Intermolecular forces in gases

What Do We Need to Know Already This chapter uses the concepts of potential energy (Section A), coulombic interactions (Section 2.4), polar molecules and dipoles (Section 3.3), and intermolecular forces in gases (Section 4.12). 

Molecular crystals are held together by van der Waals forces, the same as the intermolecular forces in gases and liquids. Because these are much weaker than ionic, metallic, and covalent bonds, the molecular crystals are usually soft, easily deformed, and have low melting points. 

Intermolecular forces in liquids are considerably stronger than intermolecular forces in gases. Particles are, on average. 

O. Sinanoglu, in Modern Quantum Chemistry—Istanbul Lectures, Part II, Interactions, O. Sinanoglu, Ed., Academic Press, New York, 1965, pp. 221-238. Intermolecular Forces in Gases and Dense Media. 

This chapter continues to explore states of matter by focusing on the gas state. Table 8.1 summarized the main features of gases. In the gaseous state, molecules are much farther apart than in either solid or liquids. Because of this distance molecules in the gaseous state have virtually no influence on each other. The independent nature of gas molecules means intermolecular forces in this state are minimal. A gas expands to fill the volume of its container. Most of the volume occupied by the gas consists of empty space. This characteristic allows gases to be compressible, and gases have only about 1/1,000 the density of solids and liquids. 

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. 

Just as the ideal gas forms a convenient point of reference in discussing the properties of real gases, so does the hard-sphere fluid in discussing the properties of liquids. This is especially true at low densities, where the role of intermolecular forces in real systems is not so important. In this limit, the hard-sphere model is useful in developing the theory of solutions, as will be seen in chapter 3. 

If intermolecular forces of attraction did not exist, most molecular substances would be gases at room temperature and pressure. There would be no attractive force between molecules causing them to cling together to form liquids or solids. You can determine the intermolecular forces in a substance using the following flowchart. 

LIQUIDS, OR GASES, BUT AMONG IONS AND MOLECULES DISSOLVED IN WATER OR OTHER SOLVENTS. In CHAPTERS 5 AND 11 WE LOOKED AT THE PROPERTIES OF GASES, LIQUIDS, AND SOLIDS. In THIS CHAPTER WE EXAMINE THE PROPERTIES OF SOLUTIONS, CONCENTRATING MAINLY ON THE ROLE OF INTERMOLECULAR FORCES IN SOLUBILITY AND OTHER PHYSICAL PROPERTIES OF SOLUTIONS. 

None of the three types of interatomic bonding discussed in the preceding chapters can be responsible for the cohesion of the inert gases in the solid state or for the intermolecular forces in organic crystals. In all of these cases the forces involved are those due to the residual or van der Waals bond. 

A rubber-like solid is unique in that its physical properties resemble those of solids, liquids, and gases in various respects. It is solidlike in that it maintains dimensional stability, and its elastic response at small strains (<5%) is essentially Hookean. It behaves like a liquid because its coefficient of thermal expansion and isothermal compressibility are of the same order of magnitude as those of liquids. The implication of this is that the intermolecular forces in an elastomer are similar to those in liquids. It resembles gases in the sense that the stress in a deformed elastomer increases with increasing temperature, much as the pressure in a compressed gas increases with increasing temperature. This gas-like behavior was, in fact, what first provided the hint that rubbery stresses are entropic in origin. 

Compare the strength of intermolecular forces in liquids with those in gases. 

As we learned in Chapter 10, the molecules in a gas cue widely separated and in a state of constant, chaotic motion. One of the key tenets of kinetic-molecular theory of gases is the assiunption that we can neglect the interactions between molecules. cco(Section 10.7) The properties of liquids and solids are quite different from those of gases largely because the intermolecular forces in Hquids and soUds are much stronger. A comparison of the properties of gases, liquids, and soUds is given in T Table 11.1. 

Most chemical reactions take place, not between pure solids, liquids, or gases, but among ions and molecules dissolved in water or other solvents. In Chapters 5 and 11 we looked at the properties of gases, liquids, and solids. In this chapter we examine the properties of solutions, concentrating mainly on the role of intermolecular forces in solubility and other physical properties of solution. 

The ideal gas is a fictitious model substance. The molecules are regarded as having no proprietary volume and exerting no intermolecular forces. In reality, there is no substance which fulfills these conditions, but the model of the ideal gas plays an important role as a starting point for the description of the PvT behavior of gases. Real substances behave very similar to ideal gases when the pressure approaches values of zero (v- - oo), because the molecular volume and the molecular interactions can be neglected at this state. 

A molecule on the surface of the crystal is held to the crystal by the forces of attraction that its neighboring molecules exert on it. Attractive forces of this kind, which are operative between all molecules when they are close together, are called van der Wools attractive forces this name is used because it was the Dutch physicist J. D. van der Waals (1837-1923) who first gave a thorough discussion of intermolecular forces in relation to the nature of gases and liquids. 

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. 

It follows from this discussion that all of the transport properties can be derived in principle from the simple kinetic dreoty of gases, and their interrelationship tlu ough k and c leads one to expect that they are all characterized by a relatively small temperature coefficient. The simple theory suggests tlrat this should be a dependence on 7 /, but because of intermolecular forces, the experimental results usually indicate a larger temperature dependence even up to for the case of molecular inter-diffusion. The Anhenius equation which would involve an enthalpy of activation does not apply because no activated state is involved in the transport processes. If, however, the temperature dependence of these processes is fitted to such an expression as an algebraic approximation, tlren an activation enthalpy of a few kilojoules is observed. It will thus be found that when tire kinetics of a gas-solid or liquid reaction depends upon the transport properties of the gas phase, the apparent activation entlralpy will be a few kilojoules only (less than 50 kJ). 

A useful property of liquids is their ability to dissolve gases, other liquids and solids. The solutions produced may be end-products, e.g. carbonated drinks, paints, disinfectants or the process itself may serve a useful function, e.g. pickling of metals, removal of pollutant gas from air by absorption (Chapter 17), leaching of a constituent from bulk solid. Clearly a solution s properties can differ significantly from the individual constituents. Solvents are covalent compounds in which molecules are much closer together than in a gas and the intermolecular forces are therefore relatively strong. When the molecules of a covalent solute are physically and chemically similar to those of a liquid solvent the intermolecular forces of each are the same and the solute and solvent will usually mix readily with each other. The quantity of solute in solvent is often expressed as a concentration, e.g. in grams/litre. 

In many cases, pressurized gases in vessels do not behave as ideal gases. At very high pressures, van der Waals forces become important, that is, intermolecular forces and finite molecule size influence the gas behavior. Another nonideal situation is that in which, following the rupture of a vessel containing both gas and liquid, the liquid flashes. 

---