Intermolecular forces in condensed

The origin of intermolecular forces in condensed phases is well known. However, the quantitative theoretical study of such intermolecular forces meets with great difficulties which are dependent on the fact that the distances between the molecules are of only the same order as the dimensions of the molecules themselves. The approximate formulae which are valid when the molecules are greater distances apart (say ten times the molecular diameter) as in a gas under normal conditions, lose all significance in a condensed phase, as we shall show later by an example. ... 

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. 

In this chapter, the basic types of chemical bonds existing in condensed phases are discussed. These interactions include ionic bonds, metallic bonds, covalent bonding (band theory), and intermolecular forces. In Chapter 10, the structures of some inorganic crystalline materials will be presented. 

The sudden expansion of liquid CO2 does work pushing back the atmosphere and overcoming intermolecular forces in the liquid. The energy to do that work comes from the molecules themselves, so the average energy of the molecules is lowered. The CO2 condenses to a solid because of this loss of energy. 

In the late 1800s and early 1900s, scientists were still struggling to understand intermolecular forces, so it is doubtful that Oscar Wilde had a clear picture of intermolecular forces in mind when he wrote of the subtle affinity between chemical atoms in The Picture of Dorian Gray. Nonetheless, his description of subtle affinity is quite apt. Intermolecular forces are complex, consisting of attractions as well as repulsions. Intermolecular attractions are those between water molecules that allow water to condense once it has been sufficiently cooled—and intermolecular repulsions are what make water feel like a solid mass when it is forcefully encountered. (Have you ever been knocked over by a wave ) If it were not for intermolecular attractions, our bodies would vaporize into gases, and if it were not for intermolecular repulsions, we would collapse into unimpressive puddles. 

What happens at the molecular level during evaporation In the beginning, the traffic is only one way Molecules are moving from the liqnid to the empty space. Soon the molecules in the space above the liquid estabhsh a vapor phase. As the concentration of molecules in the vapor phase increases, some molecules condense, that is, they return to the liquid phase. Condensation, the change from the gas phase to the liquid phase, occurs because a molecule strikes the liquid surface and becomes trapped by intermolecular forces in the liquid. 

Why does the process of vaporization require an input of energy Why is it so important that water has a large heat of vaporization What is condensation Explain how the processes of vaporization and condensation represent an equilibrium in a closed container. Define the equilibrium vapor pressure of a liquid. Describe how this pressure arises in a closed container. Describe an experiment that demonstrates vapor pressure and enables us to measure the magnitude of that pressure. How is the magnitude of a liquid s vapor pressure related to the intermolecular forces in the liquid ... 

Kinetics. The Role of Intermolecular Forces in the Condensed Phase Decomposition of Octahydro-1,3,5,7-tetranitro -1,3,5,7-tetrazocine Journal of Physical Chemistry 86, 4260-4265. 

The first follows from the improvements to the ideal-gas equation made by van der Waals in 1873, when he included a parameter 6 to account for the volume occupied by the molecules and a term a/V y to account for the intermolecular forces. In spite of its simplicity, the van der Waals equation gives a good qualitative representation of the behavior of a fluid, including the processes of condensation and evaporation, and... 

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. 

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". 

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. 

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