1 / 43

Chemical Bonding and Molecular Structure (Chapter 9)

Chemical Bonding and Molecular Structure (Chapter 9). Ionic vs. covalent bonding Molecular orbitals and the covalent bond (Ch. 10) Valence electron Lewis dot structures octet vs. non-octet resonance structures formal charges VSEPR - predicting shapes of molecules

PamelaLan
Télécharger la présentation

Chemical Bonding and Molecular Structure (Chapter 9)

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chemical Bonding and Molecular Structure (Chapter 9) • Ionic vs. covalent bonding • Molecular orbitals and the covalent bond (Ch. 10) • Valence electron Lewis dot structures • octet vs. non-octet • resonance structures • formal charges • VSEPR - predicting shapes of molecules • Bond properties • electronegativity • polarity, bond order, bond strength Bonding and structure (2)

  2. — 2 for # of PAIRS Rules for making Lewis dot structures 1. Count no. of valence electrons (- don’t forget to include the charge on molecular ions!) 2. Place a bond pair (BP) between connected atoms 3. Complete octets by using rest of e- as lone pairs (LP) 4. For atoms with <8 e-, make multiple bonds to complete octets 5. Assign formal charges : fc = Z - (#BP/2) - (#LP) Indicate equivalent (RESONANCE) structures 6. Structures with smaller formal charges are preferred - consider non-octet alternatives (esp. for 3rd, 4th row) • OCTET RULE:#Bond Pairs + #Lone Pairs = 4 • (except for H and atoms of 3rd and higher periods) #lone pairs at central atom in AXn = {(#e-) - 8*n}/2 Bonding and structure (2)

  3. Rules 1-3  O—S —O + + — — Sulfur Dioxide, SO2 These equivalent structures are called: RESONANCE STRUCTURES. The proper Lewis structure is a HYBRID of the two. Each atom has OCTET . . . . . BUT there is a +1 and -1 formal charge Bonding and structure (2)

  4. O = S = O SO2 (2) Alternate Lewis structure for SO2 uses 2 double bonds Sulfur does not obey OCTET rule BUT the formal charge = 0 This is better structure than O=S+-O- since it reduces formal charge (rule 6). 3rd row S atom can have 5 or 6 electron pairs NB: # of central atom lone pairs = (3*6 -8*2)/2 = 1 in both O=S+-O- and O=S=O structures Bonding and structure (2)

  5. A. S=C=N Calculated partial charges B. S=C - N C. S-C N -0.16 -0.32 -0.52 Thiocyanate ion, (SCN)- Which of three possible resonance structures is most important? ANSWER: C > A > B Bonding and structure (2)

  6. 6_VSEPR.mov MOLECULAR GEOMETRY Molecule adopts the shape that minimizes the electron pair repulsions. VSEPR • Valence Shell Electron Pair Repulsion theory. • Most important factor in determining geometry is relative repulsion between electron pairs. Bonding and structure (2)

  7. No. of e- Pairs Around Central Atom linear 2 F—Be—F 180o F planar trigonal 3 B F F 120o 109o H 4 tetrahedral C H H H CAChe image Example Geometry Bonding and structure (2)

  8. lone pair of electrons in tetrahedral position N H H H Structure Determination by VSEPR Ammonia, NH3 There are 4 electron pairs at the corners of a tetrahedron. The ELECTRON PAIR GEOMETRY is tetrahedral. Bonding and structure (2)

  9. lone pair of electrons in tetrahedral position N H H H VSEPR - ammonia Ammonia, NH3 Although the electron pair geometry is tetrahedral . . . . . . the MOLECULAR GEOMETRY — the positions of the atoms — is PYRAMIDAL. Bonding and structure (2)

  10. AXnEm notation • a good way to distinguish between • electron pair and molecular geometries • is theAXnEmnotation • where: • A - atom whose local geometry is of interest (typically the CENTRAL ATOM) • Xn - n atoms bonded to A • Em - m lone pair electrons at A • NH3 is AX3E system  pyramidal • (NB this notation not used by Kotz) Bonding and structure (2)

  11. •• H - O - H •• VSEPR - water Water, H2O 2. Count BP’s and LP’s = 4 3. The 4 electron pairs are at the corners of a tetrahedron. 1. Draw electron dot structure The electron pair geometry is TETRAHEDRAL. Bonding and structure (2)

  12. •• H - O - H •• VSEPR - water (2) Although the electron pair geometry is TETRAHEDRAL . . . . . . the molecular geometry is bent. H2O - AX2E2 system - angular geometry Bonding and structure (2)

  13. O O • • • • C H H • • • • C H H Formaldehyde, CH2O VSEPR - formaldehyde 1. Draw electron dot structure 2. Count BP’s and LP’s: At Carbon there are 4 BP but . . . 3. These are distributed in ONLY 3 regions. Double bond electron pairs are in same region. There are 3 regions of electron density Electron repulsion places them at the corners of a planar triangle. Both the electron pair geometry and the molecular geometry are PLANAR TRIGONAL  120o bond angles. H2CO at the C atom is an AX3 species Bonding and structure (2)

  14. 6_CH3OH.mov VSEPR - Bond Angles Methanol, CH3OH H •• Define bond angles 1 and 2 Angle 1 = H-C-H = ? Angle 2 = H-O-C = ? Answer: H—C—O—H •• H Angle 1 Angle 2 109o because both the C and O atoms are surrounded by 4 electron pairs. AXnEm designation ? at C at O AX4 = tetrahedral AX2E2 = bent Bonding and structure (2)

  15. H N H—C—C •• 1 2 H Acetonitrile, CH3CN VSEPR - bond angles (2) Angle 1 = ? 109o Define bond angles 1 and 2 Angle 2 = ? 180o Why ? : The CH3 carbon is surrounded by 4 bond charges The CN carbon is surrounded by 2 bond charges AXnEm designation ? at CH3 carbon at CN carbon AX4 = tetrahedral AX2 = linear Bonding and structure (2)

  16. What about:STRUCTURES WITH CENTRAL ATOMS THAT DO NOT OBEY THE OCTET RULE ? PF5 BF3 SF4 Bonding and structure (2)

  17. Geometry for non-octet species also obey VSEPR rules Consider boron trifluoride, BF3 The B atom is surrounded by only 3 electron pairs. Bond angles are 120o Molecular Geometry is planar trigonal BF3 is an AX3 species Bonding and structure (2)

  18. 6_VSEPR.mov Trigonal bipyramid 90° F 5 electron pairs F 120° F P F F 90° Octahedron F 6 electron pairs F F F S F 90° F Compounds with 5 or 6 Pairs Around the Central Atom AX5 system AX6 system Bonding and structure (2)

  19. •• •• • • F • •• •• •• • • • • S F F • • •• •• • • • • F • • •• F • F • S • • F F F F S F F Sulfur Tetrafluoride, SF4 Number of valence e- = 34 No. of S lone pairs = {17 - 4 b.p. - 3x4 l.p.(F)} = 1 lone pair on S There are 5 (BP + LP) e- pairs around the S THEREFORE: electron pair geometry ? = trigonal bipyramid OR AX4E system. Molecular geometry ? Bonding and structure (2)

  20. F F • S • F equatorial F axial Sulfur Tetrafluoride, SF4 (2) 90° 120° Lone pair is in the equatorial position because it requires more room than a bond pair. Molecular geometry of SF4 is “see-saw” Q: What is molecular geometry of SO2 ? Bonding and structure (2)

  21. 6_CH4.mov Bonding with Hybrid Atomic Orbitals - Carbon prefers to make 4 bonds as in CH4 4 C atom orbitals hybridize to form four equivalent sp3 hybrid atomic orbitals. But atomic carbon has an s2p2 configuration Why can it make more than 2 bonds ? Bonding and structure (2)

  22. Orbital Hybridization BONDS SHAPE HYBRID REMAIN e.g. s2p2  2 linear {2 x sp & 2 p’s} C2H2 3 trigonal {3 x sp2 & 1 p} C2H4 planar 4 tetrahedral {4 xsp3 } CH4 Bonding and structure (2)

  23. ­ ­ ­ ­ ­¯ ­ ­ p 2s 2p 3 sp2 orbital hybrid orbitals H H sp2 120° C C H H Multiple Bondss and p Bonding in C2H4 • The extra p orbital electron on each C atom overlaps the p orbital on the neighboring atom to form the p bond. C atom orbitals are COMBINED (= re-hybridized) to form orbitals better suited for BONDING • The 3 sp2 hybrid orbitals • are used to make the C-C • and two C-H  bonds 6_C2H4-sg.mov 6_C2H4-pi.mov 6_C2H4.mov Bonding and structure (2)

  24. 233 E (kJ/mol) 27 -180 0 180 C-C=C angle (o) Consequences of Multiple Bonding Restricted rotation around C=C bond in 1-butene = CH2=CH-CH2-CH3. See Butene.Map in ENER_MAP in CAChe models. P. 475 - Photo-rotation about double bonds lets us see !! Bonding and structure (2)

  25. Bond Properties - bond order - bond length - bond strength - bond polarity - MOLECULAR polarity • What is the effect of bonding and structure on molecular properties ? Buckyball in HIV-protease, see page 107 Bonding and structure (2)

  26. triple, BO = 3 H H and 2 p 1 s H C C C N double, BO = 2 single and 1 p 1 s BO = 1 1 s Bond Order • the number of bonds between a pair of atoms. CH2CHCN Acrylonitrile Bonding and structure (2)

  27. Total # of e - pairs used for a type of bond Bond order = Total # of bonds of that type 3 (e - pairs in N-O bonds) = Bond order in NO2- 2 (N - O bonds) Bond Order (2) Fractional bond orders occur in molecules with resonance structures. Consider NO2- N-O bond order in NO2- = 1.5 Bonding and structure (2)

  28. 110 pm 745 kJ 123 pm 414 kJ Formaldehye Bond Order and Bond Length Bond order is related to two important bond properties: (a) bond strength as given by DE (b) Bond length - the distance between the nuclei of two bonded atoms. Bonding and structure (2)

  29. Bond Length Molecule R(H-X) H- F 104 pm H- Cl 131 pm H- I 165 pm - depends on size of bonded atoms: - depends on bond order. Molecule R(C-O) CH3C- OH 141 pm O=C=O 132 pm C O 119 pm Bonding and structure (2)

  30. Bond Strength • Bond Dissociation energy (DE) - energy required to break a bond in gas phase. • See Table 9.5 BOND STRENGTH (kJ/mol) LENGTH (pm) H—H 436 74 C—C 347 154 C=C 611 134 CºC 837 121 NºN 946 110 The GREATER the number of bonds (bond order) the HIGHER the bond strength and the SHORTER the bond. Bonding and structure (2)

  31. Bond Strength (2) Bond Order Length Strength HO—OH 1 149 pm 210 kJ/mol O=O 2 121 498 kJ/mol 1.5 128 ? 303 kJ/mol O3 (g)  3 O(g) HOW TO CALCULATE ? Hrxn = {3xHf(O) - Hf(O3)} = {3x249.2 - 142.7} = 605 kJ/mol 2 O-O bonds in O3  DE (O3) = 605/2 = 302.5 kJ/mol Bonding and structure (2)

  32. Bond Polarity HCl is POLAR because it has a positive end and a negative end (partly ionic). Polarity arises because Cl has a greater share of the bonding electrons than H. Calculated charge by CAChe: H (red) is +ve (+0.20 e-) Cl (yellow) is -ve (-0.20 e-). (See PARTCHRG folder in MODELS.) Bonding and structure (2)

  33. Bond Polarity (2) • Due to the bond polarity, the H—Cl bond energy is GREATER than expected for a “pure” covalent bond. BOND ENERGY “pure” bond 339 kJ/mol calculated real bond 432 kJ/mol measured Difference 92 kJ/mol. This difference is the contribution of IONIC bonding It is proportional to the difference in ELECTRONEGATIVITY, c. Bonding and structure (2)

  34. Electronegativity, c c is a measure of the ability of an atom in a molecule to attract electrons to itself. Concept proposed by Linus Pauling (1901-94) Nobel prizes: Chemistry (54), Peace (63) See p. 425; 008vd3.mov (CD) Bonding and structure (2)

  35. Electronegativity, c Figure 9.7 • F has maximum c. • Atom with lowest c is the center atom in most molecules. • Relative values of c determines BOND POLARITY (and point of attack on a molecule). Bonding and structure (2)

  36. Bond Polarity Which bond is more polar ? (has larger bond DIPOLE) O—H O—F c H 2.1 O F 3.5 4.0 c(A) - c(B)3.5 - 2.1 Dc 1.4 3.5 - 4.0 0.5 (O-H) > (O-F) Therefore OH is more polar than OF Also note that polarity is “reversed.” Bonding and structure (2)

  37. Molecular Polarity • Molecules—such as HCl and H2O— can be POLAR (or dipolar). • They have a DIPOLE MOMENT. • Polar molecules turn to align their dipole with an electric field. Bonding and structure (2)

  38. Symmetric molecules Predicting molecular polarity A molecule will be polar ONLY if a) it contains polar bonds AND b) the molecule is NOT “symmetric” Bonding and structure (2)

  39. Molecular Polarity: H2O Water is polar because: a) O-H bond is polar b) water is non-symmetric The dipole associated with polar H2O is the basis for absorption of microwaves used in cooking with a microwave oven Bonding and structure (2)

  40. -0.73 +1.46 -0.73 Carbon Dioxide • CO2 is NOT polar even though the CO bonds are polar. • Because CO2 is symmetrical the BOND polarity cancels The positive C atom is why water attaches to CO2 CO2 + H2O  H2CO3 Bonding and structure (2)

  41. HBF2 is polar BF3 is NOT polar Molecular Polarity in NON-symmetric molecules B—F, B—H bonds polar molecule is NOT symmetric Atom Chg.  B +ve 2.0 H +ve 2.1 F -ve 4.0 B +ve F -ve B—F bonds are polar molecule is symmetric Bonding and structure (2)

  42. Fluorine-substituted Ethylene: C2H2F2 C—F bonds are MUCH more polar than C—H bonds. (C-F) = 1.5, (C-H) = 0.4 CIS isomer • both C—F bonds on same side  molecule is POLAR. TRANS isomer • both C—F bonds on opposite side  molecule is NOT POLAR. Bonding and structure (2)

  43. Chemical Bonding and Molecular Structure (Chapter 9) • Ionic vs. covalent bonding • Molecular orbitals and the covalent bond (Ch. 10) • Valence electron Lewis dot structures • octet vs. non-octet • resonance structures • formal charges • VSEPR - predicting shapes of molecules • Bond properties • electronegativity • polarity, bond order, bond strength Bonding and structure (2)

More Related