Chapter 11 Intermolecular Forces 11 1 Intermolecular Forces

Chapter 11 Intermolecular Forces

11. 1: Intermolecular Forces (IMF) § IMF < intramolecular forces (covalent, metallic, ionic bonds) § IMF strength: solids > liquids > gases § Boiling points and melting points are good indicators of relative IMF strength. 2

11. 2: Types of IMF 1. Electrostatic forces: act over larger distances in accordance with Coulomb’s law a. Ion-ion forces: strongest; found in ionic crystals (i. e. lattice energy) 3

b. Ion-dipole: between an ion and a dipole (a neutral, polar molecule/has separated partial charges) 4 § Increase with increasing polarity of molecule and increasing ion charge. Ex: Compare IMF in Cl- (aq) and S 2 - (aq). d+ d+ d- Cl- d+ d+ d- < d+ d+ d- S 2 - d+ d+ d- d+

c. Dipole-dipole: weakest electrostatic force; exist between neutral polar molecules § Increase with increasing polarity (dipole moment) of molecule Ex: What IMF exist in Na. Cl (aq)? 5

6 d. Hydrogen bonds (or H-bonds): § H is unique among the elements because it has a single ethat is also a valence e-. – When this e- is “hogged” by a highly EN atom (a very polar covalent bond), the H nucleus is partially exposed and becomes attracted to an e--rich atom nearby.

§ H-bonds form with H-X • • • X', where X and X' have high EN and X' possesses a lone pair of e§ X = F, O, N (since most EN elements) on two molecules: F-H : F O-H : O N-H : N 7

8 § * There is no strict cutoff for the ability to form H-bonds (S forms a biologically important hydrogen bond in proteins). § * Hold DNA strands together in double-helix Nucleotide pairs form Hbonds DNA double helix

§ H-bonds explain why ice is less dense than water. 9

Ex: Boiling points of nonmetal hydrides § Polar molecules have higher BP than nonpolar molecules ∴ Polar molecules have stronger IMF Boiling Points (ºC) Conclusions: § BP increases with increasing MW ∴ Heavier molecules have stronger IMF § NH 3, H 2 O, and HF have unusually high BP. ∴ H-bonds are stronger than dipole-dipole IMF 10

2. Inductive forces: § Arise from distortion of the e- cloud induced by the electrical field produced by another particle or molecule nearby. 11 § London dispersion: between polar or nonpolar molecules or atoms – * Proposed by Fritz London in 1930 – Must exist because nonpolar molecules form liquids Fritz London (1900 -1954)

How they form: 12 1. Motion of e- creates an instantaneous dipole moment, making it “temporarily polar”. 2. Instantaneous dipole moment induces a dipole in an adjacent atom • * Persist for about 10 -14 or 10 -15 second Ex: two He atoms

* Geckos! 13 § Geckos’ feet make use of London dispersion forces to climb almost anything. § A gecko can hang on a glass surface using only one toe. § Researchers at Stanford University recently developed a gecko-like robot which uses synthetic setae to climb walls http: //en. wikipedia. org/wiki/Van_der_Waals%27_force

14 London dispersion forces increase with: § Increasing MW, # of e-, and # of atoms (increasing # of e- orbitals to be distorted) Boiling points: Effect of MW: Effect of # atoms: pentane 36ºC Ne – 246°C hexane 69ºC CH 4 – 162°C heptane 98ºC ? ? ? effect: H 2 O D 2 O § 100°C 101. 4°C “Longer” shapes (more likely to interact with other molecules) C 5 H 12 isomers: 2, 2 -dimethylpropane 10°C pentane 36°C

Summary of IMF Van der Waals forces 15

Ex: Identify all IMF present in a pure sample of each substance, then explain the boiling points. BP(⁰C) HF HCl HBr HI 20 -85 -67 -35 16 IMF Explanation London, dipole, H-bonds Lowest MW/weakest London, but most polar/strongest dipole and has H-bonds London, dipole Low MW/weak London, moderate polarity/dipole-dipole and no H-bonds London, dipole Medium MW/medium London, moderate polarity/dipole-dipole and no H-bonds London, dipole Highest MW/strongest London, but least polar bond/weakest dipole-dipole and no H-bonds

11. 3: Properties resulting from IMF 1. Viscosity: resistance of a liquid to flow Viscosity depends on: -the attractive forces between molecules -the tendency of molecules to become entangled -the temperature 17

11. 3: Properties resulting from IMF • Surface tension: energy required to increase the surface area of a liquid 18

3. Cohesion: attraction of molecules for other molecules of the same compound 4. Adhesion: attraction of molecules for a surface 19

20 5. Meniscus: curved upper surface of a liquid in a container; a relative measure of adhesive and cohesive forces Ex: Hg (cohesion rules) H 2 O (adhesion rules)

Phase Changes • Surface molecules are only attracted inwards towards the bulk molecules. • Sublimation: solid gas. • Vaporization: liquid gas. • Melting or fusion: solid liquid. • Deposition: gas solid. • Condensation: gas liquid. • Freezing: liquid solid. Energy Changes Accompanying Phase Changes • Energy change of the system for the above processes are: 21

Intermolecular Forces Bulk and Surface 22

Phase Changes Energy Changes Accompanying Phase Changes – – – Sublimation: Hsub > 0 (endothermic). Vaporization: Hvap > 0 (endothermic). Melting or Fusion: Hfus > 0 (endothermic). Deposition: Hdep < 0 (exothermic). Condensation: Hcon < 0 (exothermic). Freezing: Hfre < 0 (exothermic). • Generally heat of fusion (enthalpy of fusion) is less than heat of vaporization: – it takes more energy to completely separate molecules, than partially separate them. 23

Phase Changes Energy Changes Accompanying Phase Changes • All phase changes are possible under the right conditions (e. g. water sublimes when snow disappears without forming puddles). • The sequence heat solid melt heat liquid boil heat gas is endothermic. • The sequence cool gas condense cool liquid freeze cool solid is exothermic. 24

Phase Changes Energy Changes Accompanying Phase Changes 25

Phase Changes Heating Curves • Plot of temperature change versus heat added is a heating curve. • During a phase change, adding heat causes no temperature change. – These points are used to calculate Hfus and Hvap. • Supercooling: When a liquid is cooled below its melting point and it still remains a liquid. • Achieved by keeping the temperature low and increasing kinetic energy to break intermolecular forces. 26

Phase Changes Heating Curves 27

Heating Curve Illustrated 28

Phase Changes Critical Temperature and Pressure • Gases liquefied by increasing pressure at some temperature. • Critical temperature: the minimum temperature for liquefaction of a gas using pressure. • Critical pressure: pressure required for liquefaction. 29

Critical Temperature, Tc 30

Transition to Supercritical CO 2 31

Supercritical CO 2 Used to Decaffeinate Coffee 32

Vapor Pressure Explaining Vapor Pressure on the Molecular Level • Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid. • These molecules move into the gas phase. • As the number of molecules in the gas phase increases, some of the gas phase molecules strike the surface and return to the liquid. • After some time the pressure of the gas will be constant at the vapor pressure. 33

Gas-Liquid Equilibration 34

Vapor Pressure Explaining Vapor Pressure on the Molecular Level • Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface. • Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium. 35

Vapor Pressure Volatility, Vapor Pressure, and Temperature • If equilibrium is never established then the liquid evaporates. • Volatile substances evaporate rapidly. • The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates. 36

Liquid Evaporates when no Equilibrium is Established 37

Vapor Pressure Volatility, Vapor Pressure, and Temperature 38

Vapor Pressure and Boiling Point • Liquids boil when the external pressure equals the vapor pressure. • Temperature of boiling point increases as pressure increases. • Two ways to get a liquid to boil: increase temperature or decrease pressure. – Pressure cookers operate at high pressure. At high pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked, reducing the cooking time required. • Normal boiling point is the boiling point at 760 mm. Hg (1 atm). 39

Phase Diagrams • Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases. • Given a temperature and pressure, phase diagrams tell us which phase will exist. • Features of a phase diagram: – Triple point: temperature and pressure at which all three phases are in equilibrium. – Vapor-pressure curve: generally as pressure increases, temperature increases. – Critical point: critical temperature and pressure for the gas. – Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid. – Normal melting point: melting point at 1 atm. 40

Phase Diagrams • Any temperature and pressure combination not on a curve represents a single phase. 41

Phase Diagrams The Phase Diagrams of H 2 O and CO 2 • Water: – The melting point curve slopes to the left because ice is less dense than water. – Triple point occurs at 0. 0098 C and 4. 58 mm. Hg. – Normal melting (freezing) point is 0 C. – Normal boiling point is 100 C. – Critical point is 374 C and 218 atm. • Carbon Dioxide: – Triple point occurs at -56. 4 C and 5. 11 atm. – Normal sublimation point is -78. 5 C. (At 1 atm CO 2 sublimes it does not melt. ) – Critical point occurs at 31. 1 C and 73 atm. 42

Phase Diagrams The Phase Diagrams of H 2 O and CO 2 43

11. 4: Phase Changes Processes: § Endothermic: melting, vaporization, sublimation § Exothermic: condensation, freezing, deposition I 2 (s) and (g) 44 Microchip

Water: Enthalpy diagram or heating curve 45

11. 5: Vapor pressure 46 Pressure cooker ≈ 2 atm Normal BP = 1 atm 10, 000’ elev ≈ 0. 7 atm 29, 029’ elev (Mt. Everest) ≈ 0. 3 atm § A liquid will boil when the vapor pressure equals the atmospheric pressure, at any T above the triple point.

11. 6: Phase diagrams: CO 2 47 § Lines: 2 phases exist in equilibrium § Triple point: all 3 phases exist together in equilibrium (X on graph) § Critical point, or critical temperature & pressure: highest T and P at which a liquid can exist (Z on graph) Temp (ºC) § For most substances, inc P will cause a gas to condense (or deposit), a liquid to freeze, and a solid to become more dense (to a limit. )

Phase diagrams: H 2 O • For H 2 O, inc P will cause ice to melt. 48

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11. 7 -8: Structures of solids § Amorphous: without orderly structure Ex: rubber, glass § Crystalline: repeating structure; have many different stacking patterns based on chemical formula, atomic or ionic sizes, and bonding 51

Cubic Unit Cells in Crystalline Solids • Primitive-cubic shared atoms are located only at each of the corners. 1 atom per unit cell. • Body-centered cubic 1 atom in center and the corner atoms give a net of 2 atoms per unit cell. • Face-centered cubic corner atoms plus half-atoms in each face give 4 atoms per unit cell. 52

Common Lattice Structures 53

Types of Crystalline Solids Type Atomic Particles Atoms 54 Forces Notable properties Examples London dispersion §Poor conductors §Very low MP Ar (s), Kr (s) A small (~2 cm long) piece of rapidly melting argon ice (the liquid is flowing off at the bottom) which has been frozen by allowing a slow stream of the gas to flow into a small graduated cylinder which was immersed into a cup of liquid nitrogen

Molecular crystals Molecules (polar or non-polar) Carbon dioxide, dry ice (g at room T) London dispersion, dipole-dipole, H -bonds §Poor conductors §Low to moderate MP Ice (liq at room T) SO 2(s) 55 CO 2 (s), C 12 H 22 O 11, H 2 O (s) Sucrose (liq at room T)

Covalent (a. k. a. covalent network) Graphite Atoms bonded in a covalent network §Very hard §Very high MP Covalent §Generally bonds insoluble §Variable conductivity Diamond C (diamond &56 graphite) Si. O 2 (quartz) Ge, Si. C, BN Si. O 2

Ionic Anions and cations Crystals shatter! Ion-ion (ionic bonding) High Lattice Energy §Hard & brittle §High MP, BP §Poor conductors §Some solubility in H 2 O 57 Na. Cl, Ca(NO 3)2

Metallic §Excellent Metallic conductors Metal cations bonds in a diffuse, Usually face §Malleable delocalized e -centered §Ductile - cloud or body §High but wide centered range of MP 58 Cu, Al, Fe (hard) Alloys Pb, Au, Na (soft)

Overall 59 • Physical properties depend on these forces. The stronger the forces between the particles, • (a) the higher the melting point. • (b) the higher the boiling point. • (c) the lower the vapor pressure (partial pressure of vapor in equilibrium with liquid or solid in a closed container at a fixed temperature). • (d) the higher the viscosity (resistance to flow). • (e) the greater the surface tension (resistance to an increase in surface area).

Practice • Determine the type of solid and the forces holding the particles together • • Si. O 2 Covalent Network Covalent Bonds Na. NO 3 Ionic Electrostatic Att. C 2 H 6 Molecular Dispersion CH 3 OH Molecular Dispersion, Dipole-Dipole, H-Bond C(diamond) Covalent Network Covalent Bonds Al Metallic Kr Atomic (Molecular) Dispersion H 2 O Molecular Dispersion, Dipole-Dipole, H-Bond 60

Extra Material • The following pages contain some additional material and review items 61

Examples 62

Ionic Solids stable, high melting points Ø held together by strong electrostatic forces between oppositely charged ions Ø larger ions are arranged in closest packing arrangement Ø smaller ions fit in the holes created by the larger ions Ø 63

Cubic Unit Cells in Crystalline Solids • Primitive-cubic shared atoms are located only at each of the corners. 1 atom per unit cell. • Body-centered cubic 1 atom in center and the corner atoms give a net of 2 atoms per unit cell. • Face-centered cubic corner atoms plus half-atoms in each face give 4 atoms per unit cell. 8– 64 64

Common Lattice Structures 65

Calculations involving the Unit Cell • • The density of a metal can be calculated if we know the length of the side of a unit cell. The radius of an metal atom can be determined if the unit cell type and the density of the metal known – Relationship between length of side and radius of atom: • Primitive 2 r = l; FCC: BCC E. g. Polonium crystallizes according to the primitive cubic structure. Determine its density if the atomic radius is 167 pm. E. g. 2 Calculate the radius of potassium if its density is 0. 8560 g/cm 3 and it has a BCC crystal structure. 8– 66 66

Figure 11. 31 • Length of sides a, b, and c as well as angles a, b, g vary to give most of the unit cells. Return to unit cells 8– 67 67

Unit Cells in Crystalline Solids • Metal crystals made up of atoms in regular arrays – the smallest of repeating array of atoms is called the unit cell. • There are 14 different unit cells that are observed which vary in terms of the angles between atoms some are 90°, but others are not. Go to Figure 11. 31 8– 68 68

Packing of Spheres and the Structures of Metals • • • Arrays of atoms act as if they are spheres. Two or more layers produce 3 -D structure. Angles between groups of atoms can be 90° or can be in a more compact arrangement such as the hexagonal closest pack (see below) where the spheres form hexagons. Two cubic arrays one directly on top of the other produces simple cubic (primitive) structure. – Each atom has 6 nearest neighbors (coordination number of 6); nearest neighbor is where an atom touches another atom. – 54% of the space in a cube is used. • Offset layers produces a-b-a-b arrangement since it takes two layers to define arrangement of atoms. – BCC structure an example. – Coordination # is 8. 8– 69 69

Packing of Spheres and the Structures of Metals • FCC structure has a-b-c-a-b-c stacking. It takes three layers to establish the repeating pattern and has 4 atoms per unit cell and the coordination number is 12. 8– 70 70

Metallic Crystals Ø can be viewed as metals atoms (spheres) packed together in the closest arrangement possible Ø closest packing- when each sphere has 12 neighbors l l l 6 on the same plane 3 in the plane above 3 in the plane below 71

Bonding of Metals Ø the highest energy level for most metal atoms does not contain many electrons Ø these vacant overlapping orbitals allow outer electrons to roam freely around the entire metal 72

Bonding of Metals Ø these roaming electrons form a sea of electrons around the metal atoms Ø malleability and ductility l l bonding is the same in every direction one layer of atoms can slide past another without friction Ø conductivity l from the freedom of electrons to move around the atoms 73

Metal Alloys substance that is a mixture of elements and has metallic properties Ø substitutional alloy Ø l l Ø host metal atoms are replaced by other metal atoms happens when they have similar sizes interstitial alloy l l metal atoms occupy spaces created between host metal atoms happens when metal atoms have large difference in size 74

Examples Ø Brass l l substitutional 1/3 of Cu atoms replaced by Zn Ø Steel l interstitial Fe with C atoms in between makes harder and less malleable 75

Chapter 11 Overview • Changes of State – Phase transitions – Phase Diagrams • Liquid State Exam on Friday We will begin Chp 14 Thursday – Properties of Liquids; Surface tension and viscosity – Intermolecular forces; explaining liquid properties • Solid State – – – Classification of Solids by Type of Attraction between Units Crystalline solids; crystal lattices and unit cells Structures of some crystalline solids Calculations Involving Unit-Cell Dimensions Determining the Crystal Structure by X-ray Diffraction 8– 76 76

Comparison of Gases, Liquids and Solids – Gases are compressible fluids. Their molecules are widely separated. – Liquids are relatively incompressible fluids. Their molecules are more tightly packed. – Solids are nearly incompressible and rigid. Their molecules or ions are in close contact and do not move. 8– 77 Figure 11. 2 States of Matter 77

Phase Transitions • Melting: change of a solid to a liquid. H 2 O(s) H 2 O(l) • Freezing: change a liquid to a solid. H 2 O(l) H 2 O(s) • Vaporization: change of a solid or liquid to a gas. Change of solid to vapor often called sublimation. H 2 O(l) H 2 O(g) or H 2 O(s) H 2 O(g) • Condensation: change of a gas to a H 2 O(g) H 2 O(l) liquid or solid. Change of a gas to a or H 2 O(g) H 2 O(s) solid often called deposition. 8– 78 78

Vapor Pressure • In a sealed container, some of a liquid evaporates to establish a pressure in the vapor phase. • Vapor pressure: partial pressure of the vapor over the liquid measured at equilibrium and at some temperature. • Dynamic equilibrium 8– 79 79

Temperature Dependence of Vapor Pressures • The vapor pressure above the liquid varies exponentially with changes in the temperature. • The Clausius-Clapeyron equation shows how the vapor pressure and temperature are related. It can be written as: 8– 80 80

Clausius – Clapeyron Equation • • • A straight line plot results when ln P vs. 1/T is plotted and has a slope of Hvap/R. Clausius – Clapeyron equation is true for any two pairs of points. Write the equation for each and combine to get: 8– 81 81

Using the Clausius – Clapeyron Equation • Boiling point the temperature at which the vapor pressure of a liquid is equal to the pressure of the external atmosphere. • Normal boiling point the temperature at which the vapor pressure of a liquid is equal to atmospheric pressure (1 atm). E. g. Determine normal boiling point of chloroform if its heat of vaporization is 31. 4 k. J/mol and it has a vapor pressure of 190. 0 mm. Hg at 25. 0°C. E. g. 2. The normal boiling point of benzene is 80. 1°C; at 26. 1°C it has a vapor pressure of 100. 0 mm. Hg. What is the heat of 8– 82 vaporization? 82

Energy of Heat and Phase Change • Heat of vaporization: heat needed for the vaporization of a liquid. H 2 O(l) H 2 O(g) H = 40. 7 k. J • Heat of fusion: heat needed for the melting of a solid. H 2 O(s) H 2 O(l) H = 6. 01 k. J • Temperature does not change during the change from one phase to another. E. g. Start with a solution consisting of 50. 0 g of H 2 O(s) and 50. 0 g of H 2 O(l) at 0°C. Determine the heat required to heat this mixture to 8– 83 100. 0°C and evaporate half of the water. 83

Phase Diagrams • Graph of pressure-temperature relationship; describes when 1, 2, 3 or more phases are present and/or in equilibrium with each other. • Lines indicate equilibrium state two phases. • Triple point- Temp. and press. where all three phases co-exist in equilibrium. • Critical temp. - Temp. where substance must always be gas, no matter what pressure. • Critical pressure- vapor pressure at critical temp. • Critical point- point where system is at its critical pressure and temp. 8– 84 84

Properties of Liquids • Surface tension: the energy required to increase the surface area of a liquid by a unit amount. • Viscosity: a measure of a liquid’s resistance to flow. • Surface tension: The net pull toward the interior of the liquid makes the surface tend to as small a surface area as possible and a substance does not penetrate it easily. • Viscosity: Related to mobility of a molecule (proportional to the size and types of interactions in the liquid). – Viscosity decreases as the temperature increases since increased temperatures tend to cause increased mobility of the molecule. 8– 85 85

Intermolecular Forces • Intermolecular forces: attractions and repulsions between molecules that hold them together. • Intermolecular forces (van der Waals forces) hold molecules together in liquid and solid phases. – Ion-dipole force: interaction between an ion and partial charges in a polar molecule. – Dipole-dipole force: attractive force between polar molecules with positive end of one molecule is aligned with negative side of other. – London dispersion Forces: interactions between instantaneously formed electric dipoles on neighboring polar or nonpolar molecules. – Polarizability: ease with which electron cloud of some substance can be distorted by presence of some electric field (such as another dipolar substance). Related to size of atom or molecule. Small atoms and molecules less easily polarized. 8– 86 86

Boiling Points vs. Molecular Weight • Hydrogen bonds: the interaction between hydrogen bound to an electronegative element (N, O, or F) and an electron pair from another electronegative element. Hydrogen bonding is the dominate force holding the two DNA molecules together to form the double helix configuration of DNA. 8– 87 87

Comparisonof Energies for Intermolecular Forces Interaction Forces : Approximate Energy Intermolecular London 1 – 10 k. J Dipole-dipole 3 – 4 k. J Ion-dipole 5 – 50 k. J Hydrogen bonding 10– 40 k. J Chemical bonding Ionic 100 – 1000 k. J Covalent 100 – 1000 k. J 8– 88 88

Structure of Solids • Types of solids: – Crystalline – a well defined arrangement of atoms; this arrangement is often seen on a macroscopic level. • Ionic solids – ionic bonds hold the solids in a regular three dimensional arrangement. • Molecular solid – solids like ice that are held together by intermolecular forces. • Covalent network – a solid consists of atoms held together in large networks or chains by covalent networks. • Metallic – similar to covalent network except with metals. Provides high conductivity. – Amorphous – atoms are randomly arranged. No order exists 8– 89 in the solid. Example: glass 89