Bonds, Van der Waals

The van der Waals bonding force results from an interaction between hydrophobic molecules (for example, between two aromatic residues such as phenylalanine (Fig. 3.11) or between two aliphatic residues such as valine). It arises from the fact that the electronic distribution in these neutral, non-polar residues is never totally even or symmetrical. As a result, there are always transient areas of high electron density and low electron density, such that an area of high electron density on one residue can have an attraction for an area of low electron density on another molecule. 

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Dipoles attract such that their energy varies as 1 /r. Thus the energy of the Van der Waals bond has the form... 

This is because rubber, like many polymers, is composed of long spaghetti-like chains of carbon atoms, all tangled together as we showed in Chapter 5. In the case of rubber, the chains are also lightly cross-linked, as shown in Fig. 5.10. There are covalent bonds along the carbon chain, and where there are occasional cross-links. These are very stiff, but they contribute very little to the overall modulus because when you load the structure it is the flabby Van der Waals bonds between the chains which stretch, and it is these which determine the modulus. 

Well, that is the case at the low temperature, when the rubber has a proper modulus of a few GPa. As the rubber warms up to room temperature, the Van der Waals bonds melt. (In fact, the stiffness of the bond is proportional to its melting point that is why diamond, which has the highest melting point of any material, also has the highest modulus.) The rubber remains solid because of the cross-links which form a sort of skeleton but when you load it, the chains now slide over each other in places where there are no cross-linking bonds. This, of course, gives extra strain, and the modulus goes down (remember, E = [Pg.61]

Fig. 6.2. How Young s modulus increases witl) increasing density of covalent cross-links in polymers, including rubbers above tbe glass temperature. Below To, be modulus of rubbers increases markedly because tbe Van der Waals bonds take hold. Above Tq they melt, and the modulus drops.

Is it possible to make polymers stiffer than the Van der Waals bonds which usually hold them together The answer is yes - if we mix into the polymer a second, stiffer, material. Good examples of materials stiffened in this way are ... 

Polymers, too, creep - many of them do so at room temperature. As we said in Chapter 5, most common polymers are not crystalline, and have no well-defined melting point. For them, the important temperature is the glass temperature, Tq, at which the Van der Waals bonds solidify. Above this temperature, the polymer is in a leathery or rubbery state, and creeps rapidly under load. Below, it becomes hard (and... 

Creep of polymers is a major design problem. The glass temperature Tq, for a polymer, is a criterion of creep-resistance, in much the way that is for a metal or a ceramic. For most polymers, is close to room temperature. Well below Tq, the polymer is a glass (often containing crystalline regions - Chapter 5) and is a brittle, elastic solid -rubber, cooled in liquid nitrogen, is an example. Above Tq the Van der Waals bonds within the polymer melt, and it becomes a rubber (if the polymer chains are cross-linked) or a viscous liquid (if they are not). Thermoplastics, which can be moulded when hot, are a simple example well below Tq they are elastic well above, they are viscous liquids, and flow like treacle. 

An advanced solution to the problem of decreasing the free mobility of the electrolyte in sealed batteries is its gel formation. By adding some 5-8 wt.% of pyrogenic silica to the electrolyte, a gel structure is formed due to the immense surface area (-200-300 m2 g ) of such silicas, which fixes the sulfuric acid solution molecules by van der Waals bonds within a lattice. These gels have thixotropic properties i.e., by mechanical stirring they can be liquefied and used to Filled into the... 

Although there are three Rji isotopes in the U- and Th-decay series, only is sufficiently long lived tm= 3.8 days) to be a useful estuarine tracer. Radioactive decay of Ra continuously produces Rn, which because of its short half-life is generally in secular equilibrium in seawater. Being chemically non-reactive except for very weak Van der Waals bonding makes this isotope a unique marine tracer in that it is not directly involved in biogeochemical cycles. 

T ike metals minerals also exhibit typical crystalline structures. As an example, the structure of molybdenite is shown in Figure 1.17. It is hexagonal with six-pole symmetry and contains two molecules per unit cell. Each sulfur atom is equidistant from three molybdenum atoms and each molybdenum atom is surrounded by six sulfur atoms located at the comers of a trigonal prism. There are two types of bonds that can be established between the atoms which constitute the molybdenite crystal stmcture. They are the covalent bonds between sulfur and molybdenum atoms and the Van der Waals bonds between sulfur-sulfur atoms. The Van der Waals bond is considerably weaker than the covalent sulfur-molybdenum bond. This causes the bonds of sulfur-sulfur to cleave easily, imparting to molybdenite the property of being a dry lubricant. Molybdenite adheres to metallic surfaces with the development of a molecular bond and the friction between metallic surfaces is replaced by easy friction between two layers of sulfur atoms. 

The presence of an immobile adsorbed film or layer on the particle surface may lead to the formation of still stronger interparticle van der Waals bonds (R6). First, surface roughness is smoothed out, increasing the apparent particle size and contact area, and second, the separation distance is effectively... 

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