1. AP CHEMISTRY CHAPTER 10 LIQUIDS AND SOLIDS

2. Intramolecular bonding- sharing electrons Intermolecular bonding- interactions between particles (atoms, molecules or ions)

3. condensed states of matter- liquids and solids *** Changes in state are due to changes in intermolecular bonding (physical changes), not intramolecular bonding(chemical changes).

4. Intermolecular Forces (imf)

5. Dipole-dipole attraction- attraction of molecules having dipole moments for each other. (negative-positive) Polar molecules have dipole-dipole attractions for each other.

6. Dipole-dipole attractions are only about 1% as strong as covalent or ionic bonds. Molecules orient themselves to minimize repulsion and maximize attractions. Examples: HCl, H 2 S

8. Hydrogen bonding- unusually strong dipole-dipole attractions involving hydrogen atoms which are covalently bonded to a very electronegative element and a small, very electronegative atom (F,O,N) ( usually on an adjacent atom) with unshared electrons.

10. Two reasons for strength: 1. small size of H atom allows closeness 2. great polarity Substances with much H bonding have higher boiling points compared to similar substances. Ex. H 2 O, NH 3, HF

11. While we normally think of hydrogen bonding only occurring between two different molecules, very large molecules such as proteins have different parts of the same Molecule forming hydrogen bonding within the same molecule. This helps to form the secondary structures of proteins.

12. London dispersion forces (LDFs) -relatively weak forces (usually) that exist between noble gas atoms and between nonpolar molecules. LDFs also exist in compounds that have dipole-dipole and/or hydrogen bonding. LDFs may be the most important force in large molecules of these types.

13. LDFs occur because of momentary electron imbalance (temporary dipole) which can induce the same in adjacent molecules.

14. At a given moment, there may be more electrons on one end of an atom or molecule than on the other end. This makes one end a little positive and the other end a little negative. This is a temporary dipole and can induce (cause) the same thing in an adjacent molecule (think chain reaction!). Polar molecules and even ions can cause an induced dipole in nearby nonpolar molecules.

15. This force is often very weak, thus the low freezing point of noble gases. The freezing point of noble gases gets higher as we go down the group because heavier atoms have lower velocity and an increased chance of temporary dipoles. These larger molecules or atoms also have more electrons and thus can more easily have the development of temporary dipoles. We say that they are “more polarizeable” when they have more electrons. This causes London dispersion forces to increase down a group.

17. The boiling point of covalent hydrides increases with molecular weight in Group 4. In other groups, the first hydride has a high boiling point because of hydrogen bonding.

19. Metallic bonding- Metallic elements have metal cations in a sea of mobile valence electrons. Ionic bonding- Cations and anions (usually metal and nonmetal) held together by electrostatic (Coulombic) attraction.

20. Covalent network bonding- This is extensive covalent bonding between atoms resulting in giant molecules. Group 4 elements often are involved (C, Si, SiO 2, SiC) Metallic, ionic and covalent network bonding will be discussed further later on in this chapter.

21. General trends in strength of attraction LDF Weakest attraction dipole-dipole H-bonding metallic bonding ionic bonding covalent network solid Strongest attraction

22. Determine the type of intermolecular force each of the following substances have: H 2 HCl Fe H 2 S SiC CO 2

23. Rank these substances from lowest boiling point to highest boiling point. Justify your answers. NaCl, Ge, NH 3, H 2 O, F 2, SO 2

24. Surface tension- resistance of a liquid to an increase in its surface area. Liquids with relatively large intermolecular forces have high surface tensions. ↑ imf = ↑ surface tension Polar molecules have more surface tension than nonpolar molecules.

25. Capillary action- spontaneous rising of a polar liquid in a narrow tube. -caused by 2 forces:

26. cohesive forces: intermolecular forces among the liquid molecules adhesive forces: forces between the liquid molecules and their container. The container must be made of polar material such as glass.

27. A concave meniscus forms because the water’s adhesive forces toward the glass are stronger than its cohesive forces. A nonpolar liquid (or liquid mercury) can produce a convex meniscus (cohesive > adhesive)

28. viscosity- resistance of a liquid to flow Liquids with large intermolecular forces tend to be highly viscous. Ex. glycerol (glycerine) ↑ imf = ↑ viscosity More complex molecules are more viscous because they tangle up.

29. SOLIDS: amorphous solids- very disordered, usually long chain- like molecules twisted up like spaghetti. (plastics, asphalt, rubber) crystalline solids- highly regular arrangement of components

30. lattice- a 3-D system of points designating the centers of the components (atoms, ions or molecules) unit cell - smallest repeating unit of the lattice. (You don’t have to memorize these.)

32. coordination number- number of nearest neighbors surrounding a particle in a crystal For particles of the same size, the higher the coordination number, the greater is the number of particles packed into a given volume of the crystal.

33. X-ray diffraction- method of determining crystal structure. X-rays of a single wavelength are directed at a crystal and are scattered by it, producing a diffraction pattern which can be used to determine the crystal structure.

34. Types of crystalline solids

35. ionic solids (NaCl) -have ions at lattice points, held together by electrostatic (Coulombic) forces.

36. Molecular solids (sucrose) - have molecules at lattice points, held together by LDF, dipole- dipole, &/or hydrogen bonding

37. metallic solids (gold)- have metal cations at lattice points, held together by metallic bonds

38. atomic solids (argon) - have noble gas atoms at lattice points, held together by LDF

39. covalent network solid (diamond and silicon compounds) (essentially one giant molecule) - covalently bonded, have an atom at each lattice point, held together by covalent bonds

40. The properties of a solid depend on the nature of the forces that hold the solid together. Diamond has very strong forces.

41. Metallic crystals - have nondirectional covalent bonding -spherical atoms packed together and bonded to each other equally -called metallic bonding

42. Bonding in Metals -strong and nondirectional -atoms are difficult to separate, but easy to move

43. Two models are used to describe metallic bonding. They are:

44. Electron sea model - This is the simplest model. The metal cations are in a sea of valence electrons. The mobile electrons can conduct heat and electricity. The cations can be easily moved around.

46. Band model (MO model) -This is a more detailed model. The electrons are assumed to travel around the metal crystal in molecular orbitals formed from the valence atomic orbitals of the metal ions. The MOs that result are very closely spaced in energy levels, thus they form a continuous band. Some MOs are empty. Mobile electrons are excited into these empty MOs. The electrons are free to move throughout the metal crystal.

48. Metal alloy - substance that contains a mixture of elements and has metallic properties

49. 2 types: substitutional alloy - some of the host metal atoms are replaced by other metal atoms of similar size. Ex. brass, sterling silver, pewter

50. interstitial alloy -holes in closest packed metal structure are filled by small atoms. The density of an interstitial alloy is higher than a substitutional alloy Ex. steel (carbon in iron)

51. Network solids -atomic solids with strong directional covalent bonds -brittle and don’t conduct heat or electricity -Ex. C and Si -strongest type of bonding

52. Allotropes -different forms of the same element Ex. diamond, graphite, and Buckminsterfullerene are all allotropes of carbon. Sulfur and phosphorus both have allotropes.

53. Diamond has a tetrahedral sp 3 arrangement. Its MOs are far apart in energy, thus no conduction of electricity.

54. Graphite has layers of six member rings (sp 2 ). The pi molecular orbitals allow it to conduct electricity. It has strong bonding within its layers. We can convert graphite to diamond at a pressure of 150,000 atm and a temperature of 2800 o C.

55. Partial Representation of the Molecular Orbital Energies in A) Diamond and B) a Typical Metal

56. Silicon compounds are to geology as carbon compounds are to biology. Most important Si compounds contain Si and O.

57. Silica, SiO 2 (empirical formula) is the fundamental Si-O compound. Quartz and some types of sand are silica. Silica is a network of SiO 4 tetrahedra. The bonds are single because the 3p orbitals are too big to bond strongly with oxygen. Silicates are anions of Si and O.

59. Glass is formed when silica is melted and cooled rapidly. Other substances are added to glass to vary its properties. Glass is amorphous and can be called a “supercooled liquid”. Ceramics contain tiny crystals of silica in a glassy cement.

60. Two Dimensional Representations of (a) a Quartz Crystal and (b) a Quartz Glass

61. Semiconductors Elemental silicon has the same structure as diamond but the energy gap between filled and empty MOs is much smaller. Thus, a few electrons can cross this gap at room temp, making Si a semiconductor. At higher temp, where more energy is available to excite electrons into the conduction bands, the conductivity of Si increases. Most metals have decreased conductivity at higher temps.

62. Doping is a process which increases the conductivity of silicon. A very few of the Si atoms are replaced by other atoms, such as arsenic.

63. Two main types of semiconductors- n-type semiconductor (n stands for negative)- The atom that is substituted for silicon has more valence electrons than silicon. The extra electron(s) are free to move throughout the structure and help it conduct electricity. Silicon with a few arsenic atoms substituted makes an n-type semiconductor.

64. p-type semiconductor (p stands for positive)- The atom being substituted in has fewer electrons than silicon. An example would be boron. This creates a “hole” where the missing electron would be. Mobile electrons can use that hole as they travel through the structure to conduct electricity.

66. Molecular solids Ex. ice, dry ice, sulfur (S 8 ), phosphorus (P 4 )

67. -strong covalent bonding within molecules but relatively weak forces between molecules. -the atoms within the molecule are closer to each other than to atoms of adjacent molecules. This indicates stronger bonding.

68. CO 2, S 8, I 2, and P 4 have no dipole-dipole forces. The last three are solids at room temperature because London dispersion forces are stronger in larger molecules.

69. Ionic Solids -stable, high melting point, held together by strong electrostatic (Coulombic) forces The smaller ions (usually cations) fit into the holes between the larger ions (anions). Attractions are maximized and repulsions are minimized.

70. Vapor Pressure and Changes in State vaporization (evaporation) -endothermic process because we add energy to break intermolecular bonding.

71. Heat of vaporization -energy required to vaporized 1 mole of a liquid at a pressure of 1 atm (  H vap ) ↑ imf = ↑ heat of vaporization Evaporation is a cooling process

72. condensation -process by which vapor molecules reform a liquid

73. dynamic equilibrium rate of condensation = rate of vaporization

76. vapor pressure - pressure of the vapor present at equilibrium -measured by a barometer P vapor = P atmosphere  P Hg column

77. volatile -liquids which evaporate rapidly and have high vapor pressure

78. Vapor pressure is affected by two factors:

79. 1. molar mass - At a given temperature, heavy molecules have lower velocities than light molecules and thus have a lower tendency to escape from the liquid surface. A liquid with a high molar mass tends to have a small vapor pressure. ↑ molar mass = ↓ vapor pressure

80. 2. intermolecular forces -Molecules with strong intermolecular forces also tend to have low vapor pressure because they need lots of energy to escape. ↑ imf = ↓ vapor pressure

82. Vapor pressure increases with temperature (higher KE) but it is not a direct relationship. *You do not need to be able to solve VP-temp problems.

83. As T increases, avg. KE increases and more particles have enough energy to vaporize.

85. Sublimation -direct change of a solid to a gas Ex. dry ice, freeze drying, iodine

86. Heating curve - a plot of temperature versus time for a process where energy is added at a constant rate.

88. Heat of fusion -  H fus enthalpy change that occurs when a solid melts ↑ imf = ↑ heat of fusion Temperature remains constant as phase changes. Below 0 o C, The VP of ice is less than the VP of liquid water.

89. Normal melting point -the temperature at which the solid and liquid have the same vapor pressure at 1 atm total pressure. ↑ imf = ↑ melting point

90. Boiling -occurs when VP of the liquid = external pressure Normal boiling point - the temperature at which the VP of the liquid is exactly 1 atm.

91. Ex. How many joules are needed to convert 5.0 g of ice at -15 o C to steam at 130 o C? (specific heat of ice =2.1 J/g o C, specific heat of water = 4.18 J/g o C, specific heat of steam =2.0 J/g o C, heat of fusion of ice = 6.01 kJ/mol, heat of vaporization of water = 40.7 kJ/mol) Phase changes and temp changes q = 5.0g (2.1J/g o C)15 o C = 160J = 0.16kJ 5.0g ice 1 mol ice 6.01 kJ 18.0 g 1 mol ice = 1.7 kJ

92. q = 5.0g (4.18J/g o C)100 o C = 2100J = 2.1kJ 5.0g water 1 mol water 40.7 kJ 18.0 g water 1 mol water = 11 kJ q = 5.0g (2.0J/g o C)30 o C = 300J = 0.30kJ q = 0.16 + 1.7 + 2.1 + 11 + 0.30 = 15 kJ

93. Phase Diagrams -a convenient way of representing the phases of a substance as a function of temperature and pressure. -represents a closed system Not tested on AP Test

95. triple point - the point on a phase diagram at which all three states of a substance are present. -For water, the triple point is at 0.0098 o C and 4.588 torr. -Solid, liquid and gas have identical VP at the triple point. Not tested on AP Test

96. Critical temperature - temperature above which the vapor cannot be liquefied no matter what pressure is applied. Above the critical temp, the substance is called a supercritical fluid and it has properties of both a liquid and a gas. Not tested on AP Test

97. Critical pressure -pressure required to produce liquefaction at the critical temperature. Not tested on AP Test

98. Critical point -critical temp and critical pressure Not tested on AP Test

100. The freezing point of H 2 O is less than 0 o C when the pressure is greater than 1 atm. (ice skating) Not tested on AP Test

103. The slope of the curve representing equilibrium between liquids and gases is always positive. Not tested on AP Test

104. The slope of the curve representing equilibrium between solids and liquids is positive if the solid is more dense than the liquid (CO 2 ) Not tested on AP Test

106. The slope of the curve representing equilibrium between solids and liquids is negative if the solid is less dense than the liquid. (H 2 O). Not tested on AP Test