy is required to disrupt the molecule. The chemical bond lowers the energy because the bonding electron pair feels simultaneously both the lithium nucleus and the fluorine nucleus. That is not to say, however, that the electrons are shared equally. After all, the lithium and fluorine atoms attract the electrons differently. This is shown by the ionization energies of these two atoms ... 

Hemin is shown on the right in Figure 22-7. It is shown beside the model of chlorophyll A to emphasize the astonishing similarity. The portions within dotted lines identify the differences. Except for the central metal atom, tl.e differences are all on the periphery of these cumbersome molecules. We cannot help wondering how nature managed to standardize on this molecular skeleton for molecules with such different functions. We cannot avoid a feeling of impatience as we await the clarification of the possible relationship, a clarification that will surely be provided by scientists of the next generation. 

Of the three principal classes of crystals, ionic crystals, crystals containing electron-pair bonds (covalent crystals), and metallic crystals, we feel that a good understanding of the first class has resulted from the work done in the last few years. Interionic distances can be reliably predicted with the aid of the tables of ionic radii obtained by Goldschmidt1) by the analysis of the empirical data and by Pauling2) by a treatment based on modem theories of atomic structure. The stability,... 

I feel now that I was influenced to some extent by my knowledge that in 1926 Goldschmidt had formulated a set of atomic radii that represented reasonably well the interatomic distances in both covalent crystals and metals (5). I was also impressed by a discussion of the properties of metals by Bernal, who, however, rejected the idea that covalent bonds are present in metals (4). Bragg also rejected this idea (5). 

---