CHAPTER 10: INTERMOLECULAR FORCES: THE UNIQUENESS OF WATER Problems: 10.2, 10.6, , , ,

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1 CHAPTER 10: INTERMOLECULAR FORCES: THE UNIQUENESS OF WATER Problems: 10.2, 10.6, , , , INTERACTIONS BETWEEN IONS Ion-ion Interactions and Lattice Energy electron affinity: energy associated with 1 mole of gaseous atoms to gain 1 mole of electrons X(g) + e X (g) A quantitative measure of an atom s ability to accept an electron The higher the electron affinity more likely an atom can gain an electron Note: Electron affinity is NOT the same as electronegativity (in Chapter 8). The following shows the energy changes associated with forming NaCl from solid Na and Cl 2 gas: So why do ionic compounds form if the process requires energy (ionization energy, etc.)? Ionic bonds form when cations and anions bond to form the 3D network (crystal lattice) of ions. Many bonds forming a lot of energy is released. The energy released is far greater than the ionization energy required. lattice energy: energy released when an ionic compound forms from gaseous ions quantitative measure of the strength of the ionic bonds in the compound Note: Since lattice energy is released, it is negative. CHEM 161 Chapter 10 page 1 of 12

2 Ex. 1: Based on the types of bonds formed, compare sodium chloride s lattice energy of -787 kj/mol with magnesium oxide s lattice energy of kj/mol. Explain why much more energy is released for the formation of the magnesium oxide lattice compared to sodium chloride. Ex. 2: As bond strength increases, lattice energy. increases decreases Thus, lattice energy and bond strength are related. directly inversely 10.2 INTERACTIONS INVOLVING POLAR MOLECULES 10.3 DISPERSIONS FORCES INTERMOLECULAR FORCES (IMFs) phase (=physical state): solid, liquid, or gas condensed states: solid and liquid phases In solids and liquids, molecules are attracted to one another. These attractions are called intermolecular forces. To undergo a change in physical state (e.g. solid liquid or liquid gas), these intermolecular forces must be overcome. Ion-Dipole Forces Attraction between an ion and the oppositely charged end of a polar molecule e.g. between Na + and the negative end (O atom) of a H 2 O molecule or between Cl and the positive end (H atoms) of a H 2 O molecule Note that when an ionic compound like NaCl dissolves in water, the formation of ion-dipole forces between the Na + (or Cl ) ions with water molecules results in the ionic bonds breaking. CHEM 161 Chapter 10 page 2 of 12

3 London Forces or Dispersion Forces (also called Induced-dipole Forces) named after a German-American physicist named Fritz London ( ) In nonpolar molecules (shown as green) the electrons may shift and concentrate on one end temporary dipole (red = + end; blue = end) The temporary dipole in one molecule causes the electrons to shift in an adjacent molecule another temporary dipole. Their temporary dipoles cause them to be attracted to one another. Attraction lasts only until the electrons shift again, and the temporary dipoles go away. Nonpolar molecules experience only LDF s. Polar molecules also experience LDFs, but they also experience other IMF s that are often stronger. The bigger the molecule the more electrons it has and the further the electrons are from nucleus the greater its polarizability (ability to distort electron clouds to get a temporary dipole) the stronger its London forces. Note: Molar mass is generally used as a measure of the size and polarizability of a molecule. The higher the molar mass of an atom/molecule, the stronger its London forces. Dipole-Dipole Forces: Attraction between polar molecules generally stronger than London dispersion forces because attraction is due to permanent dipoles that do not go away like induced dipoles Hydrogen Bonding: Especially strong type of dipole-dipole force Exist between molecules that contain the following bonds: H F, H O, H N because these are small atoms with large electronegativity differences very strong dipole in the molecule Strongest type of intermolecular force Responsible for ice being less dense than liquid water, the relatively high melting and boiling points for water, and the bending and twisting in proteins and DNA. CHEM 161 Chapter 10 page 3 of 12

4 Hydrogen bonds are responsible for the bending and twisting in proteins, DNA, and other important biological molecules. Image from Note: Hydrogen bonds are the strongest type of intermolecular forces between different molecules, BUT ionic and covalent bonds (holding ions or atoms together in compounds) are stronger than hydrogen bonds! The term van der Waals forces is also used to refer to intermolecular forces (London/dispersion forces, dipole-dipole forces, or hydrogen bonds) but not ion-dipole forces. How to determine type of intermolecular forces involved: Is the molecule polar or nonpolar? polar Any H-F, H-O, or H-N bonds in the molecule? yes hydrogen bonds and London Forces nonpolar no London Forces dipole-dipole forces and London Forces CHEM 161 Chapter 10 page 4 of 12

5 Ex. 1: Label all of the intermolecular forces between molecules in the following: London/dispersion forces (LDF), dipole-dipole forces (D-D), hydrogen bonding (H bond). i. Br 2 (l) iii. H 2 S(l) ii. NH 3 (l) iv. CCl 4 (l) Ex. 2: For each of the following, i. Identify the bond between atoms as ionic, polar covalent, or nonpolar covalent. ii. Identify the intermolecular forces as London dispersion forces, dipole-dipole forces, and/or hydrogen bonding. Water (H 2 O) A: B: HCN (H=white, C=charcoal, N=blue) A: B: CHEM 161 Chapter 10 page 5 of 12

6 Ex. 3: Indicate the bond(s) or intermolecular forces described for each below: A. ionic bond D. London forces G. ion-dipole forces B. polar covalent bond E. dipole-dipole forces C. nonpolar covalent bond F. hydrogen bonding i. What holds two I 2 molecules together in a sample of I 2 (s)? ii. What holds atoms together in HF? iii. What holds atoms together in a nitrogen molecule? iv. What holds two ammonia molecules together in liquid ammonia? v. What holds two methane molecules together in liquid methane? vi. What holds two CH 2 O molecules together in a sample of CH 2 O(l)? vii. The bonds broken when water boils. viii. The bonds broken when NaCl(s) melts. ix. The bonds formed when KBr dissolves in water. Ex. 4 Circle the molecule in each pair that experiences the stronger intermolecular forces and explain why. a. N 2 or NO Why? b. H 2 S or H 2 O Why? c. Cl 2 or F 2 Why? CHEM 161 Chapter 10 page 6 of 12

7 10.5 VAPOR PRESSURE OF PURE LIQUIDS: Vapor Pressure and Temperature vaporization: liquid gas From a molecular viewpoint, a molecule escaping from the liquid to gaseous state vapor: The gas above a liquid when the liquid and gaseous states are both present In the examples below, there is more vapor is (b) than in (a). vapor pressure: the pressure exerted by the gas molecules above a liquid In the examples above, the sample in (b) has higher vapor pressure than the sample in (a). vaporization: liquid + heat vapor condensation: vapor liquid + heat vaporization Liquid-Gas Equilibrium: liquid + heat condensation vapor When the molecules in the liquid have enough energy, they escape to the gas phase. In a closed system, when enough vapor exists above the liquid, some gaseous molecules condense back to the liquid. Ultimately, the rate of vaporization = the rate of condensation. The system has reached a state of dynamic equilibrium in which the forward process occurs at the same rate as the reverse process. In an open system, when molecules in the liquid have enough energy to escape to the gas phase, the molecules continue to escape in a process called evaporation. The vaporized molecules continue to escape few gas molecules condense to liquid Ultimately, all of the liquid is converted into a gas. Since vaporization requires energy, the liquid molecules take energy from the surroundings, so the temperature of the surroundings decreases. why evaporation is a cooling process that can reduce body temperature CHEM 161 Chapter 10 page 7 of 12

8 Vapor Pressure (v.p.): pressure exerted by gas molecules above a liquid for a molecule to go from liquid to gas, it has to break the intermolecular forces with or attraction to the other liquid molecules around it weaker the intermolecular forces are easier to break more gas molecules higher vapor pressure stronger the intermolecular forces are harder to break fewer molecules go from liquid to gas lower vapor pressure Boiling Point: temperature where vapor pressure of liquid is equal to external pressure (usually atmospheric pressure) For a liquid to boil, its vapor pressure must equal atmospheric pressure. Since intermolecular forces influence vapor pressure, they also influence boiling point. Weaker intermolecular forces more gas molecules higher vapor pressure Less energy is needed to get vapor P = atmospheric P lower boiling point Stronger intermolecular forces fewer gas molecules lower vapor pressure More energy is needed to get vapor pressure = atmospheric pressure higher boiling point normal boiling point is the boiling point at a pressure of 1 atm e.g. Water boils at 100 C at 1 atm but at ~95 C in Denver where P atm =~0.85 atm. Ex. 1: If hexane (C 6 H 14 ) molecules are nonpolar, fill in the blanks for the following statements about water and hexane: a. Hexane s intermolecular forces are. b. Water s intermolecular forces are. c. Water's intermolecular forces are than hexane's. stronger weaker d. Water has a vapor pressure compared with hexane. higher lower e. Water has a boiling point than hexane. higher lower Ex. 2: Circle the compound in each pair below with the higher boiling point: a. CO or O 2 b. HF or HCl c. CCl 4 or CF 4 Ex. 3: NH 3 molecules experience hydrogen bonds while NCl 3 molecules experience dipoledipole forces. Explain how the boiling point for NCl 3 (71 C) can be higher than the boiling point for NH 3 (-33.3 C). CHEM 161 Chapter 10 page 8 of 12

9 AN INTRODUCTION TO STRUCTURES AND TYPES OF SOLIDS Solids can be crystalline or amorphous. CRYSTALLINE SOLIDS: Have an ordered arrangement extending over a long range different types of crystalline solids: molecular solids, covalent network solids, ionic and metallic solids. Molecular Solids: molecules often arranged in recurring patterns exist as molecules held together by intermolecular forces Ionic Solids: lattice of metal and nonmetal ions (eg. NaCl, MgO, etc.) Have high melting points because of 3-D network of ions held together by ionic bonds Some are slightly soluble in water (See Solublity Rules!) Conduct electricity when melted or dissolved in solution Atomic Solids (or Covalent Network Solids) consist of one type of atom typically held together by covalent bonds e.g. diamond exists as covalently bonded carbon atoms covalent network solids never dissolve in any solvent CHEM 161 Chapter 10 page 9 of 12

10 Molecular Solids: consist of molecules held together by intermolecular forces The Structure and Properties of Ice Ice is an example of a molecular solid. The hydrogen bonds between water molecules are responsible for many unusual properties of ice and water. The density of ice (d=0.917 g/cm 3 ) is lower than the density of liquid water (d=1.00 g/cm 3 ). This is an exception! For all other substances, the solid is more dense than its liquid. When ice melts, the water molecules fill in the holes, so liquid water is more dense than ice. Because of hydrogen bonding, water molecules in the ice are in a tetrahedral arrangement "holes" or empty space. Note the hexagonal holes in the molecular-level image for ice at the right. Snowflakes have hexagonal symmetry because of the hexagonal holes in the molecular-level arrangement of water molecules in ice! IONIC CRYSTALS: lattice of metal & nonmetal ions e.g. NaCl, MgO, CaBr 2 3D network of ions held together by electrostatic attraction = ionic bonds high melting points, hard and brittle conduct electricity only when melted or dissolved in solution CHEM 161 Chapter 10 page 10 of 12

11 CARBON AND SILICON: NETWORK ATOMIC SOLIDS SOME COVALENT NETWORK SOLIDS Graphite (a) and diamond (b) are two forms of carbon that are network covalent solids. Network Covalent Solids: covalently bonded atoms forming a large network of indefinite size. (a) Graphite is made up of layers of covalently bonded sp 2 hybridized carbon atoms, and the layers are held together by π bonds that are delocalized over the entire solid. (b) Diamond consists of covalently bonded carbon atoms that form such a network of sp 3 hybridized carbon atoms in 3D tetrahedral structure. Diamond is so hard because so many covalent bonds must be broken to break up the diamond crystal. CHEM 161 Chapter 10 page 11 of 12

12 Ex. 1: Indicate the types of bonds or intermolecular forces that must be broken to melt the following substances, and indicate which solid would have the higher melting point. a. ice or diamond (a form of carbon) b. ice or sodium chloride Ex. 2: If sodium chloride s melting point is 801 C while diamond s melting point is about 3550 C, compare the relative strength of the ionic bonds in sodium chloride with the covalent bonds in diamond. Ex. 3: If graphite is the lead used in pencils, compare the strength of the π bonds between layers in graphite with the σ bonds holding carbon atoms together in a diamond. AMORPHOUS: solids lacking 3D arrangement of atoms may be similar to liquids in appearance since they lack the ordered structure of crystalline solids Silica (SiO 2 ) makes up sand and quartz glass: optically transparent solid of inorganic materials cooled to a rigid but non-crystalline arrangement of Si-O bonded atoms called quartz glass SiO 2 Glass CHEM 161 Chapter 10 page 12 of 12