Liquids & Solids. Mr. Hollister Holliday Legacy High School Regular & Honors Chemistry

Transcription

1 Liquids & Solids Mr. Hollister Holliday Legacy High School Regular & Honors Chemistry 1

2 Liquids 2

3 Properties of the States of Matter: Liquids High densities compared to gases. Fluid. The material exhibits a smooth, continuous flow as it moves. Take the shape of their container(s). Keep their volume, do not expand to fill their container(s). Cannot be compressed into a smaller volume. Particles in a Liquid 3

4 Liquids The particles in a liquid are closely packed, but they have some ability to move around. The close packing results in liquids being incompressible. But the ability of the particles to move allows liquids to take the shape of their container and to flow. However, they don t have enough freedom to escape and expand to fill the container(s). 4

5 Properties of Liquids: Surface Tension Liquids tend to minimize their surface a phenomenon we call surface tension. This tendency causes liquids to have a surface that resists penetration. The stronger the attractive force between the molecules, the larger the surface tension. Water Strider 5

6 Properties of Liquids: Viscosity Some liquids flow more easily than others. The resistance of a liquid s flow is called viscosity. The stronger the attractive forces between the molecules, the more viscous the liquid is. Also, the less round the molecule s shape, the larger the liquid s viscosity. Some liquids are more viscous because their molecules are long and get tangled in each other, causing them to resist flowing. 6

7 Evaporation Over time, liquids evaporate the molecules of the liquid mix with and dissolve in the air. The evaporation happens at the surface. Molecules on the surface experience a smaller net attractive force than molecules in the interior. All the surface molecules do not escape at once, only the ones with sufficient kinetic energy to overcome the attractions will escape. 7

8 Escaping from the Surface The process of molecules of a liquid breaking free from the surface is called evaporation. Also known as vaporization. Evaporation is a physical change in which a substance is converted from its liquid form to its gaseous form. The gaseous form is called a vapor. 8

9 Factors Effecting the Rate of Evaporation Liquids that evaporate quickly are called volatile liquids, while those that do not are called nonvolatile. Increasing the surface area increases the rate of evaporation. More surface molecules. Increasing the temperature increases the rate of evaporation. Raises the average kinetic energy, resulting in more molecules that can escape. Weaker attractive forces between the molecules = faster rate of evaporation. 9

10 Reconnecting with the Surface When a liquid evaporates in a closed container, the vapor molecules are trapped. The vapor molecules may eventually bump into and stick to the surface of the container or get recaptured by the liquid. This process is called condensation. A physical change in which a gaseous form is converted to a liquid form. 10

11 Dynamic Equilibrium Evaporation and condensation are opposite processes. Eventually, the rate of evaporation and rate of condensation in the container will be the same. Opposite processes that occur at the same rate in the same system are said to be in dynamic equilibrium. 11

12 Evaporation and Condensation When water is just added to the flask and it is capped, all the water molecules are in the liquid. Shortly, the water starts to evaporate. Initially the rate of evaporation is much faster than rate of condensation Eventually, the condensation and evaporation reach the same speed. The air in the flask is now saturated with water vapor. 12

13 Vapor Pressure Once equilibrium is reached, from that time forward, the amount of vapor in the container will remain the same. As long as you don t change the conditions. The partial pressure exerted by the vapor is called the vapor pressure. The vapor pressure of a liquid depends on the temperature and strength of intermolecular attractions. 13

14 Vapor Pressure Curves 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 or 1 atm.

15 Water: A Unique and Important Substance Water is found in all three states on Earth. As a liquid, it is the most common solvent found in nature. Without water, life as we know it could not exist. The search for extraterrestrial life starts with the search for water. 15

16 Water Liquid at room temperature. Most molecular substances that have a molar mass (18.02 g/mol) similar to water s are gaseous. Relatively high boiling point. Expands as it freezes. Most substances contract as they freeze. Causes ice to be less dense than liquid water. 16

17 Solids 17

18 Properties of the States of Matter: Solids High densities compared to gases. Nonfluid. They move as an entire block rather than a smooth, continuous flow. Keep their own shape, do not take the shape of their container(s). Keep their own volume, do not expand to fill their container(s). Cannot be compressed into a smaller volume. 18

19 Solids Some solids have their particles arranged in an orderly geometric pattern. We call these crystalline solids. Salt and diamonds. Other solids have particles that do not show a regular geometric pattern over a long range. We call these amorphous solids. Plastic and glass. 19

20 Crystalline Solids 20

21 Types of Crystalline Solids 21

22 Molecular Crystalline Solids Molecular solids are solids whose composite units are molecules. Solid held together by intermolecular attractive forces. Dispersion, dipole-dipole, or H-bonding. Generally low melting points 22

23 Ionic Crystalline Solids Ionic solids are solids whose composite units are formula units. Solid held together by electrostatic attractive forces between cations and anions. Cations and anions arranged in a geometric pattern called a crystal lattice to maximize attractions. Generally higher melting points than molecular solids. Because ionic bonds are stronger than intermolecular forces. 23

24 Atomic Crystalline Solids Atomic solids are solids whose composite units are individual atoms. Solids held together by either covalent bonds, dispersion forces, or metallic bonds. Melting points vary depending on the attractive forces between the atoms. 24

25 Practice Classify Each of the Following Crystalline Solids as Molecular, Ionic, or Atomic. H 2 O(s) molecular. Si(s) atomic. C 12 H 22 O 11 (s) molecular. CaF 2 (s) ionic. Sc(NO 3 ) 3 (s) ionic. 25

26 Types of Atomic Solids 26

27 Types of Atomic Solids: Covalent Network Covalent atomic solids have their atoms attached by covalent bonds. Effectively, the entire solid is one giant molecule. Because covalent bonds are strong, these solids have very high melting points. Because covalent bonds are directional, these substances tend to be very hard. 27

28 Types of Atomic Solids: Nonbonding Nonbonding atomic solids are held together by dispersion forces. Because dispersion forces are relatively weak, these solids have very low melting points All the noble gases form nonbonding atomic solids. 28

29 Types of Atomic Solids: Metallic Metallic solids are held together by metallic bonds. Metal atoms release some of their electrons to be shared by all the other atoms in the crystal. The metallic bond is the attraction of the metal cations for the mobile electrons. Often described as islands of cations in a sea of electrons. 29

30 Metallic Bonding The model of metallic bonding can be used to explain the properties of metals. The luster, malleability, ductility, and electrical and thermal conductivity are all related to the mobility of the electrons in the solid. The strength of the metallic bond varies, depending on the charge and size of the cations, so the melting points of metals vary as well. 30

31 Practice Decide if Each of the Following Atomic Solids Is Covalent, Metallic, or Nonbonding. diamond covalent. neon nonbonding. iron metallic. 31

32 Phase Changes 32

33 Heating Curve Diagrams 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.

34 Phase Diagrams Phase diagrams display the state of a substance at various pressures and temperatures.

35 Phase Diagrams The AB line is the liquid-vapor interface. It starts at the triple point (A), the point at which all three states are in equilibrium.

36 Phase Diagrams It ends at the critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

37 Phase Diagrams Each point along this line is the boiling point of the substance at that pressure.

38 Phase Diagrams The AD line is the interface between liquid and solid. The melting point at each pressure can be found along this line.

39 Phase Diagrams 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.

40 Phase Diagram of CO 2 40

41 Sublimation Sublimation is a physical change in which the solid form changes directly to the gaseous form. Without going through the liquid form. Like melting, sublimation is endothermic. 41

42 Intermolecular Forces 42

43 Why Are Molecules Attracted to Each Other? Intermolecular attractions are a result of attractive forces between opposite charges. + ion to ion. + end of one polar molecule to end of another polar molecule. H-bonding is especially strong. Even nonpolar molecules will have temporary induced dipoles. Larger charge = stronger attraction. 43

44 Intermolecular Forces The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.

45 Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

46 Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces.

47 van der Waals Forces London dispersion forces Dipole-dipole interactions Hydrogen bonding Johannes van der Waals

48 Types of Intermolecular Forces Type of force Relative strength Present in Example London Dispersion force Weak, but increases with molar mass All atoms and molecules H 2 Dipole Dipole force Moderate Only polar molecules HCl Hydrogen Bond Strong Molecules having H bonded to F, O, or N HF 48

49 Dispersion Forces Also known as London forces or instantaneous dipoles. Caused by distortions in the electron cloud of one molecule inducing distortion in the electron cloud on another. Distortions in the electron cloud lead to a temporary dipole. The temporary dipoles lead to attractions between molecules dispersion forces. All molecules have attractions caused by dispersion forces. 49

50 Instantaneous Dipoles 50

51 Strength of the Dispersion Force Depends on how easily the electrons can move, or be polarized. The more electrons and the farther they are from the nuclei, the larger the dipole that can be induced. Strength of the dispersion force gets larger with larger molecules. 51

52 Dispersion Force and Molar Mass Noble Gas Molar Mass (g/mol) Boiling Point (K) He Ne Ar Kr Xe

53 Boiling Point, C Relationship Between Dispersion Force and Molecular Size BP, Noble Gas BP, Halogens BP, XH Period 53

54 Practice The Following Are All Made of Non Polar Molecules. Pick the Substance in Each Pair with the Highest Boiling Point. CH 4 or C 3 H 8. BF 3 or BCl 3. CO 2 or CS 2. 54

55 Permanent Dipoles Because of the kinds of atoms that are bonded together and their relative positions in the molecule, some molecules have a permanent dipole. Polar molecules. 55

56 Dipole-to-Dipole Attraction Polar molecules have a permanent dipole. A + end and a end. The + end of one molecule will be attracted to the end of another. 56

57 Polarity and Dipole-to-Dipole Attraction Molar Mass Boiling Dipole (g/mol) Point, C size, D CH 3 CH 2 CH CH 3 -O-CH CH 3 - CH=O CH 3 -C N

58 Attractive Forces Dispersion forces All molecules Dipole-to-dipole forces Polar molecules

59 Hydrogen Bonding Hydrogen atoms bound to a N, O or F atom have strong intermolecular attractions. Unusually high melting and boiling points. Unusually high solubility in water. This kind of attraction is called a hydrogen bond. 59

60 Properties and H-Bonding Name Formula Molar mass (g/mol) Structure Boiling point, C Melting point, C Solubility in water H H Ethane C 2 H H C C H Immiscible H H H H C O H Ethanol CH 4 O Miscible H 60

61 Intermolecular H-Bonding 61

62 H-Bonds vs. Chemical Bonds Hydrogen bonds are not chemical bonds. Hydrogen bonds are attractive forces between molecules. Chemical bonds are attractive forces that make molecules. 62

63 Boiling Point, C Relationship Between H-Bonding and Intermolecular Attraction H 2 O HF H 2 Te NH 3 H 2 S H 2 Se SnH CH 4 SiH 4 GeH 4 BP, HX BP, H2X BP, H3X -200 Period BP, XH4 63