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In Chapter 10 "Gases", you learned that attractive intermolecular forces cause most gases to condense to liquids at high pressure, low temperature, or both. In this chapter, we look more closely at the intermolecular forces that are responsible for the properties of liquids, describe some of the unique properties of liquids compared with the other states of matter, and then consider changes in state between liquids and gases or solids.
The kinetic molecular theory of gases described in Chapter 10 "Gases" gives a reasonably accurate description of the behavior of gases.
In a gas, the distance between molecules, whether monatomic or polyatomic, is very large compared with the size of the molecules; thus gases have a low density and are highly compressible. Liquids exhibit short-range order because strong intermolecular attractive forces cause the molecules to pack together rather tightly.
Liquids have so little empty space between their component molecules that they cannot be readily compressed.
The intermolecular forces in liquids are strong enough to keep them from expanding significantly when heated (typically only a few percent over a 100A°C temperature range). Molecules in liquids diffuse because they are in constant motion (Figure 11.1 "Molecular Diffusion in a Liquid").
A drop of an aqueous solution containing a marker dye is added to a larger volume of water. Liquids can flow, adjusting to the shape of their containers, because their molecules are free to move. The properties of liquids can be explained using a modified version of the kinetic molecular theory of gases described in Chapter 10 "Gases". The kinetic molecular description of liquids must take into account both the nonzero volumes of particles and the presence of strong intermolecular attractive forces. How must the ideal gas law be altered to apply the kinetic molecular theory of gases to liquids?
Why are the root mean square speeds of molecules in liquids less than the root mean square speeds of molecules in gases? The properties of liquids are intermediate between those of gases and solids but are more similar to solids.
Intermolecular forces determine bulk properties such as the melting points of solids and the boiling points of liquids. Intermolecular forces are electrostatic in nature; that is, they arise from the interaction between positively and negatively charged species.
Because each end of a dipole possesses only a fraction of the charge of an electron, dipolea€“dipole interactions are substantially weaker than the interactions between two ions, each of which has a charge of at least A±1, or between a dipole and an ion, in which one of the species has at least a full positive or negative charge.
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. 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 dipolea€“induced dipole interactions between nonpolar molecules can produce intermolecular attractions just as they produce interatomic attractions in monatomic substances like Xe. 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, as we shall see when we discuss solutions in Chapter 13 "Solutions".
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.
Arrange n-butane, propane, 2-methylpropane [isobutene, (CH3)2CHCH3], 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 four compounds are alkanes and nonpolar, so London dispersion forces are the only important intermolecular forces. 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 plots of the boiling points of the covalent hydrides of the elements of groups 14a€“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? 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.
Because ice is less dense than liquid water, rivers, lakes, and oceans freeze from the top down. B Of the compounds that can act as hydrogen bond donors, identify those that also contain lone pairs of electrons, which allow them to be hydrogen bond acceptors. A Of the species listed, xenon (Xe), ethane (C2H6), and trimethylamine [(CH3)3N] 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 (CH3OH), 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. Arrange C60 (buckminsterfullerene, which has a cage structure), NaCl, He, Ar, and N2O in order of increasing boiling points. Identify the intermolecular forces in each compound and then arrange the compounds according to the strength of those forces. Electrostatic interactions are strongest for an ionic compound, so we expect NaCl to have the highest boiling point. 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. Intermolecular forces are electrostatic in nature and include van der Waals forces and hydrogen bonds.
What is the main difference between intramolecular interactions and intermolecular interactions? Describe the three major kinds of intermolecular interactions discussed in this chapter and their major features. Explain why hydrogen bonds are unusually strong compared to other dipolea€“dipole interactions. Liquid water is essential for life as we know it, but based on its molecular mass, water should be a gas under standard conditions. Describe the effect of polarity, molecular mass, and hydrogen bonding on the melting point and boiling point of a substance. Using acetic acid as an example, illustrate both attractive and repulsive intermolecular interactions. In group 17, elemental fluorine and chlorine are gases, whereas bromine is a liquid and iodine is a solid. Both water and methanol have anomalously high boiling points due to hydrogen bonding, but the boiling point of water is greater than that of methanol despite its lower molecular mass. Why is it not advisable to freeze a sealed glass bottle that is completely filled with water?
Use the melting of a metal such as lead to explain the process of melting in terms of what is happening at the molecular level. How does the Oa€“H distance in a hydrogen bond in liquid water compare with the Oa€“H distance in the covalent Oa€“H bond in the H2O molecule? Explain why the hydrogen bonds in liquid HF are stronger than the corresponding intermolecular Ha‹…a‹…a‹…I interactions in liquid HI. For which substance will hydrogen bonding have the greater effect on the boiling point: HF or H2O? Water is a liquid under standard conditions because of its unique ability to form four strong hydrogen bonds per molecule. As the atomic mass of the halogens increases, so does the number of electrons and the average distance of those electrons from the nucleus. The V-shaped SO2 molecule has a large dipole moment due to the polar S=O bonds, so dipolea€“dipole interactions will be most important. The Ha€“F bond is highly polar, and the fluorine atom has three lone pairs of electrons to act as hydrogen bond acceptors; hydrogen bonding will be most important. Although the C=O bonds are polar, this linear molecule has no net dipole moment; hence, London dispersion forces are most important.
This is a symmetrical molecule that has no net dipole moment, and the Cl atoms are relatively polarizable; thus, London dispersion forces will dominate. Water has two polar Oa€“H bonds with H atoms that can act as hydrogen bond donors, plus two lone pairs of electrons that can act as hydrogen bond acceptors, giving a net of four hydrogen bonds per H2O molecule.
Vigorous boiling causes more water molecule to escape into the vapor phase, but does not affect the temperature of the liquid.
Although you have been introduced to some of the interactions that hold molecules together in a liquid, we have not yet discussed the consequences of those interactions for the bulk properties of liquids. We stated in Section 11.1 "The Kinetic Molecular Description of Liquids" that liquids tend to adopt the shapes of their containers. Figure 11.9 "A Representation of Surface Tension in a Liquid" presents a microscopic view of a liquid droplet.
Molecules at the surface of water experience a net attraction to other molecules in the liquid, which holds the surface of the bulk sample together. The same phenomenon holds molecules together at the surface of a bulk sample of water, almost as if they formed a skin.
Such phenomena are manifestations of surface tensionThe energy required to increase the surface area of a liquid by a certain amount.
The stopwatch does not get stopped until that blue color returns (you are about to make it go away)9. Increasing the temperature increases the kinetic energy of the enzyme and substrate molecules so that they move faster and are more likely to collide. This website uses essential cookies without which it will not work, along with other harmless cookies aimed at improving your use of our website.

Conventional wind turbines use massive turbine generator systems mounted on top of a tower. Bringing the airflow from the top of the tower to ground level allows for greater power generation with much smaller turbine blades. The unit is about 50% shorter than traditional wind towers and uses a ground-based turbine with blades that are 84% smaller. INVELOX is scalable, equally suitable for use in major wind farms or micro-generation settings. Invelox triggered a massive eight out of thirteen possible red flags on my assessment of whether they are good or bad bets for investment. The Invelox is a very ineffective and very misleadingly represented wind generation device. See the license for more details, but that basically means you can share this book as long as you credit the author (but see below), don't make money from it, and do make it available to everyone else under the same terms. However, the publisher has asked for the customary Creative Commons attribution to the original publisher, authors, title, and book URI to be removed. Substances that normally are liquids are held together by exactly the same forces that are responsible for the liquefaction of gases.
By the end of the chapter, you will understand what is happening at the molecular level when you dry yourself with a towel, why you feel cold when you come out of the water, why ice is slippery, and how it is possible to decaffeinate coffee without removing important flavor components.
A similar model can be applied to liquids, but it must take into account the nonzero volumes of particles and the presence of strong intermolecular attractive forces. In contrast, the molecules in liquids are very close together, with essentially no empty space between them. Because of their higher kinetic energy compared to the molecules in a solid, however, the molecules in a liquid move rapidly with respect to one another. Compression would force the atoms on adjacent molecules to occupy the same region of space. A molecule in a liquid cannot move far before colliding with another molecule, however, so the mean free path in liquids is very short, and the rate of diffusion is much slower than in gases.
This freedom of motion and their close spacing allow the molecules in a liquid to move rapidly into the openings left by other molecules, in turn generating more openings, and so forth (Figure 11.2 "Why Liquids Flow"). Consequently, when the flask is tilted, molecules move to the left and down due to the force of gravity, and the openings are occupied by other molecules. In contrast to intramolecular forces, such as the covalent bonds that hold atoms together in molecules and polyatomic ions, intermolecular forces hold molecules together in a liquid or solid. Liquids boil when the molecules have enough thermal energy to overcome the intermolecular attractive forces that hold them together, thereby forming bubbles of vapor within the liquid. Like covalent and ionic bonds, intermolecular interactions are the sum of both attractive and repulsive components.
If the structure of a molecule is such that the individual bond dipoles do not cancel one another, then the molecule has a net dipole moment. In addition, the attractive interaction between dipoles falls off much more rapidly with increasing distance than do the iona€“ion interactions we considered in Chapter 8 "Ionic versus Covalent Bonding".
Compounds with higher molar masses and that are polar will have the highest boiling points. The first compound, 2-methylpropane, contains only Ca€“H bonds, which are not very polar because C and H have similar electronegativities. Even the noble gases can be liquefied or solidified at low temperatures, high pressures, or both (Table 11.3 "Normal Melting and Boiling Points of Some Elements and Nonpolar Compounds"). This question was answered by Fritz London (1900a€“1954), a German physicist who later worked in the United States. On average, the two electrons in each He atom are uniformly distributed around the nucleus.
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 11.6 "Mass and Surface Area Affect the Strength of London Dispersion Forces").
In general, however, dipolea€“dipole interactions in small polar molecules are significantly stronger than London dispersion forces, so the former predominate.
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. 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 14a€“17 in Figure 11.7 "The Effects of Hydrogen Bonding on Boiling 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 structure of liquid water is very similar, but in the liquid, the hydrogen bonds are continually broken and formed because of rapid molecular motion. 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. Compounds such as HF can form only two hydrogen bonds at a time as can, on average, pure liquid NH3. If a substance is both a hydrogen donor and a hydrogen bond acceptor, draw a structure showing the hydrogen bonding.
To predict the relative boiling points of the other compounds, we must consider their polarity (for dipolea€“dipole interactions), their ability to form hydrogen bonds, and their molar mass (for London dispersion forces). 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. How does the boiling point of a substance depend on the magnitude of the repulsive intermolecular interactions? What is the difference in the temperature of the cooking liquid between boiling and simmering?
As a piece of lead melts, the temperature of the metal remains constant, even though energy is being added continuously. Larger atoms with more electrons are more easily polarized than smaller atoms, and the increase in polarizability with atomic number increases the strength of London dispersion forces. In such a case, dipolea€“dipole interactions and London dispersion forces are often comparable in magnitude. Although methanol also has two lone pairs of electrons on oxygen that can act as hydrogen bond acceptors, it only has one Oa€“H bond with an H atom that can act as a hydrogen bond donor. We now turn our attention to three unique properties of liquids that intimately depend on the nature of intermolecular interactions: surface tension, capillary action, and viscosity. Why, then, do small amounts of water on a freshly waxed car form raised droplets instead of a thin, continuous film? A typical molecule in the interior of the droplet is surrounded by other molecules that exert attractive forces from all directions. When filling a glass with water, the glass can be overfilled so that the level of the liquid actually extends above the rim.
In a dynamic equilibrium the position of the equilibrium will shift in order to reverse any changes you introduce.
So increasing the temperature increases the rate of the reaction up to a certain temperature. It also allows for networking, allowing multiple towers to direct energy to the same generator. It will produce power effectively at much lower wind speeds than current systems do, which means it will have a dramatically wider geographical range of use. You may also download a PDF copy of this book (147 MB) or just this chapter (6 MB), suitable for printing or most e-readers, or a .zip file containing this book's HTML files (for use in a web browser offline). You will also learn how liquid crystal display (LCD) devices in electronic devices function, and how adhesive strips used to measure body temperature change color to indicate a fever. As in gases, however, the molecules in liquids are in constant motion, and their kinetic energy (and hence their speed) depends on their temperature. The density of a liquid is typically about the same as the density of the solid state of the substance. Thus unlike the ions in the ionic solids discussed in Chapter 8 "Ionic versus Covalent Bonding", Section 8.2 "Ionic Bonding", the molecules in liquids are not arranged in a repeating three-dimensional array. Notice from Table 11.1 "The Density of Water at Various Temperatures" that the density of water, for example, changes by only about 3% over a 90-degree temperature range. Similarly, solids melt when the molecules acquire enough thermal energy to overcome the intermolecular forces that lock them into place in the solid.
Because electrostatic interactions fall off rapidly with increasing distance between molecules, intermolecular interactions are most important for solids and liquids, where the molecules are close together.
For more information on the dissolution of ionic substances, see Chapter 4 "Reactions in Aqueous Solution" and Chapter 5 "Energy Changes in Chemical Reactions".) dipolea€“dipole interactions, London dispersion forces, and hydrogen bonds.
Molecules with net dipole moments tend to align themselves so that the positive end of one dipole is near the negative end of another and vice versa, as shown in part (a) in Figure 11.3 "Attractive and Repulsive Dipolea€“Dipole Interactions".
It should therefore have a very small (but nonzero) dipole moment and a very low boiling point. In 1930, London proposed that temporary fluctuations in the electron distributions within atoms and nonpolar molecules could result in the formation of short-lived instantaneous dipole momentsThe short-lived dipole moment in atoms and nonpolar molecules caused by the constant motion of their electrons, which results in an asymmetrical distribution of charge at any given instant., which produce attractive forces called London dispersion forcesA kind of intermolecular interaction (force) that results from temporary fluctuations in the electron distribution within atoms and nonpolar molecules. 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. 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. As a result, neopentane is a gas at room temperature, whereas n-pentane is a volatile liquid.
Of the two butane isomers, 2-methylpropane is more compact, and n-butane has the more extended shape. Methane and its heavier congeners in group 14 form a series whose boiling points increase smoothly with increasing molar mass. 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. If ice were denser than the liquid, the ice formed at the surface in cold weather would sink as fast as it formed.
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.
The three major types of intermolecular interactions are dipolea€“dipole interactions, London dispersion forces (these two are often referred to collectively as van der Waals forces), and hydrogen bonds. Identify the kind of interaction that includes hydrogen bonds and explain why hydrogen bonds fall into this category.

Draw the structures of these two compounds, including any lone pairs, and indicate potential hydrogen bonds.
These intermolecular interactions are strong enough to favor the condensed states for bromine and iodine under normal conditions of temperature and pressure. Consequently, methanol can only form two hydrogen bonds per molecule on average, versus four for water. The answer lies in a property called surface tension, which depends on intermolecular forces.
Consequently, there is no net force on the molecule that would cause it to move in a particular direction. Similarly, a sewing needle or a paper clip can be placed on the surface of a glass of water where it a€?floats,a€? even though steel is much denser than water (part (a) in Figure 11.10 "The Effects of the High Surface Tension of Liquid Water"). As C & D are formed in increasing amounts what will happen to the rate of the reverse reaction ? The sum of the partial pressures of the gases in a mixture is equal to the total pressure of the mixture.) If K c or K p are large the amount of products will be large and reactants small, ? we say the equilibrium position is over to the right hand side (RHS).
It is not “how much” of a product is made, but instead “how quickly” a reaction takes place. Instead of snatching bits of energy from the wind as it passes through the blades of a rotor, wind is captured with a funnel and directed through a tapering passageway that naturally accelerates its flow. Because of its thermal properties, water also modulates Eartha€™s temperature, maintaining a temperature range suitable for life.
We begin our discussion by examining some of the characteristic properties of liquids to see how each is consistent with a modified kinetic molecular description. Unlike the molecules in gases, however, the arrangement of the molecules in a liquid is not completely random. These interactions become important for gases only at very high pressures, where they are responsible for the observed deviations from the ideal gas law at high pressures. The first two are often described collectively as van der Waals forcesThe intermolecular forces known as dipolea€“dipole interactions and London dispersion forces.. These arrangements are more stable than arrangements in which two positive or two negative ends are adjacent (part (c) in Figure 11.3 "Attractive and Repulsive Dipolea€“Dipole Interactions").
Ethyl methyl ether has a structure similar to H2O; it contains two polar Ca€“O single bonds oriented at about a 109A° angle to each other, in addition to relatively nonpolar Ca€“H bonds.
As shown in part (a) in Figure 11.5 "Instantaneous Dipole Moments", the instantaneous dipole moment on one atom can interact with the electrons in an adjacent atom, pulling them toward the positive end of the instantaneous dipole or repelling them from the negative end.
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.
For example, part (b) in Figure 11.6 "Mass and Surface Area Affect the Strength of London Dispersion Forces" shows 2,2-dimethylpropane (neopentane) and n-pentane, both of which have the empirical formula C5H12. Consequently, we expect intermolecular interactions for n-butane to be stronger due to its larger surface area, resulting in a higher boiling point.
This is the expected trend in nonpolar molecules, for which London dispersion forces are the exclusive intermolecular forces.
Consequently, Ha€“O, Ha€“N, and Ha€“F bonds have very large bond dipoles that can interact strongly with one another. Bodies of water would freeze from the bottom up, which would be lethal for most aquatic creatures. In contrast, a molecule on the surface experiences a net attraction toward the drop because there are no molecules on the outside to balance the forces exerted by adjacent molecules in the interior.
Many insects take advantage of this property to walk on the surface of puddles or ponds without sinking (part (b) in Figure 11.10 "The Effects of the High Surface Tension of Liquid Water").
The values of the surface tension of some representative liquids are listed in Table 11.4 "Surface Tension, Viscosity, Vapor Pressure (at 25A°C Unless Otherwise Indicated), and Normal Boiling Points of Common Liquids". Dynamic equilibrium Equilibrium – Equilibrium – because of the unchanging amounts Dynamic – Dynamic – because reaction is still occurring The person stays in the same place!
The equilibrium constants K c & K p are numerical constants for a given reaction at a given temperature. This stream of kinetic energy then drives a generator that is installed safely and economically at ground level.
Maintenance should be cheaper and safer as there are no large turbine blades located 200 feet in the air and tied to the generator, also 200 feet in the air.
Other liquids are used to manufacture objects that we use every daya€”for example, a solid material is converted to a liquid, the liquid is injected into a mold, and it is then solidified into complex shapes under conditions that are carefully controlled. As a result, the Ca€“O bond dipoles partially reinforce one another and generate a significant dipole moment that should give a moderately high boiling point. The net effect is that the first atom causes the temporary formation of a dipole, called an induced dipoleA short-lived dipole moment that is created in atoms and nonpolar molecules adjacent to atoms or molecules with an instantaneous dipole moment., in the second. In small atoms such as He, the two 1s electrons are held close to the nucleus in a very small volume, and electrona€“electron repulsions are strong enough to prevent significant asymmetry in their distribution. 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. In contrast, the hydrides of the lightest members of groups 15a€“17 have boiling points that are more than 100A°C greater than predicted on the basis of their molar masses. Because a hydrogen atom is so small, these dipoles can also approach one another more closely than most other dipoles. The expansion of water when freezing also explains why automobile or boat engines must be protected by a€?antifreezea€? (we will discuss how antifreeze works in Chapter 13 "Solutions") and why unprotected pipes in houses break if they are allowed to freeze. 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. Because a sphere has the smallest possible surface area for a given volume, intermolecular attractive interactions between water molecules cause the droplet to adopt a spherical shape. Note the correlation between the surface tension of a liquid and the strength of the intermolecular forces: the stronger the intermolecular forces, the higher the surface tension. The mathematical relationship between concentration and rate of reaction is related with the orders of reaction. Thus no large blades or ice from the blades to fly off and travel hundreds of feet endangering people and property.
To understand such processes, our study of the macroscopic properties of matter must include an understanding of the properties of liquids and the interconversion of the three states of matter: gases, liquids, and solids. Because molecules in a liquid move freely and continuously, molecules always experience both attractive and repulsive dipolea€“dipole interactions simultaneously, as shown in Figure 11.4 "Both Attractive and Repulsive Dipolea€“Dipole Interactions Occur in a Liquid Sample with Many Molecules".
Thus a substance such as HCl, which is partially held together by dipolea€“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. Acetone contains a polar C=O double bond oriented at about 120A° to two methyl groups with nonpolar Ca€“H bonds. In larger atoms such as Xe, however, the outer electrons are much less strongly attracted to the nucleus because of filled intervening shells. This maximizes the number of attractive interactions and minimizes the number of water molecules at the surface.
For example, water, with its strong intermolecular hydrogen bonding, has one of the highest surface tension values of any liquid, whereas low-boiling-point organic molecules, which have relatively weak intermolecular forces, have much lower surface tensions. Fire fighting for an electrical fire in the generator could be done as it is on the ground. 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 11.2 "Relationships between the Dipole Moment and the Boiling Point for Organic Compounds of Similar Molar Mass". The Ca€“O bond dipole therefore corresponds to the molecular dipole, which should result in both a rather large dipole moment and a high boiling point.
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 the boiling points of nonpolar substances increase rapidly with molecular mass, C60 should boil at a higher temperature than the other nonionic substances.
Larger atoms tend to be more polarizable than smaller ones because their outer electrons are less tightly bound and are therefore more easily perturbed. Hence raindrops are almost spherical, and drops of water on a waxed (nonpolar) surface, which does not interact strongly with water, form round beads (see the chapter opener photo).
Mercury is an apparent anomaly, but its very high surface tension is due to the presence of strong metallic bonding, which we will discuss in more detail in Chapter 12 "Solids". The noise level should be much less than the large wind turbines, and hopefully, there would be no infrasound (0 - 20 hertz), which people cannot hear, but is a problem. Using what we learned in Chapter 9 "Molecular Geometry and Covalent Bonding Models" 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.
The ease of deformation of the electron distribution in an atom or molecule is called its polarizabilityThe ease of deformation of the electron distribution in an atom or molecule.. 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.
Hydrogen bonds are especially strong dipolea€“dipole interactions between molecules that have hydrogen bonded to a highly electronegative atom, such as O, N, or F.
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. 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 resulting partially positively charged H atom on one molecule (the hydrogen bond donor) can interact strongly with a lone pair of electrons of a partially negatively charged O, N, or F atom on adjacent molecules (the hydrogen bond acceptor). Attractive interactions between the polar substances and water cause the water to spread out into a thin film instead of forming beads.
The bridging hydrogen atoms are not equidistant from the two oxygen atoms they connect, however. Because of strong Oa‹…a‹…a‹…H hydrogen bonding between water molecules, water has an unusually high boiling point, and ice has an open, cagelike structure that is less dense than liquid water. In contrast, each oxygen atom is bonded to two H atoms at the shorter distance and two at the longer distance, corresponding to two Oa€“H covalent bonds and two Oa‹…a‹…a‹…H hydrogen bonds from adjacent water molecules, respectively. 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.

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