Intermolecular forces condensation

As also noted in the preceding chapter, it is customary to divide adsorption into two broad classes, namely, physical adsorption and chemisorption. Physical adsorption equilibrium is very rapid in attainment (except when limited by mass transport rates in the gas phase or within a porous adsorbent) and is reversible, the adsorbate being removable without change by lowering the pressure (there may be hysteresis in the case of a porous solid). It is supposed that this type of adsorption occurs as a result of the same type of relatively nonspecific intermolecular forces that are responsible for the condensation of a vapor to a liquid, and in physical adsorption the heat of adsorption should be in the range of heats of condensation. Physical adsorption is usually important only for gases below their critical temperature, that is, for vapors. 

We have seen that the pure elements may solidify in the form of molecular solids, network solids, or metals. Compounds also may condense to molecular solids, network solids, or metallic solids. In addition, there is a new effect that does not occur with the pure elements. In a pure element the ionization energies of all atoms are identical and electrons are shared equally. In compounds, where the most stable electron distribution need not involve equal sharing, electric dipoles may result. Since two bonded atoms may have different ionization energies, the electrons may spend more time near one of the positive nuclei than near the other. This charge separation may give rise to strong intermolecular forces of a type not found in the pure elements. 

Intermolecular forces are responsible for the existence of several different phases of matter. A phase is a form of matter that is uniform throughout in both chemical composition and physical state. The phases of matter include the three common physical states, solid, liquid, and gas (or vapor), introduced in Section A. Many substances have more than one solid phase, with different arrangements of their atoms or molecules. For instance, carbon has several solid phases one is the hard, brilliantly transparent diamond we value and treasure and another is the soft, slippery, black graphite we use in common pencil lead. A condensed phase means simply a solid or liquid phase. The temperature at which a gas condenses to a liquid or a solid depends on the strength of the attractive forces between its molecules. 

A 7/vap always is 3.10 kJ/mol greater than A fi vap At 298 K (25 °C, room temperature) A /Tvap always is 2.48 kJ/mol greater than A ivap The difference between A vap and A i7vap arises because, in addition to overcoming intermolecular forces in the condensed phase (A E), the escaping vapor must do work, w = A(P V ) — RT as it expands against the constant external pressure of the atmosphere. 

Gases and condensed phases look very different at the molecular level. Molecules of F2 or CI2 move freely throughout their gaseous volume, traveling many molecular diameters before colliding with one another or with the walls of their container. Because much of the volume of a gas is empty space, samples of gaseous F2 and CI2 readily expand or contract in response to changes in pressure. This freedom of motion exists because the intermolecular forces between these molecules are small. 

A substance exists in a condensed phase when its molecules have too little average kinetic energy to overcome intermolecular forces of attraction. 

A gas condenses to a liquid if it is cooled sufficiently. Condensation occurs when the average kinetic energy of motion of molecules falls below the value needed for the molecules to move about independently. Thus, the molecules in a liquid are confined to a specific volume by intermolecular forces of attraction. Although they cannot readily escape, liquid molecules remain free to move about within the liquid phase, hi this behavior, liquid molecules behave like the molecules of a gas. The large-scale consequences of the molecular-level properties are apparent. Like gases, liquids are fluid, so they flow easily from place to place. Unlike gases, however, liquids are compact, so they cannot expand or contract significantly. 

Molecular solids are aggregates of molecules bound together by intermolecular forces. Substances that are gases under normal conditions form molecular solids when they condense at low temperature. Many larger molecules have sufficient dispersion forces to exist as solids at room temperature. One example is naphthalene (Cio Hg), a white solid that melts at 80 °C. Naphthalene has a planar structure like that of benzene (see Section 10-), with a cloud of ten delocalized n electrons that lie above and below the molecular plane. Naphthalene molecules are held in the solid state by strong dispersion forces among these highly polarizable n electrons. The molecules in... 

Phase changes are characteristic of all substances. The normal phases displayed by the halogens appear in Section II-L where we also show that a gas liquefies or a liquid freezes at low enough temperatures. Vapor pressure, which results from molecules escaping from a condensed phase into the gas phase, is one of the liquid properties described in Section II-I. Phase changes depends on temperature, pressure, and the magnitudes of intermolecular forces. 

In the liquid state, the molecules are still free to move in three dimensions but stiU have to be confined in a container in the same manner as the gaseous state if we expect to be able to measure them. However, there are important differences. Since the molecules in the liquid state have had energy removed from them in order to get them to condense, the translational degrees of freedom are found to be restricted. This is due to the fact that the molecules are much closer together and can interact with one another. It is this interaction that gives the Uquid state its unique properties. Thus, the molecules of a liquid are not free to flow in any of the three directions, but are bound by intermolecular forces. These forces depend upon the electronic structure of the molecule. In the case of water, which has two electrons on the ojQ gen atom which do not participate in the bonding structure, the molecule has an electronic moment, i.e.- is a "dipole". 

Intermolecular forces can affect phase changes. Strong intermolecular forces require more kinetic energy to convert a liquid into a gas. Stronger intermolecular forces, make it easier to condense a gas into a liquid. 

The liquid state is a condensed state, so each molecule is always interacting with a group of neighbours although diffusing quite rapidly. As a result, although momentum through a shear plane still occurs, it is a small contribution when compared to the frictional resistance of the molecules in adjacent layers. It is the nature of this frictional resistance that we must now address and it will become clear that it arises from the intermolecular forces. The theories of the viscosity of liquids are still in an unfinished state but the physical ideas have been laid down. The first... 

The cohesive energy coh of a substance in a condensed state is defined as the increase in internal energy AU per mole of substance if all the intermolecular forces are eliminated. 

Our interest is in the connection between the intermolecular forces that cause condensation and/or gas phase molecular clustering and thermodynamics. To set the stage consider the following simple model ... 

The previous chapter dealt with chemical bonding and the forces present between the atoms in molecules. Forces between atoms within a molecule are termed intramolecular forces and are responsible for chemical bonding. The interaction of valence electrons between atoms creates intramolecular forces, and this interaction dictates the chemical behavior of substances. Forces also exist between the molecules themselves, and these are collectively referred to as intermolecular forces. Intermolecular forces are mainly responsible for the physical characteristics of substances. One of the most obvious physical characteristics related to intermolecular force is the phase or physical state of matter. Solid, liquid, and gas are the three common states of matter. In addition to these three, two other states of matter exist—plasma and Bose-Einstein condensate. 

Intermolecular forces are responsible for the condensed states of matter. The particles making up solids and liquids are held together by intermolecular forces, and these forces affect a number of the physical properties of matter in these two states. Intermolecular forces are quite a bit weaker than the covalent and ionic bonds discussed in Chapter 7. The latter requires several hundred to several thousand kilojoules per mole to break. The strength of intermolecular forces are a few to tens of kilojoules per... 

Throughout this chapter we have dealt with surface tension from a phenomenological point of view almost exclusively. From fundamental perspective, however, descriptions from a molecular perspective are often more illuminating than descriptions of phenomena alone. In condensed phases, in which interactions involve many molecules, rigorous derivations based on the cumulative behavior of individual molecules are extremely difficult. We shall not attempt to review any of the efforts directed along these lines for surface tension. Instead, we consider the various types of intermolecular forces that exist and interpret 7 for any interface as the summation of contributions arising from the various types of interactions that operate in the materials forming the interface. 

If the area of an insoluble monolayer is isothermally reduced still further, the compressibility eventually becomes very low. Because of the low compressibility, the states observed at these low values of a are called condensed states. In general, the isotherm is essentially linear, although it may display a well-defined change in slope as tt is increased, as shown in Figure 7.6. As menlioned above, the (relatively) more expanded of these two linear portions is the liquid-condensed state LC, and the less expanded is the solid state S. It is clear from the low compressibility of these states that both the LC and S states are held together by strong intermolecular forces so as to be relatively independent of the film pressure. 

The very fact that the vapor phase of many substances can condense to form a liquid is a consequence of the existence of attractive van der Waals forces between atoms or molecules. An attractive intermolecular force is not needed for a gas to condense into a solid solidification can occur purely as a result of excluded-volume interactions among the molecules at sufficiently large densities. The pressure in a fluid, the cohesion between materials, and the existence of surface energy or surface tension all result, partially or wholly, from van der Waals forces. 

Primitive considerations convince us that such secondary forces exist. For example, gases consist of disordered molecules, whether they be polyatomic like chlorine or ether vapour, or single atoms like helium or mercury vapour. But all gases, even helium, ultimately condense to liquids — and then to solids — if they are cooled and/or compressed sufficiently. When the molecules are forced into close proximity and have their kinetic energies diminished, the weak intermolecular forces are able to take control. Liquefaction results. The strength of these forces can be measured by the latent heat necessary to evaporate the liquid, or to sublime the solid. The equation,... 

The temperature at which a gas condenses depends on the pressure and the strength of the attractive forces between its molecules. Intermolecular forces pull molecules together and, provided the temperature is low enough, produce a condensed phase. In the gas phase, the properties of the substance are dominated by the nearly free motion of the molecules, so they are nearly independent of the identity of the gas. In condensed phases, however, the molecules are very close to one another all the time and intermolecular forces are of dominating importance. 

Physisorption (or Physical Adsorption) is adsorption in which the forces involved are intermolecular forces (van der Waals forces) of the same kind as those responsible for the imperfection of real gases and the condensation of vapours, and which do not involve a significant change in the electronic orbital patterns of the species involved. The term van der Waals adsorption is synonymous with physical adsorption, but its use is not recommended. 

A supercritical fluid exists when a substance is heated above its critical temperature and pressure and is unable to be condensed to a liquid by pressure alone. A typical supercritical fluid is carbon dioxide, which, at temperatures above 31°C and pressures above 73 atm, exists in a supercritical fluid state where individual molecules of the compound are held by less restrictive intermolecular forces and molecular movement resembles that of a gas (1 atm =101 325 Pa). 

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