Liquids & Solids

1. Liquids & Solids • Chapter 13

2. Heat of Fusion/Vaporization • H2O(s) ----> H2O(l) Hfusion = 6.02 kj/mol • H2O(l) ----> H2O(g)  Hvaporization = 40.6 kj/mol • From the  Ho values above, which two states are most similar? • How do the attractive forces between the molecules compare in these two states to the third state?

3. Liquids & Solids • The liquid and solid states are considered to be the condensed states. • Liquids & solids have much higher densities than gases and are not compressible.

4. Three States of Matter

5. Some Properties of a Liquid • Surface Tension: The resistance to an increase in its surface area (polar molecules). A sphere has the maximum volume for the minimum surface area.

6. Some Properties of a Liquid • Capillary Action: Spontaneous rising of a liquid in a narrow tube. Viscosity: Resistance to flow (molecules with large intermolecular forces).

7. Some Properties of a Liquid • Cohesive forces exist between molecules of a liquid. Adhesive forces exist between the liquid and its container.

8. Water & Its Phase Changes • Water is the most common liquid--covering about 70% of the earth’s surface. • Water is necessary for life, moderates the earth’s climate, a means of transportation, and cools many industrial processes. • Pure water is colorless, odorless, tasteless, freezes at 0 oC and boils at 100 oC at 1 atm.

9. Normal Melting/Freezing Point • The temperature at which a solid melts or freezes at 1 atm pressure--0 oC for water. • Molecules break loose from lattice points and solid changes to liquid. (Temperature is constant as melting occurs.) • vapor pressure of solid = vapor pressure of liquid

10. Normal Boiling Point • The boiling temperature of a liquid at one atmosphere pressure--100 oC for water. • Constant temperature when added energy is used to vaporize the liquid. • vapor pressure of liquid = pressure of surrounding atmosphere

11. Figure 13.3: Both liquid water and gaseous water contain H2O molecules

12. Boiling Point • What effect does altitude have on boiling point? • Where on earth would have the highest boiling point? • Where on earth would have the lowest boiling point? Higher altitude--lower b.p. Dead Sea Top of Mt. Everest--70 oC

13. Freezing of Water • When water freezes, it expands about 1/9th in volume. This causes water pipes and engine blocks to burst when frozen. • The density of ice is less than water and, therefore, ice floats. If ice were more dense than water, lakes and rivers would freeze from the bottom up and aquatic life could not survive.

14. Endothermic Melting Boiling Exothermic Condensing Freezing Physical Changes & Energy Changes Changes of state (melting, freezing, boiling, & condensing) are constant temperature processes!!!!!

15. Heats of Fusion & Vaporization • The molar heat of fusion of ice is 6.02 kJ/mol. • Hfusion = 6.02 kJ/mol • The molar heat of vaporization of water is 40.6 kJ/mol at 100 oC. • Hvaporization = 40.6 kJ/mol

16. Heating curve for water. Q = (ms t)ice + m Hf + (ms t) water + m Hv + (mst)steam Q = KE & PE + PE + KE & PE + PE + PE & KE

17. Calculating Energy Changes • Calculate the amount of energy required to melt 8.5 g of ice at 0 oC. • Q = mHfusion • Q = (8.5g HOH)(1mol/18.02g HOH)(6.02kJ/mol) • Q = 2.8 kJ

18. Calculating Energy ChangesLiquid toGas • Calculate the energy (in kJ) required to heat 25g of liquid water from 25 oC to 100 oC and change it to steam at 100 oC. The specific heat (s) of water is 4.18 J/gCo. • Q = mst + mHvaporization • Q = (25g)(4.18 J/gCo)(75 Co)(1kJ/1000J) + (25g)(1mol/18.02g)(40.6kJ/mol) • Q = 7.8 kJ + 57 kJ • Q = 65 kJ

19. Calculating Energy ChangesSolid toGas • Calculate the energy (in kJ) required to melt 15g of ice at 0 oC, heat the water to 100 oC, and vaporize it to steam at 100 oC. • Q =mHfusion+mst + mHvaporization • Q = (15g)(1mol/18.02g)(6.02kJ/mol) + (15g)(4.18 J/gCo)(100 Co)(1kJ/1000J) + (15g)(1mol/18.02g)(40.6kJ/mol) • Q = 5.0 kJ + 6.3 kJ + 34 kJ • Q = 45 kJ

20. Intramolecular within the molecule covalent bonding ionic bonding Intermolecular between molecules dipole-dipole forces hydrogen bonding London Dispersion Forces Types of Bonding When ice changes to liquid and then to vapor, the intramolecular forces (covalent bonds) stay intact, only the weaker hydrogen bonds between molecules weaken and break. These are, therefore, physical changes.

21. Intermolecular Forces • Forces between (rather than within) molecules. • dipole-dipole attraction: molecules with dipoles orient themselves so that “+” and “” ends of the dipoles are close to each other. (1 % as strong as covalent or ionic.) • hydrogen bonds: dipole-dipole attraction in which hydrogen is bound to a highly electronegative atom. (F, O, N)

22. Electrostatic interaction of two polar molecules.

23. The polar water molecule and hydrogen bonds among water molecules.

24. The boiling points of the covalent hydrides of the elements in Groups 4A, 5A, 6A, & 7A. The high boiling points of HOH, NH3, HF is due to hydrogen bonding.

25. Instantaneous and induced dipole moments between nonpolar molecules -- London Dispersion Forces. LDF forces are both weak and short-lived.

26. - - - - - - - - - - - - - - + - - + - - - - - - - - + - + + - - + - + + - - - - - - - - - - - - - - + + - - - - - - - - London Dispersion Forces • Also Known As Induced Dipoles • Size of the London Dispersion Force depends on the number of electrons and shapes of molecules • the larger the molar mass, the larger the induced dipole

27. London Dispersion Forces • relatively weakforces that exist among noble gas atoms and nonpolar molecules. (Ar, C8H18) • caused by instantaneous dipole, in which electron distribution becomes asymmetrical. • the ease with which electron “cloud” of an atom can be distorted is called polarizability.

28. Equilibrium Liquid just poured into open container, little vapor Evaporation faster than Condensation Evaporation as fast as Condensation Vapor Pressure

29. Vapor Pressure • . . . is the pressure of the vapor present at equilibriumwith its liquid. • . . . is determined principally by the size of the intermolecular forces in the liquid. • . . . increases significantly with temperature. • Volatile liquidshave high vapor pressures.

30. Vapor Pressure • Low boiling point • high vapor pressure. • weak intermolecular forces. • Low vapor pressure • high molar masses. • strong intermolecular forces.

31. Boltzman Distribution -- number of molecules in a liquid with a given energy versus kinetic energy at two different temperatures.

32. Rate of Evaporation & High Vapor Pressure • Liquids evaporate more rapidly and have a high vapor pressure when the liquid: • 1. has weak intermolecular forces. • 2. is made up of lighter molecules. • 3. is at a high temperature.

33. Vapor Pressure • Which of the following pairs have the highest vapor pressure? • 1. HOH(l) or CH3OH(l) • 2. CH3OH(l) or CH3CH2CH2CH2OH(l) Why?

34. Sublimation • Change of a solid directly to a vapor without passing through the liquid state. • Iodine • Dry Ice • Moth Balls

35. Types of Solids • Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS2)]. • Amorphous solids: considerable disorder in their structures (glass).

36. Representation of Components in a Crystalline Solid • Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

37. Representation of Components in a Crystalline Solid • Unit Cell: The smallest repeating unit of the lattice. • simple cubic -- 1 atom/cell • body-centered cubic -- 2 atoms/cell • face-centered cubic -- 4 atoms/cell

38. Three cubic unit cells and the corresponding lattices.

39. Types of Crystalline Solids • Atomic Solid: contains atoms at the lattice points (diamond) -- variable melting points. • Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl) -- high melting points, strong electrostatic forces between + & - ions. • Molecular Solid: discrete covalently bonded molecules at each of its lattice points (sucrose, ice) -- low melting points, weak attraction.

40. Crystalline Solids

41. Three crystalline solids -- a) atomic solid, b) ionic solid, and c) molecular solid.

42. The properties of solids are determined primarily by the nature of the forces that hold the solid together.

43. Molecular Solids • molecular units at each lattice position. • strong covalent bonding within molecules. • relatively weak forces between molecules. • London Dispersion Forces -- CO2, I2, P4, & S8. • Hydrogen Bonding -- H2O, NH3, & HF.

44. Figure 13.16: (Left) Sulfur crystals contain S8 molecules. (Right) White phosphorus contains P4 molecules.

45. Network Solids • Composed of strong directional covalent bonds that are best viewed as a “giant molecule”. • brittle • do not conduct heat or electricity • carbon, silicon-based • graphite, diamond, ceramics, glass

46. Network solid structure of diamond.

47. Bonding Models for Metals • Metals are malleable, ductile, good conductors of heat and electricity, have high melting points, and are durable. The bonding in metals is strong but nondirectional.

48. Bonding Models for Metals • Electron Sea Model: A regular array of metals in a “sea” of electrons. The cations are mutually attracted to the valence electrons--this attraction holds the metal together. • The mobile electrons can conduct heat and electricity and the cations are fairly easily moved--making the metal malleable & ductile.

49. + + + + + + + + + e- e- e- e- e- e- e- e- + + + + + + + + + e- e- e- e- e- e- e- e- + + + + + + + + + Electron Sea Model • Atomic Solids that are made of metal atoms • metal atoms release their valence electrons • metal cations fixed in a “sea” of mobile electrons • Leads to strong attractions that are non-directional

50. Metal Alloys • 1. Substitutional Alloy: some metal atoms replaced by others of similar size. • brass = Cu/Zn Substances that have a mixture of elements and metallic properties.