Presentation is loading. Please wait.

Presentation is loading. Please wait.

1 © 2006 Brooks/Cole - Thomson Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver CHAPTER 9 Bonding and Molecular.

Similar presentations


Presentation on theme: "1 © 2006 Brooks/Cole - Thomson Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver CHAPTER 9 Bonding and Molecular."— Presentation transcript:

1 1 © 2006 Brooks/Cole - Thomson Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver CHAPTER 9 Bonding and Molecular Structure: Fundamental Concepts © 2006 Brooks/Cole Thomson Lectures written by John Kotz

2 2 © 2006 Brooks/Cole - Thomson Chemical Bonding Problems and questions — How is a molecule or polyatomic ion held together? Why are atoms distributed at strange angles? Why are molecules not flat? Can we predict the structure? How is structure related to chemical and physical properties?

3 3 © 2006 Brooks/Cole - Thomson Structure & Bonding NN triple bond. Molecule is unreactive Phosphorus is a tetrahedron of P atoms. Very reactive! Red phosphorus, a polymer. Used in matches.

4 4 © 2006 Brooks/Cole - Thomson Forms of Chemical Bonds There are 2 extreme forms of connecting or bonding atoms:There are 2 extreme forms of connecting or bonding atoms: Ionic —complete transfer of 1 or more electrons from one atom to anotherIonic —complete transfer of 1 or more electrons from one atom to another Covalent —some valence electrons shared between atomsCovalent —some valence electrons shared between atoms Most bonds are somewhere in between.Most bonds are somewhere in between.

5 5 © 2006 Brooks/Cole - Thomson Ionic Compounds Metal Nonmetal 2 Na(s) + Cl 2 (g) ---> 2 Na Cl -

6 6 © 2006 Brooks/Cole - Thomson Covalent Bonding The bond arises from the mutual attraction of 2 nuclei for the same electrons. Electron sharing results. Bond is a balance of attractive and repulsive forces.

7 7 © 2006 Brooks/Cole - Thomson Bond Formation A bond can result from a “head-to-head” overlap of atomic orbitals on neighboring atoms. Cl HH + Overlap of H (1s) and Cl (3p) Note that each atom has a single, unpaired electron.

8 8 © 2006 Brooks/Cole - Thomson Chemical Bonding: Objectives Objectives are to understand: 1. valence e- distribution in molecules and ions. 2. molecular structures 3. bond properties and their effect on molecular properties.

9 9 © 2006 Brooks/Cole - Thomson Electron Distribution in Molecules Electron distribution is depicted with Lewis electron dot structuresElectron distribution is depicted with Lewis electron dot structures Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. G. N. Lewis

10 10 © 2006 Brooks/Cole - Thomson Bond and Lone Pairs Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS.Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. HCl lone pair (LP) shared or bond pair This is called a LEWIS ELECTRON DOT structure.

11 11 © 2006 Brooks/Cole - Thomson Valence Electrons Electrons are divided between core and valence electrons B 1s 2 2s 2 2p 1 Core = [He], valence = 2s 2 2p 1 Br [Ar] 3d 10 4s 2 4p 5 Core = [Ar] 3d 10, valence = 4s 2 4p 5

12 12 © 2006 Brooks/Cole - Thomson Rules of the Game No. of valence electrons of a main group atom = Group number For Groups 1A-4A, no. of bond pairs = group number. For Groups 5A -7A, BP’s = 8 - Grp. No. For Groups 5A -7A, BP’s = 8 - Grp. No.

13 13 © 2006 Brooks/Cole - Thomson Rules of the Game Except for H (and sometimes atoms of 2 nd and 3 rd families and 3rd and higher periods), BP’s + LP’s = 4 This observation is called the OCTET RULE

14 14 © 2006 Brooks/Cole - Thomson Building a Dot Structure Ammonia, NH 3 1. Decide on the central atom; never H. Central atom is atom of lowest affinity for electrons. Central atom is atom of lowest affinity for electrons. Therefore, N is central Therefore, N is central 2. Count valence electrons H = 1 and N = 5 H = 1 and N = 5 Total = (3 x 1) + 5 Total = (3 x 1) + 5 = 8 electrons / 4 pairs = 8 electrons / 4 pairs

15 15 © 2006 Brooks/Cole - Thomson 3.Form a single bond between the central atom and each surrounding atom H H H N Building a Dot Structure H H H N 4.Remaining electrons form LONE PAIRS to complete octet as needed. 3 BOND PAIRS and 1 LONE PAIR. Note that N has a share in 4 pairs (8 electrons), while H shares 1 pair.

16 16 © 2006 Brooks/Cole - Thomson 10 pairs of electrons are now left. 10 pairs of electrons are now left. Sulfite ion, SO 3 2- Step 1. Central atom = S Step 2. Count valence electrons S = 6 3 x O = 3 x 6 = 18 3 x O = 3 x 6 = 18 Negative charge = 2 Negative charge = 2 TOTAL = 26 e- or 13 pairs Step 3. Form bonds

17 17 © 2006 Brooks/Cole - Thomson Sulfite ion, SO 3 2- Remaining pairs become lone pairs, first on outside atoms and then on central atom. OO O S Each atom is surrounded by an octet of electrons.

18 18 © 2006 Brooks/Cole - Thomson Carbon Dioxide, CO 2 1. Central atom = _______ 2. Valence electrons = __ or __ pairs 3. Form bonds. 4. Place lone pairs on outer atoms. This leaves 6 pairs.

19 19 © 2006 Brooks/Cole - Thomson Carbon Dioxide, CO 2 4. Place lone pairs on outer atoms. The second bonding pair forms a pi (π) bond. 5. So that C has an octet, we shall form DOUBLE BONDS between C and O.

20 20 © 2006 Brooks/Cole - Thomson Double and even triple bonds are commonly observed for C, N, P, O, and S H 2 CO SO 3 C2F4C2F4C2F4C2F4

21 21 © 2006 Brooks/Cole - Thomson Sulfur Dioxide, SO 2 1. Central atom = S 2. Valence electrons = 18 or 9 pairs bring in left pair OR bring in right pair 3. Form double bond so that S has an octet — but note that there are two ways of doing this. OS O

22 22 © 2006 Brooks/Cole - Thomson Sulfur Dioxide, SO 2 This leads to the following structures. These equivalent structures are called RESONANCE STRUCTURES. The true electronic structure is a HYBRID of the two. RESONANCE STRUCTURES. The true electronic structure is a HYBRID of the two.

23 23 © 2006 Brooks/Cole - Thomson Formal Atom Charges Atoms in molecules often bear a charge (+ or -). The predominant resonance structure of a molecule is the one with charges as close to 0 as possible. Formal charge = Group number – 1/2 (no. of bonding electrons) - (no. of LP electrons)Formal charge = Group number – 1/2 (no. of bonding electrons) - (no. of LP electrons)

24 24 © 2006 Brooks/Cole - Thomson OOC +4 - (1/2)(8) - 0 = (1/2)(4) - 4 = 0 Carbon Dioxide, CO 2

25 25 © 2006 Brooks/Cole - Thomson Thiocyanate Ion, SCN (1/2)(2) - 6 = (1/2)(6) - 2 = (1/2)(8) - 0 = 0 SNC

26 26 © 2006 Brooks/Cole - Thomson Thiocyanate Ion, SCN - SNC SNC SNC Which is the most important resonance form?

27 27 © 2006 Brooks/Cole - Thomson Calculated Partial Charges in SCN - All atoms negative, but most on the S SNC

28 28 © 2006 Brooks/Cole - Thomson Violations of the Octet Rule Usually occurs with B and elements of higher periods. BF 3 SF 4

29 29 © 2006 Brooks/Cole - Thomson Boron Trifluoride Central atom = _____________Central atom = _____________ Valence electrons = __________ or electron pairs = __________Valence electrons = __________ or electron pairs = __________ Assemble dot structureAssemble dot structure The B atom has a share in only 6 pairs of electrons (or 3 pairs). B atom in many molecules is electron deficient.

30 30 © 2006 Brooks/Cole - Thomson Boron Trifluoride, BF 3 What if we form a B—F double bond to satisfy the B atom octet? F F F B ++ --

31 31 © 2006 Brooks/Cole - Thomson Sulfur Tetrafluoride, SF 4 Central atom =Central atom = Valence electrons = ___ or ___ pairs.Valence electrons = ___ or ___ pairs. Form sigma bonds and distribute electron pairs.Form sigma bonds and distribute electron pairs. 5 pairs around the S atom. A common occurrence outside the 2nd period.

32 MOLECULAR GEOMETRY

33 33 © 2006 Brooks/Cole - Thomson VSEPR VSEPR V alence S hell E lectron P air R epulsion theory.V alence S hell E lectron P air R epulsion theory. Most important factor in determining geometry is relative repulsion between electron pairs.Most important factor in determining geometry is relative repulsion between electron pairs. Molecule adopts the shape that minimizes the electron pair repulsions. MOLECULAR GEOMETRY

34 34 © 2006 Brooks/Cole - Thomson Electron Pair Geometries Active Figure 9.8

35 35 © 2006 Brooks/Cole - Thomson

36 36 © 2006 Brooks/Cole - Thomson

37 37 © 2006 Brooks/Cole - Thomson

38 38 © 2006 Brooks/Cole - Thomson Electron Pair Geometries Active Figure 9.8

39 39 © 2006 Brooks/Cole - Thomson Structure Determination by VSEPR Ammonia, NH 3 1. Draw electron dot structure 2. Count BP’s and LP’s = 4 H H H N 3. The 4 electron pairs are at the corners of a tetrahedron.

40 40 © 2006 Brooks/Cole - Thomson Structure Determination by VSEPR Ammonia, NH 3 There are 4 electron pairs at the corners of a tetrahedron. The ELECTRON PAIR GEOMETRY is tetrahedral.

41 41 © 2006 Brooks/Cole - Thomson Ammonia, NH 3 The electron pair geometry is tetrahedral. Structure Determination by VSEPR The MOLECULAR GEOMETRY — the positions of the atoms — is PYRAMIDAL.

42 42 © 2006 Brooks/Cole - Thomson Structure Determination by VSEPR Water, H 2 O 1. Draw electron dot structure The electron pair geometry is TETRAHEDRAL. 2. Count BP’s and LP’s = 4 3. The 4 electron pairs are at the corners of a tetrahedron.

43 43 © 2006 Brooks/Cole - Thomson Structure Determination by VSEPR Water, H 2 O The electron pair geometry is TETRAHEDRAL The molecular geometry is BENT.

44 44 © 2006 Brooks/Cole - Thomson Geometries for Four Electron Pairs Figure 9.9

45 45 © 2006 Brooks/Cole - Thomson Structure Determination by VSEPR Formaldehyde, CH 2 O 1. Draw electron dot structure The electron pair geometry is PLANAR TRIGONAL with 120 o bond angles. CHH O C HH O 2. Count BP’s and LP’s at C 3. There are 3 electron “lumps” around C at the corners of a planar triangle.

46 46 © 2006 Brooks/Cole - Thomson Structure Determination by VSEPR Formaldehyde, CH 2 O The electron pair geometry is PLANAR TRIGONAL The molecular geometry is also planar trigonal.

47 47 © 2006 Brooks/Cole - Thomson H-C-H = 109 o C-O-H = 109 o In both cases the atom is surrounded by 4 electron pairs. Structure Determination by VSEPR H H H—C—O—H 109˚ Methanol, CH 3 OH Define H-C-H and C-O-H bond angles

48 48 © 2006 Brooks/Cole - Thomson Structure Determination by VSEPR Acetonitrile, CH 3 CN H H H—C—CN 180˚ 109˚ H-C-H = 109 o C-C-N = 180 o One C is surrounded by 4 electron “lumps” and the other by 2 “lumps” Define unique bond angles

49 49 © 2006 Brooks/Cole - Thomson Phenylalanine, an amino acid

50 50 © 2006 Brooks/Cole - Thomson Phenylalanine

51 51 © 2006 Brooks/Cole - Thomson Structures with Central Atoms with More Than or Less Than 4 Electron Pairs Often occurs with Group 3A elements and with those of 3rd period and higher.

52 52 © 2006 Brooks/Cole - Thomson Boron Compounds Consider boron trifluoride, BF 3 Geometry described as planar trigonal The B atom is surrounded by only 3 electron pairs. Bond angles are 120 o

53 53 © 2006 Brooks/Cole - Thomson 5 electron pairs Compounds with 5 or 6 Pairs Around the Central Atom

54 54 © 2006 Brooks/Cole - Thomson Molecular Geometries for Five Electron Pairs Figure 9.11 All based on trigonal bipyramid

55 55 © 2006 Brooks/Cole - Thomson Number of valence electrons = 34Number of valence electrons = 34 Central atom = S Central atom = S Dot structureDot structure Sulfur Tetrafluoride, SF 4 Electron pair geometry --> trigonal bipyramid (because there are 5 pairs around the S) Electron pair geometry --> trigonal bipyramid (because there are 5 pairs around the S)

56 56 © 2006 Brooks/Cole - Thomson Lone pair is in the equator because it requires more room. Sulfur Tetrafluoride, SF 4

57 57 © 2006 Brooks/Cole - Thomson Molecular Geometries for Six Electron Pairs Figure 9.14 All are based on the 8- sided octahedron

58 58 © 2006 Brooks/Cole - Thomson 6 electron pairs Compounds with 5 or 6 Pairs Around the Central Atom

59 59 © 2006 Brooks/Cole - Thomson Bond Properties What is the effect of bonding and structure on molecular properties? Free rotation around C–C single bond No rotation around C=C double bond

60 60 © 2006 Brooks/Cole - Thomson Bond Order # of bonds between a pair of atoms # of bonds between a pair of atoms Bond Order # of bonds between a pair of atoms # of bonds between a pair of atoms Double bond Single bond Triple bond AcrylonitrileAcrylonitrile

61 61 © 2006 Brooks/Cole - Thomson Bond Order Fractional bond orders occur in molecules with resonance structures. Consider NO 2 - The N—O bond order = 1.5

62 62 © 2006 Brooks/Cole - Thomson Bond Order Bond order is proportional to two important bond properties: (a) bond strength (b)bond length 745 kJ 414 kJ 110 pm 123 pm

63 63 © 2006 Brooks/Cole - Thomson Bond Length Bond length depends on size of bonded atoms. H—F H—Cl H—I Bond distances measured in Angstrom units where 1 A = pm.

64 64 © 2006 Brooks/Cole - Thomson Bond length depends on bond order. Bond distances measured in Angstrom units where 1 A = pm. Bond Length

65 65 © 2006 Brooks/Cole - Thomson —measured by the energy req’d to break a bond. See Table BOND STRENGTH (kJ/mol) H—H436 C—C346 C=C602 C  C 835 N  N945 The GREATER the number of bonds (bond order) the HIGHER the bond strength and the SHORTER the bond. Bond Strength

66 66 © 2006 Brooks/Cole - Thomson BondOrderLengthStrength HO—OH O=O pm 210 kJ/mol ?? Bond Strength

67 67 © 2006 Brooks/Cole - Thomson Molecular Polarity Why do ionic compounds dissolve in water? Water Boiling point = 100 ˚C Methane Boiling point = -161 ˚C Why do water and methane differ so much in their boiling points?

68 68 © 2006 Brooks/Cole - Thomson Bond Polarity HCl is POLAR because it has a positive end and a negative end. Cl has a greater share in bonding electrons than does H. Cl has slight negative charge (-  ) and H has slight positive charge (+  )

69 69 © 2006 Brooks/Cole - Thomson Bond Polarity Three molecules with polar, covalent bonds.Three molecules with polar, covalent bonds. Each bond has one atom with a slight negative charge (-  ) and and another with a slight positive charge (+  )Each bond has one atom with a slight negative charge (-  ) and and another with a slight positive charge (+  )

70 70 © 2006 Brooks/Cole - Thomson This model, calc’d using CAChe software for molecular calculations, shows that H is + (red) and Cl is - (yellow). Calc’d charge is + or Bond Polarity

71 71 © 2006 Brooks/Cole - Thomson Due to the bond polarity, the H—Cl bond energy is GREATER than expected for a “pure” covalent bond. BONDENERGY “pure” bond339 kJ/mol calc’d real bond432 kJ/mol measured BONDENERGY “pure” bond339 kJ/mol calc’d real bond432 kJ/mol measured Difference = 92 kJ. This difference is proportional to the difference in ELECTRONEGATIVITY, . Bond Polarity

72 72 © 2006 Brooks/Cole - Thomson Electronegativity,   is a measure of the ability of an atom in a molecule to attract electrons to itself. Concept proposed by Linus Pauling

73 73 © 2006 Brooks/Cole - Thomson Linus Pauling, The only person to receive two unshared Nobel prizes (for Peace and Chemistry). Chemistry areas: bonding, electronegativity, protein structure

74 74 © 2006 Brooks/Cole - Thomson Electronegativity Figure 9.14

75 75 © 2006 Brooks/Cole - Thomson F has maximum .F has maximum . Atom with lowest  is the center atom in most molecules.Atom with lowest  is the center atom in most molecules. Relative values of  determine BOND POLARITY (and point of attack on a molecule).Relative values of  determine BOND POLARITY (and point of attack on a molecule). Electronegativity,  See Figure 9.14

76 76 © 2006 Brooks/Cole - Thomson Bond Polarity Which bond is more polar (or DIPOLAR)? O—HO—F O—HO—F   OH is more polar than OF OH is more polar than OF and polarity is “reversed.”

77 77 © 2006 Brooks/Cole - Thomson Molecular Polarity Molecules—such as HI and H 2 O— can be POLAR (or dipolar). They have a DIPOLE MOMENT. The polar HCl molecule will turn to align with an electric field.

78 78 © 2006 Brooks/Cole - Thomson Molecular Polarity The magnitude of the dipole is given in Debye units. Named for Peter Debye ( ). Rec’d 1936 Nobel prize for work on x- ray diffraction and dipole moments.

79 79 © 2006 Brooks/Cole - Thomson Dipole Moments Why are some molecules polar but others are not?

80 80 © 2006 Brooks/Cole - Thomson Molecular Polarity Molecules will be polar if a)bonds are polar AND AND b)the molecule is NOT “symmetric” All above are NOT polar

81 81 © 2006 Brooks/Cole - Thomson Polar or Nonpolar? Compare CO 2 and H 2 O. Which one is polar?

82 82 © 2006 Brooks/Cole - Thomson Carbon Dioxide CO 2 is NOT polar even though the CO bonds are polar.CO 2 is NOT polar even though the CO bonds are polar. CO 2 is symmetrical.CO 2 is symmetrical Positive C atom is reason CO 2 and H 2 O react to give H 2 CO 3

83 83 © 2006 Brooks/Cole - Thomson Polar or Nonpolar? Consider AB 3 molecules: BF 3, Cl 2 CO, and NH 3.

84 84 © 2006 Brooks/Cole - Thomson Molecular Polarity, BF 3 B atom is positive and F atoms are negative. B—F bonds in BF 3 are polar. But molecule is symmetrical and NOT polar

85 85 © 2006 Brooks/Cole - Thomson Molecular Polarity, HBF 2 B atom is positive but H & F atoms are negative. B—F and B—H bonds in HBF 2 are polar. But molecule is NOT symmetrical and is polar.

86 86 © 2006 Brooks/Cole - Thomson Is Methane, CH 4, Polar? Methane is symmetrical and is NOT polar.

87 87 © 2006 Brooks/Cole - Thomson Is CH 3 F Polar? C—F bond is very polar. Molecule is not symmetrical and so is polar.

88 88 © 2006 Brooks/Cole - Thomson CH 4 … CCl 4 Polar or Not? Only CH 4 and CCl 4 are NOT polar. These are the only two molecules that are “symmetrical.”

89 89 © 2006 Brooks/Cole - Thomson Substituted Ethylene C—F bonds are MUCH more polar than C—H bonds.C—F bonds are MUCH more polar than C—H bonds. Because both C—F bonds are on same side of molecule, molecule is POLAR.Because both C—F bonds are on same side of molecule, molecule is POLAR.

90 90 © 2006 Brooks/Cole - Thomson Substituted Ethylene C—F bonds are MUCH more polar than C—H bonds.C—F bonds are MUCH more polar than C—H bonds. Because both C—F bonds are on opposing ends of molecule, molecule is NOT POLAR.Because both C—F bonds are on opposing ends of molecule, molecule is NOT POLAR.


Download ppt "1 © 2006 Brooks/Cole - Thomson Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver CHAPTER 9 Bonding and Molecular."

Similar presentations


Ads by Google