INTERMOLECULAR FORCES: LIQUIDS, SOLIDS & PHASE CHANGES (Silberberg, Chapter 12)

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1 INTERMOLECULAR FORCES: LIQUIDS, SOLIDS & PASE CANGES (Silberberg, Chapter 12) Intermolecular interactions Ideal gas molecules act independently PV=nRT Real gas molecules attract/repulse one another 2 n a ( P )( V nb) nrt 2 V Liquid molecules stick to one another

2 Centrosymmetric formamide dimer at the crystal geometry. Blue shows regions of the molecular surface with a partial positive charge, red those regions with a partial negative charge and white regions which are neutral.

3 Types of Intermolecular interactions (intermolecular forces) Note: van der Waals forces is the generic term for ALL intermolecular forces. It is NOT interchangeable for dispersion forces. 1. London force ( dispersion force ) Electron cloud gives instantaneous dipole moment (e- build up at one end of molecule leaving the nucleus at the other end partially exposed) Molecules stick together because their partial charges attract one another London forces act between all types of molecules (polar and non-polar). Strength increases with molar mass : heavier molecules have more electrons (further from the nuclei) bigger fluctuations in electron movement (they are more polarizable) Eg. F 2 and Cl 2 are gases, Br 2 is a liquid, I 2 a solid at room temperature. Strength influenced by molecule shape Rod-shaped molecules have greater London forces of attraction because the instantaneous dipoles can get closer

4 2. Dipole-dipole interactions: Polar molecules have permanent partial charges. Polar molecules that are brought close to one another tend to orient their dipole moments so that the plus end of one molecule faces the minus end of another. This is a dipole-dipole interaction: Strength depends on magnitude of bond dipoles and shape of molecule. If bond dipoles cancel one another within a molecule, then the molecule itself has no dipole moment.

5 3. ydrogen bonding: Occurs when a hydrogen atom is bonded to a strongly electronegative atom with lone pairs (N, O or F) R-X- Y-R (X, Y = F, O or N) covalent bond Eg. water hydrogen bond O O O O C 3 C O O C C 3 N O O acetic acid amine N N ydrogen bonding is the strongest of the intermolecular interactions

6 In hexagonal ice I, (natural form of ice on Earth), each water molecule is -bonded to 4 neighbouring molecules in a tetrahedral arrangement. -bonds are less ordered at 20 C but they still account for the cohesiveness of liquid water where each molecule remains -bonded to an average of 3.5 neighbours (at that temperature water molecules exchange their positions about times per second!). Finally, the bonds weaken at higher temperatures and water vaporizes because thermal motion causes an uncertain random orientation of the water dipoles. Vapour is like a gas where molecules are too distant and move too fast to be able to interact.

7 Table of intermolecular forces and energies Type of interaction Intermolecular Typical energy (kj mol -1 ) Interacting species London (dispersion) Polarizable e - clouds (all atoms/molecules) dipole-dipole 5-25 dipole-dipole (polar molecules) ydrogen bonding Polar bond to dipole (molecules where is bonded to electronegative atom with lone pairs) Chemical bonding Ionic Cation-anion Covalent atoms sharing electrons in a bond

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9 Properties of liquids; influence of intermolecular interactions: Viscosity Viscosity = resistance to flow igh viscosity slow flow rate Strong intermolecular forces molecules held together can t move past one another easily high viscosity Viscosity as temperature because molecules have higher E k O vs But London forces can add up to quite significant degree: C 3 (C 2 ) 4 C=CC 2 C=C(C 2 ) 7 COO Linoleic acid C 3 (C 2 ) 7 C=C(C 2 ) 7 COO Oleic acid

10 Properties of liquids: Surface tension In a liquid, molecules experience equal forces from all directions except at the surface (experience a net inward force) intermolecular interactions occur in 3 dimensions int. interactions occur across & below surface net vector for attractive forces is down Strong intermolecular forces surface tension capillarity Water meniscus curves upward because forces between water molecules and the oxygen atoms and O groups in glass are stronger than the forces between water molecules. Mercury meniscus curves downwards because forces between mercury atoms are stronger than between g atoms and the glass

11 Phase transitions Phase transition Name Example Gas to liquid Condensation or liquefication Dew Liquid to gas Vaporisation Boiling water steam Gas to solid Condensation or deposition Frost Solid to gas Sublimation Evaporation of CO 2 Liquid to solid Freezing water ice Solid to liquid Melting or fusion ice water sublimination exothermic melting vaporizing solid liquid gas endothermic freezing condensing

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13 Vapour Pressure At a given T, molecules in liquid have a distribution of speeds. Some move fast enough to escape liquid (enter gas phase) At higher T, average kinetic energy is higher so greater fraction of molecules can escape liquid

14 Consider liquid in sealed container: Some fast moving molecules escape into gas phase These gas molecules exert pressure on the liquid surface The pressure as more molecules enter gas phase Some gas molecules collide with surface and stick to it, re-entering the liquid phase Eventually, rate of molecules evaporating equals rate of molecules condensing (equilibrium is reached) vapour pressure : partial pressure exerted by a vapour over a liquid at a fixed temperature Vapour pressure is characteristic for a given liquid (or solid) Function of the intermolecular forces Vapour pressure as temperature for a given liquid Example: Which would you expect to have a higher boiling point, p-dichlorobenzene or o-dichlorobenzene? Cl Cl Cl Cl

15 Vapour pressure as function of temperature and intermolecular forces Boiling point and melting point In open container, atmosphere exerts a pressure on the liquid surface. As T liquid molecules move more quickly & leave surface more often. At some temperature, KE is great enough for bubbles of vapour to appear in the liquid ( the liquid boils ). The boiling point is the temperature at which the vapour pressure equals the external pressure The normal boiling point, T b, of a liquid is defined as the boiling point at 1 atm external pressure Strong intermolecular forces low vapour pressure high boiling point The normal boiling point, T b, is defined as the boiling point at 1 atm external pressure The normal freezing point, T f, of a liquid is the temperature at which it freezes at 1 atm pressure (T f is only slightly dependent on pressure)

16 eat of phase transitions 1. Gas cools: 2. Gas condenses: 3. Liquid cools: 4. Liquid freezes: 5. Solid cools: Note: Within a phase, a change in heat is accompanied by a change in temperature (because of change in E k ). eat lost or gained depends on amount of substance, molar heat capacity for that phase and change in temperature. During a phase change, a change in heat occurs at constant temperature (energy needed to cause change in phase) Melting or boiling is endothermic; Condensation or freezing is exothermic.

17 Phase diagrams summarise phases of a substance under different conditions of temperature and pressure Phase diagram of CO 2 : solid liquid critical point gas triple point

18 Phase diagram of 2 O:

19 The Solid State (Silberberg, Chapter 12) Amorphous: atoms/molecules in random arrangement ( frozen liquid ) Crystalline: atoms/molecules in ordered pattern. Often have flat faces and definite angles at edges (formed by orderly stacks of atoms) Amorphous and crystalline silicon dioxide. When molten silica cools quickly it becomes a glass. The atoms (red=o; black=si) are arranged in a disorderly fashion. Quartz is a crystalline form of silica, SiO 2. The atoms (red=o; black=si) are arranged in an orderly network.

20 Crystalline solids can be classified in terms of their bonding: Class Examples Characteristics metallic s- and d- block elements malleable, ductile, lustrous, electrically and thermally conducting ionic NaCl, CsCl, KNO 3, CuSO O network (covalent bonding) molecular (covalent bonding; intermolecular interactions) hard, rigid, brittle, high melting points, those soluble in water conduct electrical current; when molten they are electrical conductors B, C, BN, SiO 2 hard, rigid, brittle, very high melting points, insoluble in water 2 O (ice), S 8, I 2, glucose, sucrose, naphthalene relatively low melting and boiling points, brittle if pure

21 The crystal lattice and unit cell lattice point unit cell All particles in a crystal are packed in an orderly way Imagine placing a point at each repeating part of the pattern the points form a regular array called a crystal lattice All lattice points are identical portion of a 3-D lattice The smallest part of the crystal that, if repeated in 3-D gives the whole crystal, is called the unit cell

22 Seven crystal systems (different shapes of unit cells) and fourteen kinds of unit cells exist in nature We consider only the cubic crystal system (all sides are equal length (called a) and all angles 90 ) In the cubic system there are 3 possible packing arrangements: Simple cubic / Primitive cubic (P) atom at each of the 8 corners of cube Body-centred cubic (I) atom at each corner & atom at centre of cube Face-centred cubic (F) atom at each corner & atom in centre of each face of cube

23 Primitive cubic Body-centred cubic Face-centred cubic Coordination number = 6 Coordination number = 8 Coordination number = atom at 8 corners 1 1 atom 8 8 atom at 8 corners at 8 corners 1 atom at center 1 2 atom at 6 faces 1 Atoms/unit cell = x 8 8 = 1 Atoms/unit cell = ( 1 x 8) + 1 = 2 8 Atoms/unit cell = ( 1 x 8) + ( 1 x 6)=4 8 2

24 Metals cations (identical spheres) packed tightly together and surrounded by a sea of electrons denser than most solids malleable and ductile conduct heat and electricity many pack in cubic unit cells Copper: face-centred cubic

25 Figure Packing of spheres. simple cubic (52% packing efficiency) body-centered cubic (68% packing efficiency)

26 Figure (continued) layer a layer a hexagonal closest packing layer b cubic closest packing layer c closest packing of first and second layers abab (74%) abcabc (74%) hexagonal unit cell expanded side views face-centered unit cell

27 Ionic solids formed from oppositely charged ions (overall crystal is neutral) held together by attraction of cations and anions hard, rigid, brittle conduct electricity when molten or in solution; not as solids anions close-pack and cations fit into interstitial spaces expanded view space-filling

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29 Allotropes (polymorphs) of Carbon Diamond and graphite are examples of Network solids atoms covalently bonded entire crystal is one network hard, brittle, high melting points

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31 Molecular crystal structures of carbon individual molecules, held in crystal structure by intermolecular forces only physical properties depend on types & strengths of intermolecular interactions relatively low melting points ( C) C 60 Face centred cubic

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33 The future? Manipulating atoms. nanotube gear

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