Sobre las multas fijas

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

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H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Rh Db Sg Bh Hs Mt Uun Uuu Uub

Figure 8.1

States of elements under normal conditions. Light gray indicates gas, dark gray indicates liquids, all the rest are solids.

mole. One of the strongest and most impor-tant types of intermolecular forces is the hydrogen bond. Its strength can be as high as 40 kilojoules per mole. The hydrogen bond was introduced in Chapter 7. Let’s examine the hydrogen bond before looking at several other types of intermolecular attractions.

Hydrogen bonds are formed between molecules that contain hydrogen covalently bonded to atoms of oxygen, fluorine, or nitrogen. Hydrogen bonded to any one of these elements produces a highly polar molecule. The hydrogen takes on a partial positive charge, and the electronegative ele-ment has a partial negative charge. The hydrogen on one molecule has a high attrac-tion for the electronegative atom on an adja-cent molecule. It is important to remember that the hydrogen bond is an intermolecular force.

Water is the most common substance displaying hydrogen bonding. The hydrogen on one molecule is attracted to an unshared pair of electrons held by oxygen on an adja-cent water molecule (Figure 8.2). Because

each oxygen in water has two unshared pairs of electrons and there are also two hydrogen atoms in water, a network of water molecules forms, held together by hydrogen bonds.

Hydrogen bonding has a pronounced effect on the physical properties of water. For example, the boiling points of compounds Table 8.1

Characteristics of Solids, Liquids, and Gases

Figure 8.2

Hydrogen bonding in water. Hydrogen bonds are represented by dashed lines. Only several water molecules are shown and not all bonds are displayed. The pattern for the central oxygen atom continues to form a net-work of hydrogen-bonded water molecules.

generally decrease up a group in the periodic table. The reason for this decrease involves intermolecular forces and is explained shortly. In the oxygen group, the boiling points of H₂Te, H₂Se, H₂S, and H₂O are graphed in Figure 8.3.

The trend follows the expected trend for the first three compounds with the boiling point decreasing from the largest compound H₂Te to H₂S. According to the normal pat-tern, water should have the lowest boiling point. If the trend displayed for the first three compounds continued, then water’s boiling point should be lower than that of H₂S. The reason that water has an abnormally high boiling point compared to similar com-pounds is that water molecules exhibit hydrogen bonding. Because the water molecules are hydrogen bonded, it takes more energy to cause water to leave the liq-uid state and enter the gas state. It’s just as though you were with a group of friends, and someone tried to pulled you away from the group. It would be a lot harder to pull you, or any of your friends, away from the group if you were all holding on to each other. In essence, hydrogen bonding is how water molecules hold on to each other. A pattern similar to that shown in Figure 8.3 exists for Groups 15 and 17 hydrogen com-pounds with ammonia, NH₃and hydrogen fluoride, HF, having the highest boiling points in these two groups, respectively. This

is not surprising because both ammonia and hydrogen fluoride exhibit hydrogen bonding.

Another interesting property of water is that its maximum density is at 4°C. In gen-eral, almost all substances have a greater density in their solid state than in their liq-uid state. Hydrogen bonding is also respon-sible for this unique property of water.

Because each water molecule has two hydro-gen atoms and two lone pairs of electrons, the water molecules can form a three dimen-sional network of approximately tetrahe-drally bonded atoms. Each oxygen is covalently bonded to two hydrogen atoms and also hydrogen bonded to two oxygen atoms (Figure 8.2). When water exists as ice, the molecules form a rigid three-dimensional crystal. As the temperature of ice increases to the melting point of 0°C, hydrogen bond-ing provides enough attractive force to main-tain the approximate structure of ice. The increase in temperature above water’s melt-ing point causes thermal expansion, and this would normally lead to an increase in vol-ume and a corresponding decrease in den-sity. Remember, density is mass divided by volume so a larger volume results in a smaller density. The reason water actually becomes more dense is that at the melting point and up to 4°C there is enough energy for some of the water molecules to overcome the intermolecular attraction provided by the hydrogen bonding. These water molecules occupy void space in the still-present approximate crystalline structure. Because more molecules of water are occupying the same volume, the density increases. This phenomenon continues until the maximum density is reached at 4°C. At this point, the trapping of additional water molecules in the void space is not great enough to overcome the thermal expansion effect that lowers the density; therefore, solid ice is less dense than liquid water.

Figure 8.3

Trend in Boiling of Several Group 16 Compounds

The fact that the maximum density of water is at 4°C has significant environmen-tal impacts. Consider the freezing of a fresh-water lake in winter. As the temperature decreases, the surface water becomes pro-gressively denser and sinks to the bottom.

This process helps carry oxygen from the surface to deeper water. The cycling of a lake’s waters also helps to bring nutrients to the surface. This process is sometimes referred to as the fall overturn. Once the lake reaches a temperature of 4°C, the water is at its maximum density. Further cooling results in less dense water, which doesn’t sink, eventually forming a layer of ice if temperatures are cold enough. The ice layer effectively insulates the rest of the lake.

Because the water was replenished with oxygen during the fall overturn, the lake generally does not become anoxic. If water behaved like most substances, the entire lake would cool down to its freezing point, and then the entire lake would freeze solid.

Several other unique properties of water can be attributed to hydrogen bonding. We discuss these after we have had a chance to examine some of the other intermolecular forces. The hydrogen bond is actually an unusually strong form of what is known as a dipole-dipole force. Polar molecules give rise to dipole-dipole forces. A dipole-dipole force results from the electrostatic attraction between the partial positive and negative charges present in polar molecules. The strength of the dipole-dipole force is directly proportional to the strength of the dipole moment of the molecule.

Just as two polar molecules, like oppo-site ends of a magnet, are attracted to each other, a polar molecule may be attracted to an ion. This gives rise to an ion-dipole force. The negative ends of polar molecules are attracted to cations and the positive end to anions. The charge on the ion and the strength of the dipole moment determine the

strength of the ion-dipole force. When an ionic compound such as salt, NaCl, dis-solves in water, ion-dipole forces come into play. The positive hydrogen end of water and negative chloride ion are attracted to each other, while the negative oxygen end and positive sodium ion are attracted to each other (Figure 8.4). Polar molecules tend to be soluble in water, while nonpolar molecules are insoluble. This fact is illus-trated by a mixture of oil and water. The nonpolar oil does not dissolve in water, and two separate layers result. A similar effect occurs with the polar vinegar and nonpolar oil portions of a salad dressing.

Hydrogen bonding, dipole-dipole, and ion-dipole forces all involve polar mole-cules, and all involve electrostatic attrac-tions between opposite charges. A force that is present in both polar and nonpolar molecules is known as a dispersion or London force. The London force is named for Fritz London (1900–1954) who described this force in 1931. Dispersion forces are a consequence of the quantum nature of the atom. Electrons move around atomic nuclei in a probabilistic fashion, but they occupy specific energy levels. In non-polar molecules, on average, the movement of the electrons are such that the centers of positive and negative charge are the same.

It is important to realize, though, that this is an average state and does not represent what is occurring at any one instance.

Because the electrons are in constant motion, the electrons assume an asymmet-rical distribution producing an temporary dipole. An instant later, this situation Figure 8.4

Ion-Dipole Forces. M
represents a cation and X^an anion.

changes. Because electrons are constantly moving and changing position, temporary dipoles are continually being formed and dissolved (Figure 8.5). Temporary dipoles induce other temporary dipoles in adjacent atoms. That is, the negative and positive ends of the dipole will attract or repel electrons in adjacent atoms. This continual process leads to an ever-present attraction between nonpolar (as well as polar) molecules.

Dispersion forces depend on temporary dipoles inducing dipoles in adjacent atoms or molecules. The ease with which electrons in an atom or molecule can be distorted to form a temporary dipole is known as polar-izability. The larger the atom or molecule is, the greater its polarizability. This is because larger atoms have more electrons, and these electrons are located farther from their respective nuclei. Electrons farther from the nucleus are held less tightly, and this results in their greater polarizability.

This is why larger molecules in a series tend to have higher boiling points. For example, the boiling points of methane (CH₄), propane (C₃H₈), and butane (C₄H₁₀) are

161°C,
42°C, and 0°C, respectively.

An example of how dispersion forces and polarizability affect physical properties is seen in the halogens. Moving down the

halogen group fluorine and chlorine are gases, bromine is a liquid, and iodine a solid. This is just what would be expected.

Dispersion forces become stronger moving down the halogen group as the atoms increase in size and are more polarizable.

Another illustration is seen in Table 8.2 which lists the freezing point of the Nobel gases. The freezing point increases down the group. If the temperature of a mixture of the gases listed in the table was lowered, Xenon would freeze first because of the presence of greater dispersion forces.

Dipole-dipole, ion-dipole, and disper-sion forces are collectively known as van der Waals forces. Johannes Diederick van der Waals (1837–1923) received the 1910 Nobel Prize in physics for his work on flu-ids. We have seen how hydrogen bonding and van der Waals forces affect the physical properties of substances, and more is said about these forces as we examine the differ-ent states of matter.