3-1 Lewis Electron-Dot Diagram

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2 3-1 Lewis Electron-Dot Diagram Lewis-dot diagrams, are diagrams that show the bonding between atoms of a molecule, & the lone pairs of electrons that may exist in the molecule. Duet rule Octet rule beyond octet rule ( )

3 3-1-1 Resonance A single Lewis structure often cannot represent the true electronic structure of a molecule. While one can only show an integral number of covalent bonds between two & only two atoms using these diagrams, one often finds that the experimentally deduced or calculated (from Quantum mechanics) structure of the molecule does not match any of the possible Lewis structures but rather has properties in some sense intermediate to these.

4 Bond order =? CO 3 2-, NO 3- and SO 3 are isoelectronic.

5 3-1-2 Expanded shells Beyond the oct rule, an option limited to elements of the third and higher periods is to use d orbitals for this expansion (ClF 3 & SF 6 ).

6 So far we have assumed that third-row & heavier atoms can exceed the octet rule by using their valence d orbitals to accommodate the extra electrons. However, recent calculations indicate that because the 3d orbitals are so much higher in energy than the 3s & 3p orbitals for a given atom it is not feasible to use them. Hyperconjugation might be another explanation.

7 Hyperconjugation covalent & ionic bonds 2/3 covalent bond per S-F interaction 4/5 covalent bond per P-Cl interaction

8 3-1-3 Formal charge formal charge (FC) is a partial charge on an atom in a molecule assigned by assuming that electrons in a chemical bond are shared equally between atoms, regardless of relative electronegativity.

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12 3-1-4 Multiple bonds in Be and B compounds A few molecules, such as BeF 2, BeCl 2, & BF 3, seem to require multiple bonds to satisfy the octet rule for Be & B, even though we do not usually expect multiple bonds for fluorine & chlorine.

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14 B 2 H 6 H B H B H H H H Diborane

15 3-2 Valence shell electron pair repulsion theory (VSEPR) It provides a method for predicting the shape of molecules, based on the electron pair electrostatic repulsion. Steric number (SN = m + n) for AX m E n.

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19 How to determine the positions of lone pairs?

20 How to determine the positions of lone pairs?

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25 3.2.3 Electronegativity and Atomic Size Effects Electronegativity (EN) can act as a guide in the use of formal charge, & also play an important role in determining the arrangement of outer atoms around a central atom & in influencing bond angles.

26 The effects of electronegativity (EN) & atomic size frequently parallel ( 匹敵 ) each other, but in some cases, the sizes of outer atoms & groups may play the more important role.

27 Electronegativity scales : the concept of EN was first introduced by Linus Pauling in the 1930s as a means of describing bond energies. Pauling s definition : Bond energies of polar bonds (HCl : 428 kj/mol) are larger than the average of the bond energies of the two homonuclear species (H 2 : 432 kj/mole; Cl 2 : 240 kj/mol). From the data like these, Pauling calculated EN values that could be used to predict other bond energies (set 4.0 for F- arbitrary).

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29 Configuration energies (CEs), formerly called spectroscopic electronegativities (ENs), are an attempt to quantum mechanically define, & extend, the important chemical concept of ENs. CEs for the elements are defined as the average ionization energies for ground-state free atoms : CE = (n s + m p )/(n + m) n & m : the number of s & p electrons s & p (ev) : the multiplet averaged one-electron energies (NIST energy level tables)

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31 4 EN of the noble gases can be calculated more easily from ionization energies than from bond energies. Because the noble gases have higher ionization energies than the halogens, calculations have suggested that the EN of the noble gases may exceed those of the halogens.

32 The noble gas atoms are somewhat smaller than the neighboring halogen atoms - for example, Ne is smaller than F as a consequence of a greater effective nuclear charge. The charge, which is able to attract noble gas electrons strongly toward the nucleus, is also likely to exert a strong attraction on electrons of neighboring atoms; hence, the high ENs predicted for the noble gases are reasonable.

33 Electronegativity & bond angles : By the VSEPR approach, trends in many bond angles can be explained by EN. Lone pair electrons dominate smallest As the EN of the halogen increases, the halogen exerts a stronger pull on electron pair it shares with the central atom (reduce e-density near the central atom, allowing the lone pair to spread out & reducing the bond angles).

34 The outer atom with larger EN draws the electrons toward itself & away from the central atom, reducing the repulsive effect of these electrons. The compounds of the halogens in Table 3.5 show this effect; the compounds containing F have smaller angles than those containing Cl (and Br, I). As a result, the lone pair effect is relatively larger & forces smaller bond angles. (II)

35 It is worth noting that the above trend is also obtained if size is considered: as the size of the outer atom increases in the order F < Cl < Br <I, the bond angle increases. (II )

36 Similar considerations can be made in situations where the outer atoms remain the same, but the central atom is changed. In these cases, as the central atoms become more electronegative, it pulls electrons in bonding pairs more strongly toward itself, increasing the e-density near the central atom.therefore, the bond angle increases. (I)

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38 Effects of size : the most electronegative atoms have also been the smallest (i.e., F, Cl, Br, & I). Thus, size & EN might have opposite effects (CH 3 v.s. CF 3 ), since a smaller outer group is less electronegative.

39 Molecules having steric number = 5 : in PCl 5, SF 4, and ClF 3, the central atom-axial distances are longer than the distances to equatorial atoms (figure 3.17). This effect has been attributed to the greater repulsion of lone & bonding pairs with atoms in axial positions (three 90 interactions) than with atoms in equatorial positions (two 90 interactions).

40 In addition, there is a tendency for less electronegative groups to occupy equatorial positions, similar to lone pairs & multiply bonded atoms (figure 3.18).

41 3-2-4 Ligand close-packing The ligand close-packing (LCP) model (developed by Gillespie) uses the distances between the outer atoms in molecules as a guide. For a series of molecules with the same central atom, the nonbonded distances between the outer atoms are consistent, but the bond angles & lengths change.

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43 Ligand close-packing and bond distances : The LCP model predicts that non-bonded atom-atom distances in molecules remain approximately the same, even if the bond angles around the central atoms are changed. For example, the F-F distances in NF 4+ and NF 3+ are both 212 pm, even though the F-N-Fbond angles are significantly different (figure 3.22).

44 VSEPR predicts that NF 3 should have the smaller bond angle, and it does : v.s (NF 4+ ). Because the F F distance remain essentially unchanged, the N-F distance in NF 3 must be longer than the 130 pm in NF 4+. (the results of the LCP approach are in many ways consistent with those of the VSEPR model).

45 xsin(102.3 /2)=106 X=136.11pm

46 3-3 Molecular Polarity Dipole moment : = Qr 1 debye = x Cm

47 3-4 Hydrogen bonding

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50 A protein helix

51 The pleated sheet arrangement