CHAPTER 10. Liquids and solids

Transcription

1 CHAPTER 10 Liquids and solids

2 Forces Solids and Liquids Gases can be compressed because gas particles are far apart. (ideal gases do not interact with one another NO IMF s) Liquid and solids are not compressible because the particles are close together. What holds solid and liquid particles close together?

3 Forces INTERmolecular forces Inside molecules (INTRAmolecular) the atoms are bonded to each other. INTERmolecular refers to the forces between the molecules. These are what hold the particles together in the condensed states.

4 Molecular forces Forces Weak Forces Big Picture Strong Forces London < Dipole- < Hydrogen < Metallic < Ionic < Covalent Dispersion Dipole Bonding Bonding Network INTER INTRA

5 Forces INTER Weaker (attractions) 1. London Dispersion Forces (LDF) (aka: Van der Waals forces or induced dipole forces) 2. Dipole dipole polar attractions (Coulombs Law) 3. H- bonding (strong dipole) INTRA Stronger (bonds) 4. Metallic 5. Ionic 6. Covalent Phase Changes the molecules stay intact (Intra). Energy used to overcome INTERmolecular forces

6 Forces INTER - + INTRA

7 Forces Dipole - Dipole Polar molecules have dipoles (δ + and δ - ). The opposite ends of the dipole can attract each other so the molecules stay close together. 1% as strong as covalent bonds Weaker with greater distance. Small role in gases because.they are far apart

8 Forces Maximize a7rac9on Minimize repulsion Coulombic A7rac9on

9 Example 10.1 A What is the molecular shape of each of the following molecules? Which molecules exhibit dipole- dipole forces? a. I - 3 b. BrF 5 c. SF 4 d. CO 2

10 100 H 2 O Example 10.1 B Boiling Points Predict the loca9ons of NH 3 H 2 O, and HF HF 0ºC H NH 3 H 2 S H 2 Se 2 Te SbH 3 HCl AsH 3 HI PH HBr 3 GeH SnH 4 4 SiH 4 CH 200 4

11 Forces Hydrogen Bonding Especially strong dipole-dipole forces when H is attached to F, O, or N These three because- They have high electronegativity. They are small enough to get close together. Between two molecules (HF and HF) Or Within a single molecule (DNA, protiens, Amino Acids)

12 Forces O O δ + H H δ + δ + H H δ + O Water is special δ + H H δ + This gives water an especially high melting and boiling point. The density of most liquids and solids doesn t change with temperature Water s density does!

13 Forces

14 Example 10.1 C Draw each of the following molecules. Which of the following molecules exhibit H- bonding? a. HBr b. NH 3 c. CH 4 d. CH 3 OH e. CH 3 OCH 3

15 Forces London Dispersion Forces Non - polar molecules also exert forces on each other. Otherwise, no solids or liquids. Electrons are not evenly distributed at every instant in time!! Have an instantaneous dipole. (temporarily polar!!) Induces a dipole in the atom next to it. Induced dipole- induced dipole interaction.

16 Forces Example H H H H Non polar

17 Forces Example δ+ δ H H H H Instantaneous Dipole b/c of uneven e- distribution

18 Forces Example δ+ δ δ+ δ H H H H INDUCES a Dipole in a neighboring molecule London Dispersion Forces = Induced Dipole Induced Dipole

19 Forces London Dispersion Forces Weak, short lived. All substances have LDF forces but they are only important in NONPOLAR molecules. Lasts longer at low temperature, Eventually long enough to make liquids. Two Factors 1. More electrons, more polarizable! (Bigger = Stronger) 2. Larger Surface Area, (Bigger) higher melting and boiling points. Conisder phases of halogens I 2 -solid Br 2 liquid Cl 2 - gas **Polarizability enhanced by presence of PI BONDS** Also called Van der Waal s forces.

20 Forces Strength of IMF s increases with Magnitude of Dipole Polarizability of nonpolar molecule (size) Mix and Match! H 2 S and H 2 O = dipole-dipole interactions CH 4 and H 2 S = dipole-induced dipole interaction!! (non) (polar) just very weak

21 Example 10.1 D Forces London Dispersion Forces The boiling point of argon is O C. Why is it so low? How does this boiling point help prove that London dispersion forces exist? The boiling point of xenon is O C. Why is it higher than that of argon?

22 Forces Example 10.1 E The Effect of Intermolecular Forces Iden9fy the substance with the strongest LDF H 2 Br 2 CH 4 Kr Xe

23 Forces Recap Intermolecular Force Molecules Cause London Dispersion Nonpolar Induced polarity Dipole- Dipole Polar Polarity Hydrogen Bonds Polar with small molecules H- N, H- O or H- F high EN F,O,N **THIS IS THE BASIS FOR LIKE DISSOLVES LIKE : )

24 Example 10.1 F The Effect of Intermolecular Forces Forces Iden9fy the dominant force in each of the following molecules H 2 O CH 4 CH 3 CH 3 OH H 2 S C 4 H 10 HF SiH 4 N 2 HBr CO 2 Kr NaCl

25 Example 10.1 G The Effect of Intermolecular Forces Forces Put the following substances in order from lowest to highest boiling point. C 2 H 6 NH 3 F 2

26 Example 10.1 H The Effect of Intermolecular Forces Forces Iden9fy the substance with the Lowest Boiling Point in each set 1. Kr O 2 C 2 H 6 Br 2 2. CO 2 H 2 O HCl Ar

27 Liquids Liquids Particles allowed to flow past one another TRANSLATION Their arrangement and properties are affected by 1. Type of IMF 2. Strength of IMF

28 PROPERTIES Liquids Surface Tension: The resistance to an increase in its surface area Capillary Ac9on: Spontaneous rising of a liquid in a narrow tube. Beading: Internal pull from intermolecular forces Viscosity: Resistance to flow (molecules with large intermolecular forces). Vapor Pressure: The gas pressure coming off a liquid Vola9lity: the tendency to vaporize *Stronger intermolecular forces cause each of these to increase*

29 Liquids Molecules at the the top are only pulled inside. Molecules in the middle are attracted in all directions. Surface tension Minimizes surface area.

30 Liquids

31 Liquids Capillary Action Liquids spontaneously rise in a narrow tube. Cohesion: intermolecular forces among the molecules of the liquid Adhesion: forces between the liquid and its container. Glass is polar. It attracts water molecules (polar).

32 Liquids

33 Liquids Beading If a polar substance is placed on a non-polar surface. There are cohesive, But no adhesive forces.

34 Questions Liquids 1. Why do liquids tend to bead up when on solid surfaces? 2. What are cohesive forces? Adhesive forces? What causes them? 3. What is surface tension? Why does it arise? 4. Why does water form a concave meniscus? Why do mercury form a convex meniscus? 5. What is viscosity?

35 Example 10.2 A Liquids Properties of Liquids Which would have a higher surface tension, H 2 O or C 6 H 14? Why? Would the shape of the H 2 O meniscus in a glass tube be the same or different than C 6 H 14?

36 Liquids Example 10.2 B Properties of Liquids Iden9fy the molecule in each set with. Highest surface most highest tension beading capillary ac9on N 2 LDF F 2 H 2 S dipole- dipole PH 3 CH 3 OH dipole- induced dipole NH 3

37 Liquids Vapor pressure The pressure above the liquid at equilibrium. Liquids with high vapor pressures evaporate easily. They are called volatile. Decreases with increasing intermolecular forces. Bigger molecules (bigger LDF) More polar molecules (dipole-dipole)

38 Liquids Vaporization is an endothermic process - it requires heat. Energy is required to overcome intermolecular forces. Responsible for cool earth. Why we sweat, cools you down Increases with Temperature

39 Liquids Vapor Pressure Vaporization - change from liquid to gas at boiling point. Evaporation - change from liquid to gas below boiling point Heat (or Enthalpy) of Vaporization (ΔH vap )- the energy required to vaporize 1 mol at 1 atm.

40 Liquids Dynamic equilibrium When first sealed the molecules gradually escape Some condense back to a liquid. Rate of vaporization remains constant but the rate of condensation increases as more molecules are available. Until.

41 Liquids Dynamic equilibrium When first sealed the molecules gradually escape Some condense back to a liquid. Rate of vaporization remains constant but the rate of condensation increases as more molecules are available. Until.

42 Liquids Dynamic equilibrium When first sealed the molecules gradually escape Some condense back to a liquid. Rate of vaporization remains constant but the rate of condensation increases as more molecules are available. Until.

43 Liquids Dynamic equilibrium Rate of Vaporization = Rate of Condensation Molecules are constantly changing phase Dynamic The total amount of liquid and vapor remains constant Equilibrium

44 Vacuum Liquids A barometer will hold a column of mercury 760 mm high at one atm P atm = 760 torr Dish of Hg

45 Vacuum Liquids A barometer will hold a column of mercury 760 mm high at one atm. P atm = 760 torr If we inject a volatile liquid in the barometer it will rise to the top of the mercury. Dish of Hg

46 Water Liquids A barometer will hold a column of mercury 760 mm high at one atm. P atm = 760 torr If we inject a volatile liquid in the barometer it will rise to the top of the mercury. There it will vaporize and push the column of mercury down. Dish of Hg

47 Liquids Water Vapor 736 mm Hg The mercury is pushed down by the vapor pressure. P atm = P Hg + P vap P atm - P Hg = P vap = 24 torr Dish of Hg

48 Liquids

49 Temperature Effect Liquids # of molecules T 1 Energy needed to overcome intermolecular forces Kine<c energy

50 Liquids At higher temperature more molecules have enough energy - higher vapor pressure. # of molecules T 1 Energy needed to overcome intermolecular forces T 2 Kine<c energy

51 Liquids Boiling Point Reached when the vapor pressure of the liquid equals the external pressure. Normal boiling point is the boiling point at 1 atm pressure. Super heating - Heating above the boiling point. Supercooling - Cooling below the freezing point.

52 Liquids Example 10.8 A The Effect of Intermolecular Forces III Iden9fy the molecule in each set with. Highest Vapor Lowest vapor lowest Pressure Pressure Boiling Point CO 2 C 2 H 4 C 2 H 4 H 2 O C 4 H 10 C 4 H 10 CH 3 OCH 3 CH 4 CH 4 NaCl O 2 O 2

53 Solids SOLIDS do NOT undergo Translation Two major types Packing Arrangements Amorphous- those with much disorder in their structure. (examples: glass, peanut brittle) Crystalline- have a regular 3-D structure. Stronger forces = increased hardness = more structure

54 Solids ATOMIC and MOLECULAR SOLIDS Covalent compounds (between nonmetals) LDF & Dipole-Dipole & H-Bonding Individual molecules/atoms ATTRACTED to each other and can form a solid Molceuclar Solids are a collection of individual molecules Atomic Solids are a collection of individual atoms

55 Molecular forces Weak Forces Big Picture Strong Forces London < Dipole- < Hydrogen < Metallic < Ionic < Covalent Dispersion Dipole Bonding Bonding Network INTER Atomic or molecular solids INTRA Large single repeating molecule

56 Questions Solids 1. Define molecular solid? 2. Give some examples of molecular solids. 3. Why are some larger nonpolar molecular substances solid at 25 C? 4. How would an ionic solid differ from a molecular solid?

57 Solids INTRA molecular forces Not individual molecules! form a large repeating pattern of bonds Metallic bonds metal + metal (repeating) Ionic bond Lattice of ions Covalent Network repeating covalent bonds

58 Solids Metallic Bonding Metallic Metal + Metal ( +Metal + Metal + Metal ) Shell Model - The outmost orbitals of metal atoms overlap allowing electrons to easily move between atoms (a sea of electrons); result is delocalized valence electrons that move easily between atoms The Electron Sea Model -positive metal ions -free flowing electrons

59 Solids Metal Properties The Electron Sea Model explains why properties of metals: Conduct flowing electrons Malleable bendable or can pound into a shape Ductile can stretch into wire These properties result because deforming the solid does not change the sea, some translation (slide past one another) with force Low Volatility individual atoms can t escape

60 Solids Alloys Alloys are a mixture of metals Properties of allows are based on size of component atoms Alloy Examples: Brass Steel Nichrome Cast Iron Stainless Steel Pig Iron White gold Solder Most jewelry Sterling Silver Bronze Pewter

61 Solids Alloys Properties based on size of component atoms Interstitial Alloy putting smaller atoms between spaces This kind of alloy strengthens metallic bonds making a more rigid structure; but this decrease malleability.

62 Solids Alloys Properties based on size of component atoms Substitutional Alloy substituting atoms of similar size Result similar density still malleable

63 Solids Alloys Stainless Steel Steel is made of Fe and C When Cr is used in steel, it reacts with air (O 2 ) at the surface and forms a coating New chemical properties at surface Stainless resists corrosion

64 Solids Example 10.4 A Metal Properties 1. Why do metals conduct electricity? 2. Why are metals malleable

65 Solids IONIC (a review) Strong coulombic attraction produce a very low vapor pressure (none?) Q 1 Q 2 r form regular (predictable) lattice structures ionic substances interact with substances with dipole-dipole forces making them soluble in polar compounds

66 Example 10.5 A The Effect of Intermolecular Forces Solids Iden9fy the substance with the highest mel9ng point in each set 1. NaCl, MgCl 2, CaCl 2, AlCl 3 2. NaCl, KBr, LiF, RbI

67 Solids Example 10.5 B Ionic Properties 1. Do ionic solids conduct electricity? Why or why not? 2. Are ionic solids malleable? Why or why not? (what happens if you hit a salt crystal with a hammer?)

68 Covalent Network Solids Nonmetal atoms that are covalently bonded in a network Properties: HIGH melting point HARDNESS incredibly rigid bonds Generally form in the carbon group because it can form 4 covalent bonds

69 Solids Carbon There are many allotropes of solid carbon. Amorphous- coal, uninteresting. Graphite- slippery, conducts electricity. Diamond- hardest natural substance on earth, insulates both heat and electricity. Buckyball - C 60 basis for nanotechnology Nanotubes - nanotechnology How the atoms in these network solids are connected explains properties.

70 Graphite - 2-D molecule Solids sp 2 hybridized room for pi bonding delocalized pi bonding network! Properties: High melting point network bonds Conductivity delocalized e- region Soft layers can slide past one another the attraction between layers is LDF

71 Solids Diamond. sp 3 hybridized locked into tetrahedral shape. Properties hardness strong σ bonds insulator - no pi bond network, e - can t flow

72 Solids Silicon Silicon is an important component in the Earth s crust (rock) CO 2 - is a gas (has 2 pi bonds) SiO 2 - (silica) is solid and is TOO BIG for pi bonds Forms a covalent network similar to diamond (SiO 2 is a common AP answer on tests as an example of a covalent network solid)

73 Solids Silicon Pure silicon is an insulator (like carbon) BUT it is a metalloid and can be a semiconductor when DOPED DOPING injecting impurities (other elements) to the silicon 1. creates conductivity 2. more impurities possible at higher temps 3. more conductive at higher temps

74 Solids Silicon IMPURITIES two types 1. Elements with one more valence electron :: 5 total e- (P, As) After all four bonds are taken, extra electron to freely move N-TYPE semiconducting (negative type)

75 Solids Silicon IMPURITIES two types 2. Elements with one less valence electron :: 3 total e- (B, Al) Only 3 covalent bonds can form. A positive hole is formed in fourth spot holes can be filled by stealing e- from neighbor moving electrons P-TYPE semiconductor (positive type)

76 Solids Silicon DOPING adding impurities (other elements) to the silicon Junctions between n-doped and p-doped materials control e- flow Basis of modern electronics Just Watch this Video. Transistors

77 Solids Silicon 1. What makes Silicon an N-doped semiconductor? 2. What makes Silicon a P-doped semiconductor?

78 Solids Example 10.5 A Combined Practice Iden9fy the dominant force in each of the following molecules SiO 2 C (graphite) NaCl CH 3 CH 3 NH 2 H 2 S C 4 H 10 W CH 3 (CH 2 ) n CH 3 SiH 4 Ne C (diamond) Ca Ca(NO 3 ) 2 Fe PH 3 PCl 6

79 Solids Example 10.7 A Types of Solids Classify each of the following substances as to the type of solid it forms. Fe C 2 H 6 CaCl 2 Graphite

80 10.8- Energy Changes of state The graph of temperature versus heat applied is called a heating curve. Temp a solid turns to a liquid = melting point. the energy is the enthalpy of fusion ΔH fus the temperature at which the vapor pressure of the liquid equals 1 atm is the (normal) boiling point of a substance energy required is enthalpy of vaporization ΔH vap

81 Hea<ng Curve for Water Energy Water and Steam Steam Water Ice Water and Ice

82 Hea<ng Curve for Water Energy Slope is Heat Capacity Heat of Vaporiza<on Heat of Fusion

83 10.8- Energy Heating Curves Hea9ng/cooling a pure substance will either Change its state of ma7er Make the temperature change q = nδh q = mc p ΔT H fus = 6.0 kj/mol specific heat of ice = 2.1 J/g C H vap = 43.9 kj/mol specific heat liquid = 4.2 J/g C specific heat of water vapor = 1.8 J/g C Total Energy change required can be obtained by adding up each component separately

84 Example 10.8 B Heating Curve Energy How much energy does it take to convert 130. g of ice at - 40 O C to steam at 160 O C?