PART TWO: Electrostatic Interactions

Transcription

1 II-1 1 PART TWO: Electrostatic Interactions In the first section of this course we were more concerned with structural aspects of molecules. In this section the emphasis is on bonding. Bonds in molecules are formed by the interactions between electrons. One way of probing the bonds is to look at the electrons in the bonds. We can do this by the process of ionization. Ionization is the basis of Photoelectron Spectroscopy. You will meet this again in (Structure and Spectroscopy). Ionization Energies (p. 73, DeKock & Gray and p , DeKock & Gray) When light (photons) of energy E = hν strikes a gas (or solid, or liquid) it can cause electrons to move from one orbital to another. this is called excitation

2 II-2 2 If the energy is high enough the electron can escape from the proton. this is called ionization For the hydrogen atom, the ionization process is: H(g) + E(=hν) H + (g) + e - ( 2 1 mev 2 ) E H = E nii - E ni The ejected electron can have a Kinetic energy of 2 1 mev 2. The Ionization Energy, IE, of an atom or molecule is the minimum energy required to remove an electron from the gaseous atom or molecule in its ground state. The ejected electron will then have zero kinetic energy.

3 II-3 3 Einstein: Correctly explained the Photoelectric effect. hν = IE mev 2 NOBEL PRIZE IN PHYSICS Einstein Photoelectric Law When Ionization occurs as a result of the interaction of photons with the molecule it is called photoionization. The resulting electrons are called photoelectrons. See DeKand G, Ch. 2 & Ch. 1, p. 16.

4 II-4 4 The Photoelectric Effect is a very important process in Physics and Chemistry. Basis of Photoelectron Spectroscopy (19-207, Analytical 2/3) To calculate the ionization energy for the H atom we must first recognize that when the atom is ionized its change in energy is the ionization energy. nii H + E H = IE = E nii - E ni ni ground State. If we know the photon energy and can measure the KE of the electron (this is photoelectron spectroscopy). We can measure the ionization energy. IE = hν mν 2 = E orb. This is known as Koopmans Theorem.

5 II-5 5 The experiment: simplified The energies of the photoelectrons can be measured by placing a negative voltage on the mesh grid. This will repel electrons unless they have enough kinetic energy to get through. This is over-simplified but it will do for now.

6 II-6 6 Ionization Energies are almost always given in electron volts: This is a useful unit of energy. The electron volt is the energy acquired by an electron when it is accelerated by a potential difference of 1 VOLT (See APPENDIX 1 DEK and G for More. The Ionization Energy for H is ev. (This is sometimes called a Rydberg) 1eV = kj mol -1 For Sodium Na(g) Na + (g) + e - IE 1 = ev In all atoms except hydrogen, further ionizations are possible (1 for each electron).

7 II-7 7 For example Li has 3 ionization energies from the configuration 1s 2 2s 1 : Li(g) Li + (g) + e - 2s IE 1 = ev Li + (g) Li 2+ (g) + e - 1s IE 2 = ev Li 2+ (g) Li 3+ (g) + e - 1s IE 3 = ev xt/li/econ.html

8 II-8 8 At the moment we can interpret this by saying that the two electrons in the 1s orbital are closer to the nucleus than the 2s electron and are harder to remove. Why is it harder to remove the last 1s electron than the first? We will need part 4 of this course to really understand this properly. We will return to this again.

9 II-9 9 Periodic Trends In Ionization Energies First ionization energies vary systematically through the periodic table. DeKock and Gray, p. 78

10 II General Points â In any row IE s as Z (Z = atomic number). Largest change He Xe: Smallest Li Rb. ã Across a period IE (e.g. Na Ar). (but some exceptions)

11 II Exceptions IE Be > B and IE N > O Atoms with half full e.g., N: 2s 2 2p 3 and full shells e.g., Be: 2s 2 have larger IE s than one would expect. why? B has a higher nuclear charge than Be.

12 II However in B, the outermost electron is in a 2p orbital and is less strongly bound. i.e., 2s 1 2s 2 2s 2 2p 1 Li Be B IE/eV Recall the difference between s and p orbitals. s p p electrons further from nucleus See part 4 of course.

13 II What About N And O? IE N > O Valence configurations are: N 2s 2 2p 1 2p 1 2p 1 x y z O 2s 2 2p 2 2p 1 2p x y 1 z this is the cause: The 2 electrons are close and repel each other and so, the electron is helped out by repulsion of the other p X electron. electon-electron repulsions are extremely important in many aspects of bonding as we will see during this course.

14 II Let s Look At The General Trends Lithium Neon (e.g. across) Increase in IE. due to a steady increase in the effective +ve nuclear charge. Li Rb (down) Gradual (but slight decrease) in IE. Electrons in Rb are further away from nucleus and are screened by the electrons closer to the nucleus. Screened: Don t feel full force of extra nuclear charge, because other electrons are in the way.

15 II Ionization Energies For Core Electrons (p. 79, DeKock & Gray) Core electrons are those situated close to the nucleus and not in the valence shell. So far we have measured how hard it is to remove a core electron after the valence electrons have been removed. Ionization energies can be measured for removing an electron from a neutral atom. Ionization of Lithium 1s. Li(1s 2 2s 1 ) Li + (1s 1 2s 1 ) IE = ev Now Look at Figure 2-3 DeKock & Gray

16 II Further Insights Li(1s 2 2s 1 ) Li + (1s 2 ) + e IE 1 = ev Li + (1s 2 ) Li 2+ (1s 1 ) + e IE 2 = ev Li(1s 2 2s 1 ) Li + (1s 1 2s 1 ) + e IE 1s = ev In the case of IE 1s. The energy is less than IE 2. Why is it harder to remove the 1s electron when the 2s is not there?? It is much more complicated but we must wait for Part Four (Quantum Chemistry)

17 II So far, we have discussed the basics of ionization energy. We will meet some more advanced material on this later.. Two key characteristics of an atom or molecule that are the highest occupied molecular orbital (HOMO) and the lowest unoccupied molecular orbital (LUMO). Together, these two orbitals are called the frontier orbitals. The HOMO can be found by locating the outer most orbital containing an electron. The LUMO then is the first orbital that does not contain an electron. See the diagram below: We will see later how these orbitals dictate the chemical reactivity of organic molecules

18 II Electron Affinity Chemical reactivity is largely controlled by electrons. We are trying to understand how strongly these are held by atoms. We have looked at IE s now EA s (Electron Affinities). Electron Affinity (p. 81, DeKock and Gray) The electron affinity (EA) of an atom is defined as the energy required to remove an electron from the negative ion. i.e. ion s (g) atom (g) + e - E = EA

19 II Be Very Careful With Signs if energy is required for A - (g) A(g) + e - EA is +ve if energy is released A - (g) A(g) + e EA is ve then A(g) is more stable than A - (g).

20 II Element Name/SymbolZ Electron Element Name/Symbol Z Electron Affinity Affinity ev ev Hydrogen H Ruthenium Ru Helium He Rhodium Rh Lithium Li Palladium Pd Beryllium Be Silver Ag Boron B Cadmium Cd 48 NA Carbon C Indium In Nitrogen N Tin Sn Oxygen O Antimony Sb Fluorine F Tellurium Te Neon Ne Iodine I Sodium Na Xenon Xe Magnesium Mg Cesium Cs Aluminum Al Barium Ba 56 NA Silicon Si Lanthanum La Phosphorus P Cerium Ce Sulfur S Praseodymium Pr Chlorine Cl Neodymium Nd Argon Ar Promethium Pm Potassium K Samarium Sm Calcium Ca Europium Eu Scandium Sc Titanium Ti Vanadium V 23 Chromium Cr Manganese Mn Iron Fe Cobalt Co Nickel Ni Copper Cu Zinc Zn 30 ~0 Gallium Ga Germanium Ge Arsenic As Selenium Se Bromine Br Krypton Kr 36 NA Rubidium Rb Strontium Sr Yttrium Y Zirconium Zr Niobium Nb Molybdenum Mo Technetium Tc

21 II General Trends halogens large electron affinities (~3.5 ev) completes closed shell s 2 p 6 config. closed shells small EA s He, Ne, Be, Mg, Zn closed shells or subshells give very stable electronic IE s and EA s: First Row

22 II Overall And now.. Electronegativity (pp , Dekock and Gray) Electron Affinity is a very useful thermodynamic quantity you will use it again.

23 II However electron distribution in molecules involves the question of the relative tendencies of each atom to acquire control of shared electrons. Molecular Property is usually discussed using Atomic property of Electronegativity. (symbol is χ ) power of an atom to attract electrons when part of a molecule My grandmother had a Boston Terrier that was obsessed with playing tug of war. You couldn't sit down in her house without having him push a nasty damp sock into your hand. If you were bored enough to accept it, he'd growl (ferociously, he thought) and pull. He was a little dog, though, and you could easily pick up the sock with dog still attached. He'd dangle like a little Christmas tree ornament until he got tired and let go. The contest was futile because my mass was about ten times his. If I'd been a Boston Terrier, the match would have been different. A tug-of-war also goes on between atoms involved in a chemical bond. The bonding electrons are the sock. The atom that can pull on the bonding electrons more strongly will get them. The winner is expected to be the atom with the higher effective nuclear charge.

24 II FOR EXAMPLE: Na Cl almost complete transfer of electron (ionic bond) Small X Large X ( + ve) ( _ ve) Mulliken Electronegativity Mulliken proposed that electronegativity was proportional to the sum of the IE and EA. EN = c(ie + EA) (c = proportionality constant.) IE ability of atom to hold electron EA ability of atom to attract an electron. Seems to make sense Cl : large IE and large EA large χ

25 II However this doesn t always work that well â EA s are not known that accurately. ã This can artificially give noble gases a large χ because of their large IE. Compare χ MULL for F & Ne EN = c(ie + EA) F Ne IE + EA = IE + EA = = = Since (IE Ne + EA Ne ) > (IE F + EA F ) χ Ne > χ F

26 II χ PAULING Pauling Electronegativity - Uses comparison of bond energies. - A 2 molecule and B 2 molecule versus AB molecule.

27 II Bond Energy would be average unless there is ionic bonding. Example HF Bond Energy HF 135 kcalmol -1 H Av = 62 kcalmol F 37 2 Extra bond energy assumed to be a consequence of ionic bonding because of electronegativity differences in H and F. Pauling then set the electronegativity difference as EN A EN B = /2 where = DE AB = [(DE )(DE )] 1/2 A 2 converts to ev B 2 D = Bond Dissociation Energy -1

28 II PAULING set χ F = 3.98 Pauling AND Mulliken, χ s agree fairly well. one way to convert them is χ p = 1.35 X 1/ Q: Fluorine is the most electronegative element and yet chlorine has a larger electron affinity: Why? Hint: think about the relative sizes. M

29 II Electronic Configurations We can now see that bonding is affected strongly by ideas of how electrons are held by atoms in molecules. We know that electrons are held in orbitals. Now we review this material and see how electrons affect each other. Configurations We will meet the mathematical basis for orbitals and what they look like in part four of this course. Now we merely state the results in a rather pictorial way. Atomic Quantum Numbers

30 II Each Atomic Orbital (AO) is defined by 3 Quantum Numbers n, R, m R n Energies R shape m R (orientation) subsets of R We will define these somewhat more accurately later. n: 1, 2, 3, 4, R : 0, 1,, n-1 m R : R, R-1, R-2,, -R R = 0, 1, 2, 3, 4, s p d f g

31 II s Subshell R = 0, m R = 0 p : R = 1 d : R = 2 m R = +1, 0, -1 (p orbitals) m R = +2, +1, 0, -1, -2 (d orbitals) So the number of orbitals in a particular shell is n 2 where n is the principal quantum number. Configurations For Atoms And Ions So we can write down the configuration for any atom or ion (or can we?). - In first year we based this on the building up or AUFBAU principle. - Method for determining the lowest energy configuration for an atom (ground state).

32 II Pauli exclusion principle: no more than 2 electrons may occupy a single orbital and if two occupy a single orbital their spins must be paired. However it is not quite as simple. Penetration And Shielding - orbital energies differ from those in hydrogen. - Electron in orbital experiences. â coulombic attraction of Nucleus. ã coulombic repulsion from other electrons. - we approximate: Electons experience a CENTRAL FIELD. - sum of field from nucleus and average of all electrons.

33 II reduces Nuclear Charge from true value (Ze) to effective nuclear charge Z eff e (for a particular electron) Shielding + These electrons are screened by Shielding of inner electrons. NUCLEUS Shielding Parameter Z eff = Z - σ

34 II If you like think of it this way: electrons in outer orbitals experience attraction from nucleus but repulsion from other electrons. Effective nuclear charges Z eff H He Z 1 2 1s Li Be B C N O F Ne Z s s p Na Mg Al Si P S Cl Ar Z s s p s p NOTE: ns electrons are less shielded than np electrons.

35 II Due to penetration s electrons get closer to nucleus than p electrons. - s electrons less shielded. Similar differences for the orbitals. s orbitals most penetrating followed by p, d, f. - However the ordering depends on the number of electrons present. - Penetration effects v. marked for 4s e s of K and Ca.

36 II Penetration effects are very marked for 4s electrons of K and Ca. Q: Why is 2s and 2p same energy on the y-axis (no electrons)

37 II Ground State Electron Configurations (DeKock and Gray, p ) In your reading of DeKock and Gray you will come across the Schrödinger equation. This is the mathematical basis of quantum mechanics. For now do not need the equation. We will cover it in PART FOUR. GROUND STATE for 1 st 5 elements. H He Li Be B 1s 1 1s 2 1s 2 2s 1 1s 2 2s 2 1s 2 2s 2 2p 1 What About Carbon? CARBON (Z = 6): Removed from chem* 2060 Fall s 2 2s 2 2p 2

38 II WAIT! - We have a problem. - where do we put the electron? Let s look at the 2p shell THESE ARE OF DIFFERENT ENERGY - choice can be made by using Hund s rule. - maximum no of parallel spins results in lowest e - - e - repulsion.

39 II Ground State can be 1s 2 2s 2 2p 1 1 x2p y. We will return to p 2 a bit later. electron-electron repulsions and the Aufbau principle (DeKock and Gray, p. 51) We have already seen that e-e repulsions affect. â Ionization Energies (See Li) ã Electron Affinities (See Cl & F) ä Ground State Configurations (e.g. C) We now look briefly at the ordering of 3d and 4s orbitals of transition metal ions.

40 II Transition Elements - Importance of d electrons, also see , (Bioinoganic Chemistry.)

41 II The electronic Configurations of the Elements. Electronic Configuration Z neutral +ve ion 1 H 1s^1-2 He 1s^2 1s^1 3 Li [He] 2s^1 1s^2 4 Be [He] 2s^2 [He] 2s^1 5 B [He] 2s^2 2p^1 [He] 2s^2 6 C [He] 2s^2 2p^2 [He] 2s^2 2p^1 7 N [He] 2s^2 2p^3 [He] 2s^2 2p^2 8 O [He] 2s^2 2p^4 [He] 2s^2 2p^3 9 F [He] 2s^2 2p^5 [He] 2s^2 2p^4 10 Ne [He] 2s^2 2p^6 [He] 2s^2 2p^5 11 Na [Ne] 3s^1 [He] 2s^2 2p^6 12 Mg [Ne] 3s^2 [Ne] 3s^1 13 Al [Ne] 3s^2 3p^1 [Ne] 3s^2 14 Si [Ne] 3s^2 3p^2 [Ne] 3s^2 3p^1 15 P [Ne] 3s^2 3p^3 [Ne] 3s^2 3p^2 16 S [Ne] 3s^2 3p^4 [Ne] 3s^2 3p^3 17 Cl [Ne] 3s^2 3p^5 [Ne] 3s^2 3p^4

42 II Ar [Ne] 3s^2 3p^6 Ne] 3s^2 3p^5 19 K [Ar] 4s^1 [Ne] 3s^2 3p^6 20 Ca [Ar] 4s^2 [Ar] 4s^1 21 Sc [Ar] 3d^1 4s^2 [Ar] 3d^1 4s^1 22 Ti [Ar] 3d^2 4s^2 [Ar] 3d^2 4s^1 23 V [Ar] 3d^3 4s^2 [Ar] 3d^4 24 Cr [Ar] 3d^5 4s^1 [Ar] 3d^5 25 Mn [Ar] 3d^5 4s^2 [Ar] 3d^5 4s^1 26 Fe [Ar] 3d^6 4s^2 [Ar] 3d^6 4s^1 27 Co [Ar] 3d^7 4s^2 [Ar] 3d^8 28 Ni [Ar] 3d^8 4s^2 [Ar] 3d^9 29 Cu [Ar] 3d^10 4s^1 [Ar] 3d^10 30 Zn [Ar] 3d^10 4s^2 [Ar] 3d^10 4s^1 31 Ga [Ar] 3d^10 4s^2 4p^1 [Ar] 3d^10 4s^2 32 Ge [Ar] 3d^10 4s^2 4p^2 [Ar] 3d^10 4s^2 4p^1 33 As [Ar] 3d^10 4s^2 4p^3 [Ar] 3d^10 4s^2 4p^2 34 Se [Ar] 3d^10 4s^2 4p^4 [Ar] 3d^10 4s^2 4p^3 35 Br [Ar] 3d^10 4s^2 4p^5 [Ar] 3d^10 4s^2 4p^4 36 Kr [Ar] 3d^10 4s^2 4p^6 [Ar] 3d^10 4s^2 4p^5 37 Rb [Kr] 5s^1 [Ar] 3d^10 4s^2 4p^6 38 Sr [Kr] 5s^2 [Kr] 5s^1 39 Y [Kr] 4d^1 5s^2 [Kr] 5s^2 40 Zr [Kr] 4d^2 5s^2 [Kr] 4d^2 5s^1 41 Nb [Kr] 4d^4 5s^1 [Kr] 4d^4 42 Mo [Kr] 4d^5 5s^1 [Kr] 4d^5 43 Tc [Kr] 4d^5 5s^2 [Kr] 4d^5 5s^1 44 Ru [Kr] 4d^7 5s^1 [Kr] 4d^7 45 Rh [Kr] 4d^8 5s^1 [Kr] 4d^8 46 Pd [Kr] 4d^10 [Kr] 4d^9 47 Ag [Kr] 4d^10 5s^1 [Kr] 4d^10 48 Cd [Kr] 4d^10 5s^2 [Kr] 4d^10 5s^1 49 In [Kr] 4d^10 5s^2 5p^1 [Kr] 4d^10 5s^2 50 Sn [Kr] 4d^10 5s^2 5p^2 [Kr] 4d^10 5s^2 5p^1 Here we see several exceptions to the Aufbau principle.

43 II We have seen (p. 89) that the filling rules that we learnt in first year chemistry are not quite accurate. The simple picture for filling orbitals depends on â atomic number (see p. 86) ã charge

44 II Let s Look At The 3d & 4s More Closely (P. 52, DeKock and Gray) 4s is lower in energy than 3d (so fills first) but: Chromium [Ar] 3d 5 4s 1 Copper [Ar] 3d 10 4s 1 Also Scandium [Ar] 3d 1 4s 2 s fills first Scandium 2+ [Ar] 3d 1 d fills first

45 II The situation is even stranger Ionization Energies Scandium has 1 d electron and 2s electrons in the valence orbitals. IE s for these are: Sc(3d 1 4s 2 ) Sc + (3d 1 4s 1 ) + e - Sc(3d 1 4s 2 ) Sc + (3d 0 4s 2 ) + e ev 7.98 ev - d orbital is more stable. - Then why does s shell fill first? - Larger e-e repulsion in d orbitals. (Less diffuse than 4s orbital) s electrons penetrate better. (Wait until part 4)

46 II So for most of the d block the ground state is 3d n 4s 2 (s fills first) because of large e-e repulsions in d orbitals. Note exceptions usually occur with ½ and full shells (see Chromium d 5 s 1 ) (see Copper d 10 s 1 ) Complications are not important when orbital energy of 3d is much less than 4s.

47 II Electrostatics Of Atoms & Molecules In the previous section we looked at Coulombic attractions and repulsions and their importance in electrons in atoms. We now look at coulombic attractions between atoms and molecules. Interacting Charges Interactions between atoms and molecules are electrical in nature. Classical Electrostatics can be used to predict energy of interaction. All based on simple coulombic law.

48 II Types Of Interaction a) Monopole Monopole b) Monopole Dipole c) Dipole Dipole a) Ion-Ion Bonds (Monopole Monopole) coulombic interaction = r For correct use and units see Prob. Set (Q s 1, 4) q+ q - 4πεr PLEASE NOTE 4πε - This gets units right. We leave it out in the notes.

49 II Note Signs: attractive has ve sign. -q q 1 r E = 2 note: 4 1 term left out επ for convenience Monopole Monopole - Non Directional - Strong over atomic distances. - Good model for ionic bonding e.g. LiF estimate 686 kj mol -1 (measured 755 kj mol -1 ) (see later) b) monopole dipole

50 II Monopole Z + Dipole R E r = -Zq + Zq = ZqR r - R r R r2 R attraction repulsion - Zµ r 2 i.e., Zq Zq r- R + r+ R 2 2 = -Zq(r + R) + Zq(r- ) (r- R 2 R 2 )(r+ R) 2 2 = -ZqR (but µ = r2 R2 4 qr) and if R is small compared to r : E = Zµ r 2

51 II net attraction: even though there is no charge on the dipole. Monopole orients the dipole in favourable direction. Example: Solvated Ions: When NaCl dissolves in water Na + is surrounded by H 2 O because of dipole-monopole interaction. In aqueous solution Na + is hydrated by an octahedron of water molecules. Note the negatively charged Cl - (from NaCl) does not participate in the hydration shell. δ H O δ δ H δ δ H H O O δ Na δ H δ δ O H δ δ O H δ H δ H δ δ H δ O H δ H δ

52 II Polar nature of water makes it an excellent solvent for ionic solids like NaCl. Energy needed to separate ions is provided by formation of hydrated ions (monopole-dipole interactions). Non polar solvents (e.g., gasoline) cannot form such strong bonds. So NaCl and other salts are insoluble in gasoline. Strengths Of Bonds Monopole Monopole ~ 400 kj mol -1 Monopole Dipole ~ 40 kj mol -1 Smaller Energy because of distance dependence.

53 II Soaps (and detergents) work because their structures combine in one molecule a hydrocarbon chain which is hydrophobic (rejects water) and lipophilic (attracts oily materials) with an end which is hydrophilic (attracts water) and lipophobic (rejects oily materials). The long "fatty" chains provide solubility in hydrocarbons (grease) and the polar, usually ionic, heads provide solubility in water. If both oily and watery materials are present, a soap provides a "bridge" by dissolving its hydrocarbon chain in a droplet of oil in such a way that the ionic, hydrophilic, end sticks out into the surrounding water. This arrangement is called a micelle and permit soapy water to "wash away" greasy materials.

54 II The hydrophilic (polar) end of the detergent binds strongly to water the non-polar part bonds to the grease. c) dipole-dipole r R µ = qr

55 II E = -q r 2 + ATT. q2 + r2 R REP q2 r2 + R q r 2 ATT E r 3 q2r2 + rr 2 2 for R << r E µ 2 r3 Magnitude ~ 5 kj mol -1 VERY SHORT RANGE INTERACTION (AND WEAK) E = µ 2 for dipole-dipole is a special case. If r3 there is an angle θ between the dipoles then: -µ µ E = (1-3cos2 r3 2 1 θ)

56 II IMPORTANCE: Polar Liquids, Dissolving of Molecular Solids (see Later) + _ + + _ + + +_ _ + _ Energy Required + + Energy Released + _ SOLVATION

57 II Induced Dipoles & Van Der Waals Forces We have seen now that ionic type bonds can occur when a dipole moment exists in a molecule. A dipole can be set up in a non-polar molecule by an electric field. For example: induced dipole moment _ + _ + + _ _ voltage off Electrons attracted to r Nuclei attracted to s voltage on

58 II electric field can be from neighbouring atom or ion q + Monopole µ The magnitude of the induced dipole moment (µ) is related to the polarizability of the molecule. µ IND = α E Polarizability Strength of Electric field. α is the ease of deformation of the electron cloud around the molecule. It is a measure of how floppy the electrons are.

59 II note: interaction of induced dipole with another charge can only be an attraction since it is automatically created with the correct geometry. If species is atom no orienting effect. If species is molecule, α may be different in one direction. Molecule will tend to orient itself to create largest induced dipole moment. Polarizability And Electronegativity Polarizability increases with volume. Also depends on electronegativity. Electronegative atoms not very polarizable.

60 II Monopole-Induced Dipole E = - q α 2r (~ 5 kj mol -1 ) Dipole Induced Dipole E = -µ α r ~ 0.05 kj mol -1 Induced Dipole-Induced Dipole E = I I 1 2 I + I 1 2 α 1 α r6 2 I = ionization energy E ~ 0.5 kj mol -1

61 II In solids where there are no permanent dipoles (e.g. Xe, Kr, He, H 2, O 2, N 2 ) the solid is held together by weak forces. These are called Van der Waals Forces. Van Der Waals Forces (DeKock and Gray, p. 431) attractive and repulsive Induced Dipole Repulsion between 1 electrons on r 6 (London or dispersion forces) neighboring atoms

62 II Attractive Part Interaction (Potential) Energies that have a 1 r6 dependence (very short range) are usually lumped together as Van der Waals Forces. Repulsive Part Must be a repulsion (or everything would collapse into itself). Repulsion Energy = be -ar Importance Of Fluctuating Dipoles (Also Fig.7-13 DeK &G) F For the instant that this situation occurs there is an attraction between the instantaneous and induced dipoles. The effect is felt by both each induces a polarization in the other.

63 II The electron clouds repel each other at very small distances. - end up with a balance. One way of writing this is the Lennard-Jones 6-12 potential. E = C 1 r r 6 Repulsive Attractive Writing the repulsive part as an exponential is actually more realistic because of the exponential nature of radial wave functions (see part four) i.e. E = be -ar - d r6 (SEE DeKock and Gray, p. 432)

64 II IT HELPS TO LOOK AT THIS PICTORIALLY POTENTIAL ENERGY Equilibrium Internuclear Distance REPULSIVE req rab D AB TOTAL A B ATTRACTIVE r D AB = BOND ENERGY

65 II Actually well Depth for a crystal refers to the enthalpy of sublimation. solid gas Short & Long Range Forces Potential Energy Shortest Range r 12 1 r 6 1 r 4 r 2 1 Longest Range Distance

66 II Van Der Waals Radii Size of Molecule (or Atom) revealed by LJ potential is much larger than that revealed by other measures of atomic size (see PART 3) such as covalent radius or ionic radius. Because electron clouds stop interpenetration for a non-bonding interaction (or no electron sharing). Covalent and Ionic Bonding will be discussed later PARTS 3, 4 and 5.

67 II Van Der Waals Solids Good reading here: A comparison between a covalent molecular bond and a much weaker VDW bond in He

68 II molecular solids in which only Van der Waals intermolecular bonding exists generally melt at low temperatures. Because thermal energy is able to overcome VDW attraction very easily. Liquid and solid He exist only below 4.6 K. Van der Waals bonds get stronger as atoms get bigger (more polarizable).

69 II (see DeK&G 435 et seq) DW attractions responsible for liquid state.

70 II Hydrogen Bonding (DeKock & Gray, p. 436) Polar Molecules held together in molecular solids by dipoles. i.e. Opposite Ends of Dipole attract each other. A VERY IMPORTANT Kind of Polar Interaction is the Hydrogen Bond. - Relatively Weak ~ 20 kj mol -1 (Covalent & Ionic Bonds ~ 400 kj mol -1. EXAMPLES: δ + H δ_ HYDROGEN BOND θ δ + δ _ F H F 1.87D 0.92 D

71 II o 2.02 D O H O H o H 0.96 D H HYDROGEN BOND F H H N H H HYDROGEN BOND ALL GASEOUS DIMERS

72 II Features Common To H-Bonded Systems â Molecular Units Retain Their Integrity. e.g. H-F bond Lengths are same as in monomer. ã F a - - H b F b bond is linear. ä H atom is asymmetric (Only in very strong H bonds e.g. FHF - is this untrue). å Angle θ (see HF) is between 100 and 120. In solid HF the bonding is zig-zag. F F b H H a H b H F a F solid is held together by H-bonds

73 II ICE Each H 2 O molecule is bonded to 4 others (Tetrahedral). Although bonds are weak they are important. Hydrogen bonding in water is responsible for many of its important properties.

74 II Melting & Boiling Temperatures of water are unexpectedly high due to H bonding.

75 II Since H-bonding creates an open network ice is less dense than water. Only about 1/3 of H-bonds are broken when ice melts. In liquid phase water there are still H-bonds. As T clusters of H bonded water break up. volume continues to shrink. As T further thermal expansion occurs. - molecules need more room. (This then dominates over shrinkage caused by collapse of H-bonds). So liquid water has minimum volume (Max Density) at 4 C. (WHY Lakes Don t Freeze Solid) We will look at H-bonds again when we know more about bonding.

76 II HYDROGEN Bonds are unbelievably important in biochemistry. e.g. DNA Sequence and Replication A H H C N N H O CH 3 C C C C N C N H N C H H N C C N O H Thymine A H H C N C N C N O H N H C C C N H N C C C N N H O H H H Cytostine H One of the biggest sources of difficulty for a chemistry student is the distinction between chemical bonds and intermolecular forces. If you are having trouble try: html

77 II-77 77

78 II-78 78

79 II-79 79