Some Thermodynamic Aspects of Intermolecular Forces. Interactions Involving Polar Molecules. Interactions Involving the Polarization of Molecules. Special Interactions. Hydrogen-Bonding, Hydrophobic, and Hydrophilic Interactions. The Forces between Particles and Surfaces. Contrasts between Intermolecular, Interparticle, and Intersurface Forces. Electrostatic Forces between Surfaces in Liquids. Solvation, Structural and Hydration Forces. Steric and Fluctuation Forces. These arrangements are more stable than arrangements in which two positive or two negative ends are adjacent part c in Figure 2.
A few examples are shown in Table 4. Krypton is a noble gas. E Fluctuating dipoles London Dispersion. The greater the number of accessible states, the greater the entropy. But when a single cation and a single anion are close together, within a protein, or within a folded RNA, those interactions are considered to be non-covalent electrostatic interactions.
Hence dipole—dipole interactions , such as those in part b in Figure 2. Because molecules in a liquid move freely and continuously, molecules always experience both attractive and repulsive dipole—dipole interactions simultaneously, as shown in Figure 2.
On average, however, the attractive interactions dominate. In addition, the attractive interaction between dipoles falls off much more rapidly with increasing distance than do the ion—ion interactions. Thus a substance such as HCl, which is partially held together by dipole—dipole interactions, is a gas at room temperature and 1 atm pressure, whereas NaCl, which is held together by interionic interactions, is a high-melting-point solid.
Within a series of compounds of similar molar mass, the strength of the intermolecular interactions increases as the dipole moment of the molecules increases, as shown in Table 2. Using what we learned about predicting relative bond polarities from the electronegativities of the bonded atoms, we can make educated guesses about the relative boiling points of similar molecules. Their structures are as follows:.
Asked for: order of increasing boiling points. Compare the molar masses and the polarities of the compounds. Compounds with higher molar masses and that are polar will have the highest boiling points. Show Answer.
Thus far we have considered only interactions between polar molecules, but other factors must be considered to explain why many nonpolar molecules, such as bromine, benzene, and hexane, are liquids at room temperature, and others, such as iodine and naphthalene, are solids. Even the noble gases can be liquefied or solidified at low temperatures, high pressures, or both Table What kind of attractive forces can exist between nonpolar molecules or atoms? This question was answered by Fritz London — , a German physicist who later worked in the United States.
In , London proposed that temporary fluctuations in the electron distributions within atoms and nonpolar molecules could result in the formation of short-lived instantaneous dipole moments , which produce attractive forces called London dispersion forces between otherwise nonpolar substances.
Consider a pair of adjacent He atoms, for example. On average, the two electrons in each He atom are uniformly distributed around the nucleus. Because the electrons are in constant motion, however, their distribution in one atom is likely to be asymmetrical at any given instant, resulting in an instantaneous dipole moment.
As shown in part a in Figure The net effect is that the first atom causes the temporary formation of a dipole, called an induced dipole , in the second. Interactions between these temporary dipoles cause atoms to be attracted to one another. These attractive interactions are weak and fall off rapidly with increasing distance.
Doubling the distance therefore decreases the attractive energy by 2 6 , or fold. Figure 3 Instantaneous Dipole Moments. The formation of an instantaneous dipole moment on one He atom a or an H2 molecule b results in the formation of an induced dipole on an adjacent atom or molecule. Instantaneous dipole—induced dipole interactions between nonpolar molecules can produce intermolecular attractions just as they produce interatomic attractions in monatomic substances like Xe. This effect, illustrated for two H 2 molecules in part b in Figure The reason for this trend is that the strength of London dispersion forces is related to the ease with which the electron distribution in a given atom can be perturbed.
In small atoms such as He, the two 1 s electrons are held close to the nucleus in a very small volume, and electron—electron repulsions are strong enough to prevent significant asymmetry in their distribution. In larger atoms such as Xe, however, the outer electrons are much less strongly attracted to the nucleus because of filled intervening shells. As a result, it is relatively easy to temporarily deform the electron distribution to generate an instantaneous or induced dipole. The ease of deformation of the electron distribution in an atom or molecule is called its polarizability.
Because the electron distribution is more easily perturbed in large, heavy species than in small, light species, we say that heavier substances tend to be much more polarizable than lighter ones. Note: For similar substances, London dispersion forces get stronger with increasing molecular size. The polarizability of a substance also determines how it interacts with ions and species that possess permanent dipoles.
Thus London dispersion forces are responsible for the general trend toward higher boiling points with increased molecular mass and greater surface area in a homologous series of compounds, such as the alkanes part a in Figure 2. The strengths of London dispersion forces also depend significantly on molecular shape because shape determines how much of one molecule can interact with its neighboring molecules at any given time.
For example, part b in Figure 2. Neopentane is almost spherical, with a small surface area for intermolecular interactions, whereas n -pentane has an extended conformation that enables it to come into close contact with other n -pentane molecules. As a result, the boiling point of neopentane 9. As a result, neopentane is a gas at room temperature, whereas n-pentane is a volatile liquid. All molecules, whether polar or nonpolar, are attracted to one another by London dispersion forces in addition to any other attractive forces that may be present.
In general, however, dipole—dipole interactions in small polar molecules are significantly stronger than London dispersion forces, so the former predominate. Arrange n -butane, propane, 2-methylpropane [isobutene, CH 3 2 CHCH 3 ], and n -pentane in order of increasing boiling points.
Determine the intermolecular forces in the compounds and then arrange the compounds according to the strength of those forces. The substance with the weakest forces will have the lowest boiling point. The four compounds are alkanes and nonpolar, so London dispersion forces are the only important intermolecular forces. These forces are generally stronger with increasing molecular mass, so propane should have the lowest boiling point and n -pentane should have the highest, with the two butane isomers falling in between.
Of the two butane isomers, 2-methylpropane is more compact, and n -butane has the more extended shape. Consequently, we expect intermolecular interactions for n -butane to be stronger due to its larger surface area, resulting in a higher boiling point. Molecules with hydrogen atoms bonded to electronegative atoms such as O, N, and F and to a much lesser extent Cl and S tend to exhibit unusually strong intermolecular interactions. These result in much higher boiling points than are observed for substances in which London dispersion forces dominate, as illustrated for the covalent hydrides of elements of groups 14—17 in Figure 2.
Methane and its heavier congeners in group 14 form a series whose boiling points increase smoothly with increasing molar mass. This is the expected trend in nonpolar molecules, for which London dispersion forces are the exclusive intermolecular forces. These plots of the boiling points of the covalent hydrides of the elements of groups 14—17 show that the boiling points of the lightest members of each series for which hydrogen bonding is possible HF, NH3, and H2O are anomalously high for compounds with such low molecular masses.
Why do strong intermolecular forces produce such anomalously high boiling points and other unusual properties, such as high enthalpies of vaporization and high melting points? The answer lies in the highly polar nature of the bonds between hydrogen and very electronegative elements such as O, N, and F. The large difference in electronegativity results in a large partial positive charge on hydrogen and a correspondingly large partial negative charge on the O, N, or F atom. Because a hydrogen atom is so small, these dipoles can also approach one another more closely than most other dipoles.
The combination of large bond dipoles and short dipole—dipole distances results in very strong dipole—dipole interactions called hydrogen bonds , as shown for ice in Figure 2. A hydrogen bond is usually indicated by a dotted line between the hydrogen atom attached to O, N, or F the hydrogen bond donor and the atom that has the lone pair of electrons the hydrogen bond acceptor.
Because each water molecule contains two hydrogen atoms and two lone pairs, a tetrahedral arrangement maximizes the number of hydrogen bonds that can be formed.
mta-sts.builttospill.reclaimhosting.com/cyberhome-dvd-manual-de-jugador.php In the structure of ice, each oxygen atom is surrounded by a distorted tetrahedron of hydrogen atoms that form bridges to the oxygen atoms of adjacent water molecules. The bridging hydrogen atoms are not equidistant from the two oxygen atoms they connect, however. Instead, each hydrogen atom is pm from one oxygen and pm from the other.
The resulting open, cagelike structure of ice means that the solid is actually slightly less dense than the liquid, which explains why ice floats on water rather than sinks. Each water molecule accepts two hydrogen bonds from two other water molecules and donates two hydrogen atoms to form hydrogen bonds with two more water molecules, producing an open, cagelike structure.
The structure of liquid water is very similar, but in the liquid, the hydrogen bonds are continually broken and formed because of rapid molecular motion. Note: Hydrogen bond formation requires both a hydrogen bond donor and a hydrogen bond acceptor. Because ice is less dense than liquid water, rivers, lakes, and oceans freeze from the top down. In fact, the ice forms a protective surface layer that insulates the rest of the water, allowing fish and other organisms to survive in the lower levels of a frozen lake or sea. If ice were denser than the liquid, the ice formed at the surface in cold weather would sink as fast as it formed.
Bodies of water would freeze from the bottom up, which would be lethal for most aquatic creatures. Compounds such as HF can form only two hydrogen bonds at a time as can, on average, pure liquid NH 3. Consequently, even though their molecular masses are similar to that of water, their boiling points are significantly lower than the boiling point of water, which forms four hydrogen bonds at a time.
Draw the hydrogen-bonded structures. Asked for: formation of hydrogen bonds and structure. A Of the species listed, xenon Xe , ethane C 2 H 6 , and trimethylamine [ CH 3 3 N] do not contain a hydrogen atom attached to O, N, or F; hence they cannot act as hydrogen bond donors. B The one compound that can act as a hydrogen bond donor, methanol CH 3 OH , contains both a hydrogen atom attached to O making it a hydrogen bond donor and two lone pairs of electrons on O making it a hydrogen bond acceptor ; methanol can thus form hydrogen bonds by acting as either a hydrogen bond donor or a hydrogen bond acceptor.
The hydrogen-bonded structure of methanol is as follows:.
Identify the intermolecular forces in each compound and then arrange the compounds according to the strength of those forces. Molecules in liquids are held to other molecules by intermolecular interactions, which are weaker than the intramolecular interactions that hold the atoms together within molecules and polyatomic ions. Transitions between the solid and liquid or the liquid and gas phases are due to changes in intermolecular interactions but do not affect intramolecular interactions. The three major types of intermolecular interactions are dipole—dipole interactions, London dispersion forces these two are often referred to collectively as van der Waals forces , and hydrogen bonds.
London dispersion forces are due to the formation of instantaneous dipole moments in polar or nonpolar molecules as a result of short-lived fluctuations of electron charge distribution, which in turn cause the temporary formation of an induced dipole in adjacent molecules. Larger atoms tend to be more polarizable than smaller ones because their outer electrons are less tightly bound and are therefore more easily perturbed.