Intermolecular Forces © 2009, Prentice-Hall, Inc. Chapter 11 Intermolecular Forces, Liquids, and Solids Chemistry, The Central Science, 11th edition Theodore L. Brown, H. Eugene LeMay, Jr., and Bruce E. Bursten

Intermolecular Forces © 2009, Prentice-Hall, Inc. Overview of Chapter 11 – Intermolecular Forces Review/On your own: compare liquids and solids with gases in the context of the KMT describe the three intermolecular forces and apply this knowledge to predictions about substances’ properties, such as melting and boiling points, solubilities, surface tension, etc. explain what vapor pressure is, and how it differs between contained and uncontained containers interpret heating curves, vapor pressure graphs, and phase diagrams calculate enthalpy changes associated with phase changes relate vapor pressure to volatility of a liquid and the intermolecular forces interpret vapor pressure graphs

Intermolecular Forces © 2009, Prentice-Hall, Inc. Overview of Chapter 11 – Intermolecular Forces More depth: interpret phase diagrams describe the three intermolecular forces and apply this knowledge to predictions about substances’ properties, such as melting and boiling points, solubilities, surface tension, etc. NOTHING ON SOLIDS (11.7 and 11.8)

Intermolecular Forces © 2009, Prentice-Hall, Inc. States of Matter ShapeVolumeDensityCompressible?Mass? Solid Liquid Gas

Intermolecular Forces © 2009, Prentice-Hall, Inc. States of Matter The fundamental difference between states of matter is the distance between particles.

Intermolecular Forces © 2009, Prentice-Hall, Inc. States of Matter Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.

Intermolecular Forces © 2009, Prentice-Hall, Inc. The States of Matter 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 attractions between the particles.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Intermolecular Forces The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces. Note: Sometimes scientists use the term van der Waals force to refer to dispersion forces only. Be sure to describe the interaction as thoroughly as possible to clearly distinguish amongst them.

Intermolecular Forces © 2009, Prentice-Hall, Inc. van der Waals Forces London dispersion forces (weakest, present in all molecules) Dipole-dipole interactions Hydrogen bonding

Intermolecular Forces © 2009, Prentice-Hall, Inc. Ion-Dipole Interactions Ion-dipole interactions (a fourth type of force), are important in solutions of ions. The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

Intermolecular Forces © 2009, Prentice-Hall, Inc. 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.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Dipole-Dipole Interactions The more polar the molecule, the higher is its boiling point. Note: There is a stronger relationship between the strength of IMFs and boiling points, more so than with melting points.

Intermolecular Forces © 2009, Prentice-Hall, Inc. London Dispersion Forces While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.

Intermolecular Forces © 2009, Prentice-Hall, Inc. London Dispersion Forces At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

Intermolecular Forces © 2009, Prentice-Hall, Inc. London Dispersion Forces Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

Intermolecular Forces © 2009, Prentice-Hall, Inc. London Dispersion Forces London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

Intermolecular Forces © 2009, Prentice-Hall, Inc. London Dispersion Forces These forces are present in all molecules, whether they are polar or nonpolar. The tendency of an electron cloud to distort in this way is called polarizability.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Factors Affecting London Forces The shape of the molecule 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. Note the differences in boiling points

Intermolecular Forces © 2009, Prentice-Hall, Inc. Factors Affecting London Forces The strength of dispersion forces tends to increase with increased molecular weight. BECAUSE larger atoms have larger electron clouds which are easier to polarize.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Sample Exercise 11.1 (p. 443) The dipole moments of acetonitrile, CH 3 CN, and methyl iodide, CH 3 I, are 3.9 D and 1.62 D, respectively. a)Which of these substances will have the greater dipole-dipole attractions among its molecules? b)Which of these substances will have the greater London dispersion attractions? c) The boiling points of CH 3 CN and CH 3 I are K and K, respectively. Which substance has the greatest overall attractive forces?

Intermolecular Forces © 2009, Prentice-Hall, Inc. Practice Exercise 11.1 Of Br 2, Ne, HCl, HBr, and N 2, which is likely to have a)the largest intermolecular dispersion forces? b)the largest dipole-dipole attractive forces?

Intermolecular Forces © 2009, Prentice-Hall, Inc. Which Have a Greater Effect? Dipole-Dipole Interactions or Dispersion Forces If two molecules are of comparable size and shape, dipole-dipole interactions will likely be the dominating force. If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

Intermolecular Forces © 2009, Prentice-Hall, Inc. How Do We Explain This? The nonpolar series (SnH 4 to CH 4 ) follow the expected trend. The polar series follows the trend from H 2 Te through H 2 S, but water is quite an anomaly.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Hydrogen Bonding The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong. We call these interactions hydrogen bonds.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Hydrogen Bonding Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed, small and can approach an electronegative atom very closely, making it much easier to interact with it.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Sample Exercise 11.2 (p. 444) In which of the following substances is hydrogen bonding likely to play an important role in determining physical properties: methane (CH 4 ), hydrazine (H 2 NNH 2 ), methyl fluoride (CH 3 F), or hydrogen sulfide (H 2 S)?

Intermolecular Forces © 2009, Prentice-Hall, Inc. Practice Exercise 11.2 In which of the following substances is significant hydrogen bonding possible: methylene chloride (CH 2 Cl 2 ), phosphine (PH 3 ), hydrogen peroxide (HOOH), or acetone (CH 3 COCH 3 )?

Intermolecular Forces © 2009, Prentice-Hall, Inc. Summarizing Intermolecular Forces

Intermolecular Forces © 2009, Prentice-Hall, Inc. Sample Exercise 11.3 (p. 447) List the substances BaCl 2, H 2, CO, HF, and Ne in order of increasing boiling points.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Practice Exercise 11.3 a)Identify the intermolecular forces present in the following substances, and b)select the substance with the highest boiling point: CH 3 CH 3, CH 3 OH, CH 3 CH 2 OH.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Intermolecular Forces Affect Many Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution.

Intermolecular Forces © 2009, Prentice-Hall, Inc. 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.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Surface Tension Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Changes

Intermolecular Forces © 2009, Prentice-Hall, Inc. Energy Changes Associated with Changes of State The heat of fusion is the energy required to change a solid at its melting point to a liquid.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Energy Changes Associated with Changes of State The heat of vaporization is defined as the energy required to change a liquid at its boiling point to a gas. Note how much more energy is needed to vaporize a substance than to melt it.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Energy Changes Associated with Changes of State 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 a phase change.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Sample Exercise 11.4 (p. 451) Calculate the enthalpy change upon converting 1.00 mol of ice at -25 o C to water vapor (steam) at 125 o C under a constant pressure of 1 atm. The specific heats of ice, water, and steam are 2.03 J/g-K, 4.18 J/g-K, and 1.84 J/g-K, respectively. For H 2 O,  H fus = 6.01 kJ/mol, and  H vap = kJ/mol. Note: multi-step problem, tedious, but not hard (56.0 kJ)

Intermolecular Forces © 2009, Prentice-Hall, Inc kJ 6.01 kJ  melting 7.52 kJ 40.7 kJ  vaporization 0.83 kJ

Intermolecular Forces © 2009, Prentice-Hall, Inc. Practice Exercise 11.4 What is the enthalpy change during the process in which g of water at 50.0 o C is cooled to ice at o C? (Use the specific heats and enthalpies for phase changes given in Sample Exercise 11.4.) (-60.4 kJ)

Intermolecular Forces © 2009, Prentice-Hall, Inc. Vapor Pressure At any temperature some molecules in a liquid have enough energy to escape. As the temperature rises, the fraction of molecules that have enough energy to escape increases.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Vapor Pressure As more molecules escape the liquid, the pressure they exert increases.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Vapor Pressure The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Vapor Pressure The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure. The normal boiling point is the temperature at which its vapor pressure is 760 torr.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Sample Exercise 11.5 (p. 455) Use the preceding figure to estimate the boiling point of diethyl ether under an external pressure of 0.80 atm.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Practice Exercise 11.5 At what external pressure will ethanol have a boiling point of 60 o C?

Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagrams Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagrams The circled line is the liquid-vapor interface. It starts at the triple point (T), the point at which all three states are in equilibrium.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagrams It ends at the critical point (C); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagrams Each point along this line is the boiling point of the substance at that pressure.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagrams The circled line in the diagram below is the interface between liquid and solid. The melting point at each pressure can be found along this line.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagrams Below the triple point the substance cannot exist in the liquid state. Along the circled line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.

Intermolecular Forces © 2009, Prentice-Hall, Inc. 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.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagram of Water 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. –Why?

Intermolecular Forces © 2009, Prentice-Hall, Inc. Read from passage from “Dirt”.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO 2 sublimes at normal pressures.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagram of Carbon Dioxide The low critical temperature and critical pressure for CO 2 make supercritical CO 2 a good solvent for extracting nonpolar substances (like caffeine)

Intermolecular Forces © 2009, Prentice-Hall, Inc. Complete the rest of the Chapter 11 packet on your own. Note that the section on packing of crystalline structures is not required. (most of Sec. 11.7)

Intermolecular Forces © 2009, Prentice-Hall, Inc. Sample Exercise 11.6 (p. 458) Referring to the figure at right, describe any changes in the phases present when H 2 O is a) kept at 0 o C while the pressure is increased from that at point 1 to that at point 5 (vertical line); b)kept at 1.00 atm while the temperature is increased from that at point 6 to that at point 9 (horizontal line).

Intermolecular Forces © 2009, Prentice-Hall, Inc. Practice Exercise 11.6 Using the phase diagram of CO 2, describe what happens when the following changes are made in a CO 2 sample initially at 1 atm and -60 o C: a) pressure increases at constant temperature of -60 o C from 1 atm to 60 atm, b) temperature increases from -60 o C to -20 o C at constant 60 atm pressure.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Solids We can think of solids as falling into two groups: –crystalline, in which particles are in highly ordered arrangement.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Solids We can think of solids as falling into two groups: –amorphous, in which there is no particular order in the arrangement of particles.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Attractions in Ionic Crystals In ionic crystals, ions pack themselves so as to maximize the attractions and minimize repulsions between the ions.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Crystalline Solids Because of the ordered in a crystal, we can focus on the repeating pattern of arrangement called the unit cell.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Crystalline Solids There are several types of basic arrangements in crystals, like the ones depicted above.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Crystalline Solids We can determine the empirical formula of an ionic solid by determining how many ions of each element fall within the unit cell.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Ionic Solids What are the empirical formulas for these compounds? –(a) Green: chlorine; Gray: cesium –(b) Yellow: sulfur; Gray: zinc –(c) Gray: calcium; Blue: fluorine CsClZnSCaF 2 (a)(b)(c)

Intermolecular Forces © 2009, Prentice-Hall, Inc. Types of Bonding in Crystalline Solids

Intermolecular Forces © 2009, Prentice-Hall, Inc. Covalent-Network and Molecular Solids Diamonds are an example of a covalent- network solid, in which atoms are covalently bonded to each other. –They tend to be hard and have high melting points.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Covalent-Network and Molecular Solids Graphite is an example of a molecular solid, in which atoms are held together with van der Waals forces. –They tend to be softer and have lower melting points.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Metallic Solids Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces. In metals valence electrons are delocalized throughout the solid.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Sample Integrative Exercise (p. 469) The substance CS 2 has a melting point of o C and a boiling point of 46.3 o C. Its density at 20 o C is 1.26 g/cm 3. It is highly flammable. a) What is the name of this compound? b) List the intermolecular forces that CS 2 molecules would have with each other. c) Predict what type of crystalline solid CS 2(s) would form.

Intermolecular Forces © 2009, Prentice-Hall, Inc. Sample Integrative Exercise (p. 469) d) Write a balanced equation for the combustion of this compound in air. (You will have to decide on the most likely oxidation products.) e)The critical temperature and pressure for CS 2 are 552 K and 78 atm, respectively. Compare these values with those for CO 2 and discuss the possible origins of the differences. f) Would you expect the density of CS 2 at 40 o C to be greater or less than at 20 o C? What accounts for the difference?