1 Intermolecular Forces 11. 2 INTERMOLECULAR FORCES Van der Waals’ forces Hydrogen bonds Dipole-dipole forces London Dispersion forces

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  • 1 Intermolecular Forces 11
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  • 2 INTERMOLECULAR FORCES Van der Waals forces Hydrogen bonds Dipole-dipole forces London Dispersion forces
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  • 3 Johannes van der Waals (1837 1923). Fritz London (1900 1954).
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  • 4 3 types of dipoles Permanent dipole Permanent dipole Instantaneous dipole Instantaneous dipole Induced dipole Induced dipole
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  • 5 Permanent dipole A permanent dipole exists in all polar molecules as a result of the difference in the electronegativity of bonded atoms.
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  • 6 Instantaneous dipole An instantaneous dipole is a temporary dipole that exists as a result of fluctuation in the electron cloud.
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  • 7 Instantaneous dipole An instantaneous dipole is a temporary dipole that exists as a result of fluctuation in the electron cloud.
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  • 8 Induced dipole An induced dipole is a temporary dipole that is created due to the influence of neighbouring dipole (which may be a permanent or an instantaneous dipole). Permanent dipole
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  • 9 11.2 Van der Waals Forces
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  • 10 Van der Waals Forces Van der Waals forces Dipole- Dipole Interaction Dipole- Dipole Interaction Dipole- Induced Dipole Interaction Dipole- Induced Dipole Interaction Instantaneous Dipole- Induced Dipole Interaction Instantaneous Dipole- Induced Dipole Interaction London dispersion forces
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  • 11 Dipole-dipole interactions Electrostatic interactions between polar molecules
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  • 12 Dipole-dipole interactions In a sample containing many polar molecules A balance of attraction and repulsion holding the molecules together
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  • 13 Dipole-induced dipole interactions When a non-polar molecule approaches a polar molecule (with a permanent dipole), a dipole will be induced in the non-polar molecule. Dispersion forces exist among all molecules and contribute most to the overall van der Waals forces.
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  • 14 Polarizability : - A measure of how easily the electron cloud of an atom/molecule can be distorted to induce a dipole Polarization
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  • 15 In general, size of electron cloud electron cloud is less controlled by positive nuclei extent of electron cloud distortion polarizability stronger dispersion forces
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  • 16 Instantaneous dipole-induced dipole interactions 11.2 Van der Waals forces (SB p.277) The instantaneous dipole arises from constant movement of electrons. Induces dipoles in neighbouring atoms or molecules
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  • 18 Instantaneous dipole-induced dipole interactions
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  • 20 Evidence for the presence of London dispersion forces 1. Condensation of noble gases at low temperatures to form liquids and solids presence of attractive forces between non-polar atoms E.g. Xe(g) Xe(s) H sub = -14.9 kJ mol 1
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  • 21 Evidence for the presence of London dispersion forces 2.The non-ideal behaviour of gases van der Waals equation
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  • 22 Strength of van der Waals forces 11.2 Van der Waals forces (SB p.279) Much weaker than covalent bonds Less than 10% the strength of covalent bonds van der Waals radius > covalent radius I2I2
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  • 23 Q.59 The electron clouds of adjacent iodine molecules would repel each other strongly until the equilibrium van der Waals distance is restored.
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  • 24 The strength of van der Waals forces can be estimated by melting point, boiling point, enthalpy change of fusion or enthalpy change of vapourization. Higher m.p./b.p./ H fusion / H vap stronger van der Waals forces
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  • 25 Strength of van der Waals forces Depends on three factors (in decreasing order of importance) : - 1. Size of molecule 2. Surface area of molecule 3. Polarity of molecule
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  • 26 Size of molecule Size of electron cloud MoleculeBoiling point ( o C) Helium Neon Argon -269 -246 -186 Fluorine Chlorine Bromine -188 -34.7 58.8 Methane Ethane Propane -162 -88.6 -42.2 1. Size of Molecule Polarizability Dispersion forces Rel. molecular mass Sometimes !
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  • 27 The van der Waals forces also increase with the surface area of the molecule. 2. Surface area of molecule van der Waals' forces are short-ranged forces Atoms or molecules must come close together for significant induction of dipoles.
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  • 28 Pentane (C 5 H 12 ) 2,2-dimethylpropane (C 5 H 12 ) Boiling point: 36.1C Boiling point: 9.5C Both are non-polar Same no. of electrons
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  • 29 2,2-dimethylpropane molecules pentane molecules larger contact area smaller contact area rod-shapedspherical in shape
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  • 30 Pentane (C 5 H 12 ) Larger contact surface area Higher chance of forming induced dipoles stronger dispersion forces Boiling point = 36.1 C
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  • 31 2,2-dimethylpropane (C 5 H 12 ) 2,2-dimethylpropane (C 5 H 12 ) Smaller contact surface area lower chance of forming induced dipoles weaker dispersion forces Boiling point = 9.5 C
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  • 32 3. Polarity of molecules For molecules with comparable molecular sizes and shapes, dispersion forces are approximately equal. Polar/polar > polar/non-polar > non-polar/non-polar Then, strength of van der Waals forces depends on the polarity of molecules involved
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  • 33 RMM = 58.0, ++ ++ Dipole-dipole forces + Dispersion forces Dispersion forces only b.p. = 50 C b.p. = 0 C
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  • 34 Other examples : - 1.Graphite layers of large surface area strong van der Waals forces 2.Polyethene vs ethene (m.p. > 100 C) (m.p. = 169 C)
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  • 35 Molecule % contribution to the overall van der Waals' forces Dipole- dipole interaction Dipole- induced dipole interaction Instantaneous dipole- induced dipole interaction C 4 H 10 00100 HCl15481
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  • 36 Q.60(a) CH 3 Cl < CH 3 Br < CH 3 I b.p./ C -24.2 3.56 42.4 The strength of dispersion forces increases with molecular size/mass. Thus, b.p. increases with molecular size/mass Although chloromethane is more polar, the effect of dispersion forces outweights that of dipole-dipole forces.
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  • 37 Q.60(b)
  • 38 Q.60(c) F 2 Cl 2 ClFCH 2 Cl 2 F 2 < ClF < Cl 2 F 2. It is because 1.ClF has a greater molecular size than F 2 and thus has stronger dispersion forces than F 2 2. ClF is polar and its molecules are held by both dipole-dipole forces and dispersion forces.
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  • 39 Q.60(c) F 2 CH 2 Cl 2 -188 C -100 C -34.0 C 39.6 C Cl 2 > ClF. It is because 1.Cl 2 has a greater molecular size than ClF and thus has stronger dispersion forces than ClF. 2.Although ClF is polar, the effect of dispersion forces outweights that of dipole-dipole forces.
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  • 40 Q.60(c) F 2 CH 2 Cl 2 -188 C -100 C -34.0 C 39.6 C CH 2 Cl 2 > Cl 2. It is because 1.CH 2 Cl 2 has a greater molecular size than Cl 2 and thus has stronger dispersion forces than Cl 2. 2.CH 2 Cl 2 is polar and its molecules are held by both dipole-dipole forces and dispersion forces.
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  • 41 Q.60(d) NO < C 2 H 6 RMM 28.0 28.0 b.p./ C -151 -89
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  • 42 1 pm = 0.001 nm 1 nm = 10 9 m
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  • 43 C 2 H 6 > NO. It is because 1.C 2 H 6 has a greater molecular size and contact surface area than NO and thus has stronger dispersion forces than NO. 2.Although NO is polar, the effect of dispersion forces outweights that of dipole-dipole forces. NO < C 2 H 6 RMM 28.0 28.0 b.p./ C -151 -89
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  • 44 The melting of a solid involves the separation of molecules from a regularly packed molecular crystal. Thus, m.p. of a solid depends on 1. The strength of van der Waals forces 2. Packing efficiency of molecules in the crystal lattice
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  • 45 Symmetry of molecule Packing efficiency m.p.
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  • 46 Q.61
  • 82 B.p. as molecular size (dispersion > dipole-dipole) However, H 2 O, HF and NH 3 have abnormally high b.p. There exist unusually strong dipole-dipole forces (H-bond) All are polar
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  • 83 Formation of hydrogen bonding When a hydrogen atom is directly bonded to a highly electronegative atom (e.g. fluorine, oxygen and nitrogen), a highly polar bond is formed. 2.1 4.03.53.0
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  • 84 Electrostatic attractions exist between this partial positive charge and the These attractions are called hydrogen bonds lone pair electrons on a highly electronegative atom (i.e. fluorine, oxygen or nitrogen) of another molecule.
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  • 85 hydrogen bond
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  • 86 Formation of hydrogen bonds between H 2 O molecules. hydrogen bond
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  • 87 Reasons f

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