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Chapter 11 Intermolecular Forces, Liquids, and Solids. 11.1 Comparing Gases, Liquids, and Solids. The fundamental difference between states of matter is the distance between particles.
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11.1 Comparing Gases, Liquids, and Solids • The fundamental difference between states of matter is the distance between particles. • Because particles are closer together in the solid and liquid states, they are referred to as condensed phases.
Properties of states of matter: See Table 11.1 • A gas is compressible, expands to assume the volume and shape of its container, flows readily, and diffusion within a gas occurs readily. • A liquid assumes the shape of the portion of the container it occupies, does not expand to fill its container, is virtually incompressible, flows readily, and diffusion within a liquid occurs slowly. • A solid retains its own shape and volume, is virtually incompressible, does not flow, and diffusion within a solid occurs extremely slowly. • The state a substance is in at a particular temperature and pressure depends on two antagonistic entities: • The kinetic energy of the particles • The strength of the intermolecular forces between the particles • In solids, the intermolecular forces between particles are strong enough to hold them close and virtually lock them in place. • In liquids, the intermolecular forces hold particles close but do not lock them in place—they are free to move past one another. • In gases, the particles are far apart and lack strong attractive forces between one another.
11.2 Intermolecular Forces • They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities. • Three types of intermolecular attractive forces exist between neutral molecules. These forces are also called van der Waals forces. They are: • Dipole-dipole interactions • Hydrogen bonding • London dispersion forces • The intermolecular attractions between molecules are not nearly as strong as the intramolecular attractions (bonding forces) that hold compounds together.
1. Dipole-Dipole Interactions • Molecules that have permanent dipoles are attracted to each other. • The positive end of one is attracted to the negative end of the other and vice-versa. • These forces are only important when the molecules are close to each other.
The more polar the molecule, the higher is its boiling point.
2. London Dispersion Forces • There are no dipole-dipole forces between nonpolar atoms and molecules. • But, the motion of electrons in an atom or molecule can create an instantaneous, or momentary, dipole moment. • In He atoms, for example, the average distribution of electrons around a nucleus is spherically symmetrical. The atom has no dipole moment. • However, an instantaneous distribution of the He 1s orbital electrons may find both electrons on one side of the nucleus. • In that instant, the atom would have an instantaneous dipole moment. • The helium atom is now polar, with an excess of electrons on one side as compared to the other.
London dispersion forces (or dispersion forces) are attractions between an instantaneous dipole and an induced dipole. • These forces, like dipole-dipole forces, are significant only when molecules are close together. • The temporary dipole moment on one atom induces a temporary dipole – an induced dipole -- on an adjacent atom, causing the atoms to be attracted to each other.
Factors Affecting Dispersion Forces • The strength of the dispersion force depends on the ease with which the charge distribution in a molecule can be distorted to induce a momentary dipole. • The tendency of an electron cloud to distort in this way is called polarizability. • The more polarizable a molecule is, the greater its dispersion forces. • Dispersion forces tend to increase in strength with increasing molecular weight. • Larger molecules tend to be more polarizable because of a greater number of electrons located further from the nucleus.
The effect of molecular weight and size on boiling points is shown below: • Dispersion forces operate between all molecules, whether they are polar or not. • Polar molecules can experience dipole-dipole and dispersion forces at the same time. • Dispersion forces between polar molecules commonly contribute more to intermolecular attractions than do dipole-dipole forces.
The shape of the molecule also affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane). • This is due to the increased surface area in n-pentane. • Which Have a Greater Effect: Dipole-Dipole Interactions or Dispersion Forces? • If two molecules of a substance are of comparable size and shape, dipole-dipole interactions will likely be the dominating force. • If one molecule of a substance is much larger than another, dispersion forces will likely determine its physical properties.
3. Hydrogen Bonding • The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong. • These interactions are called hydrogen bonds. • Due to hydrogen bonds, there are very strong intermolecular forces between the H and O atoms on adjacent water molecules. • Hydrogen bonding is the cause of water’s • high boiling point • high melting point • high specific heat • high heat of vaporization
Boiling point as a function of molecular weight • The nonpolar series (SnH4 to CH4) follow the expected trend. • The polar series follows the trend from H2Te through H2S. • Water, despite its relatively low molecular weight, is an anomaly due to strong hydrogen bonding attractions.
Hydrogen bonds are unique dipole-dipole attractions • Bonds between H and O, F or N are very electronegative with H at the positive (+) end of the dipole. • The hydrogen proton is nearly exposed. • The small size of the hydrogen atom allows it to approach an electronegative atom very closely, allowing it to strongly interact with it. • The energies of hydrogen bonds vary from about 4 kJ/mol to about 25 kJ/mol. • Hydrogen bonds are much weaker than ordinary chemical bonds, but are generally stronger than other dipole-dipole force or dispersion forces.
Hydrogen bonds and water/ice densities • The lower density of ice compared to liquid water can be understood in terms of hydrogen-bonding interactions. • In ice, the lower temperature slows down the motion of water molecules. • The water molecules assume an orderly, open arrangement with each water molecule forming hydrogen bonds to four other water molecules. • These hydrogen bonds create open cavities in the ice structure. The water molecules are kept further apart than in liquid water and, hence, the density of ice is less than in liquid water. Ice floats!
As the temperature increases, molecular motions increase and the cavities of the ice structure collapse. • The hydrogen bonding in the liquid water is more random than in ice, but it is strong enough to hold water molecules closer together giving it a more dense structure than in ice.
Ion-Dipole Interactions • A fourth type of force, ion-dipole interactions are an important force in solutions of ions. • The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.
11.3 Some Properties of Liquids The strength of the attractions between particles can greatly affect the properties of a substance or solution.
Viscosity • Resistance of a liquid to flow is called viscosity. • It is related to the ease with which molecules can move past each other. • Viscosity increases with stronger intermolecular forces and decreases with higher temperature.
Surface Tension • Surface tension results from the net inward force experienced by the molecules on the surface of a liquid. • Surface molecules are attracted only by other surface molecules and the molecules below the surface. • Surface molecules pack closer together than molecules below the surface. • Surface tension is the energy required to increase the surface area of a liquid by a unit amount. • Unlike surface molecules, molecules in interior of the liquid experience attraction in all directions • Surface tension of water: • Surface tension of mercury:
Energy Changes Accompanying Phase Changes • Heat of Fusion (ΔHfus): Energy required to change a solid at its melting point to a liquid. Temp stays constant. • Heat of Vaporization (ΔHvap): Energy required to change a liquid at its boiling point to a gas. Temp stays constant. • Heat of Sublimation (ΔHsub): Energy required to change a solid directly to the gaseous state. The sum of ΔHfus+ ΔHvap.
Heating Curves • The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other. • The temperature of the substance does not rise during the phase change. • Heating curve for water (at right). • Heating curves graph temperature versus amount of heat added. • Calculate energy (heat) added: • BC and DE: Phase changes (kJ/mol) • AB, CD, EF: Temp increase within phase (q=mcΔt) • See sample exercise 11.4
Critical temperature: the highest temperature at which a gas can form a distinct liquid phase. It is the highest temperature at which a liquid can exist. • Critical pressure: the pressure required to bring about liquefaction at the critical temperature. • The critical temperature and pressure are important to engineers and others who need to liquify gases for storage. • See Table 11.5. Nonpolar, low M.W. substances with weak intermolecular attractions have lower c.p. and c.t. than those that are polar or have higher M.W. • Water and ammonia have exceptionally high c.p. and c.t. as a consequence of strong intermolecular hydrogen-bonding forces.
Supercritical fluid • A supercritical fluid (SCF) is any substance at a temperature and pressure above its critical point, where distinct liquid and gas phases do not exist. • It can effuse through solids like a gas, and dissolve materials like a liquid. • Supercritical fluids are used for a range of industrial and laboratory processes. • Carbon dioxide and water are the most commonly used supercritical fluids, being used for decaffeination and power generation, respectively. • Read Chemistry At Work, page 459.
11.5 Vapor Pressure • Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid or solid in a closed container. • At any temperature, some molecules in a liquid have enough energy to escape by evaporation. • As the temperature rises, the fraction of molecules that have enough energy to escape increases.
As more molecules escape the liquid, the pressure they exert increases. • The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.
Volatility, Vapor Pressure, and Temperature • Substances with high vapor pressures evaporate more quickly than substances with lower vapor pressure. • Liquids that evaporate more readily are said to be volatile. • Volatility is the tendency of a liquid to evaporate. • Liquids that evaporate quickly from an open container have high volatility and high vapor pressure. • In contrast, liquids with low vapor pressures have low volatility.
Vapor pressure and boiling point • The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure. • At this point bubbles of vapor are able to form in the fluid. • The temperature at which a given liquid boils increases or decreases with increases or decreases in pressure. • The boiling point of a liquid at 1 atm pressure is called its normal boiling point.
11.6 Phase Diagrams • Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.
The AB line is the liquid-vapor interface. • It starts at the triple point (A), where all three states are in equilibrium. • It ends at the critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other. • Each point along the AB line is the boiling point of the substance at that pressure.
The AD line is the interface between liquid and solid. • The melting point at each pressure can be found along this line. • Below A the substance cannot exist in the liquid state. • Along the AC line the solid and gas phases are in equilibrium • The sublimation point at each pressure is along this line.
Phase Diagram of Water • Note the high critical temperature and critical pressure. • These are due to the strong van der Waals forces between water molecules. • The slope of the solid–liquid line is negative. • This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.
Phase Diagram of Carbon Dioxide • Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.
11.7 Solids Crystalline solids • The atoms, ions, or molecules are ordered in well-defined arrangements. • These solids have flat surfaces, or faces —particles are in highly order, that make definite angles with each other. • These solids have highly regular shapes (see Fig. 11.29). There are two groups of solids:
Amorphous solids • A solid whose particles have no orderly structure. • They lack well-defined faces and shapes. • Examples include glasses, rubber, etc.
Covalent-network solids • Molecular solids • Consist of atoms or molecules held together by intermolecular forces. • Graphite is an example: layers of atoms are held together with van der Waals forces. • Also, sucrose, dry ice, water, for example • They tend to be softer and have lower melting points. • Consist of atoms held together by covalent bonds. • Diamonds are an example: they are hard and have high melting points.
Ionic solids • Consist of ions held together by ionic bonds. • The strength of the ionic bond (see discussion of lattice energy in sec 8.2) depends greatly on the charges of the ions. • NaCl, with charges of 1+ and 1-, melts at 801oC, whereas MgO, with charges of 2+ and 2-, melts at 2852oC. • Ionic compounds have an orderly arrangement of ions. Their structures depends on the charges and relative sizes of their ions. Metallic Solids • Consist entirely of metal atoms. • Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces. • Bonding due to delocalized valence electrons throughout the solid.
