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Molecular Geometry and Bonding Theories AP Chemistry – Ch 9 Mr. Christopherson.

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2 Molecular Geometry and Bonding Theories AP Chemistry – Ch 9 Mr. Christopherson

3 Molecular Geometry and Bonding Theories AP Chemistry – Ch 9 Mr. Christopherson

4 Bonding Theories & Geometry Molecular Geometry (shapes) VSEPR Theory Lewis Structures Molecular Polarity (dipoles) Covalent Bonds Hybridization Ionic Bonds

5 Episode 9 Episode 9 – Molecular Architecture Episode 8 Episode 8 – Chemical Bonds VIDEO ON DEMAND Elements bond to form compounds by giving, taking, or sharing electrons. The differences between ionic and covalent bonds are explained by the use of scientific models and examples from nature. World of Chemistry The Annenberg Film Series VIDEO ON DEMAND The shape and physical properties of a molecule are determined by the electronic structure of its elements and their bonds. How living organisms distinguish between similar molecules (isomers) is revealed.

6 molecular formula structural formula molecular shape ball-and-stick model CH 4 C H H HH H H H H o C tetrahedron tetrahedral shape of methane C H H H H

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8 109.5 o

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10 Tetrahedron

11 Central Atom Central Atom

12 Central Atom Central Atom

13 Substituents

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18 Methane, CH 4

19 Tetrahedral geometry Tetrahedral geometry Methane, CH 4 Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

20 Methane & Carbon Tetrachloride molecular formula structural formula molecular shape ball-and-stick model CH 4 C H H HH H H H H o C CCl 4 space-filling model C Cl

21 Molecular Geometry H H H H o C Linear Trigonal planar Tetrahedral Trigonal pyramidal Bent o o o H 2 O CH 4 AsCl 3 AsF 5 BeH 2 BF 3 CO o

22 A Lone PairA Lone Pear

23 C o H H H H N 107 o H H H.. O o H H.. CH 4, methaneNH 3, ammoniaH 2 O, water.. O O O lone pair electrons O OO O 3, ozone

24 Molecular Shapes Three atoms (AB 2 )Four atoms (AB 3 ) Five atoms (AB 4 ) Six atoms (AB 5 ) Seven atoms (AB 6 ) Linear (180 o ) Bent Trigonal planar (120 o ) Trigonal pyramidal T-shaped Tetrahedral ( o ) Square planar Seesaw Trigonal bipyramidal (B e AB e, 120 o ) & (B e AB a, 90 o ) Square pyramidal Octahedral BBA B B A B lineartrigonal planar B A B B B tetrahedral A BeBe BeBe BeBe BaBa BaBa Trigonal bipyramidal B B B B B B A Bailar, Moeller, Kleinberg, Guss, Castellion, Metz, Chemistry, 1984, page 313.

25 Bonding and Shape of Molecules Number of Bonds Number of Unshared Pairs ShapeExamples Linear Trigonal planar Tetrahedral Pyramidal Bent BeCl 2 BF 3 CH 4, SiCl 4 NH 3, PCl 3 H 2 O, H 2 S, SCl 2 -Be- B C N : O : : Covalent Structure

26 Molecular Shapes AB 2 Linear AB 3 Trigonal planar AB 4 Tetrahedral AB 5 Trigonal bipyramidal AB 6 Octahedral AB 2 E Angular or Bent AB 3 E Trigonal pyramidal AB 2 E 2 Angular or Bent AB 4 E Irregular tetrahedral (see saw) AB 3 E 2 T-shaped AB 2 E 3 Linear AB 6 E Square pyramidal AB 5 E 2 Square planar

27 V alence S hell E lectron P air R epulsion Theory Planar triangular Tetrahedral Trigonal bipyramidal Octahedral

28 V alence S hell E lectron P air R epulsion Theory Planar triangular Tetrahedral Trigonal bipyramidal Octahedral

29 .. The VSEPR Model OO C Linear The Shapes of Some Simple AB n Molecules OO S Bent OO S O Trigonal planar F F F N Trigonal pyramidal T-shapedSquare planar FF Cl F F F Xe FF F F F P F F Trigonal bipyramidal Octahedral F F F S F F F SF 6 SO 2 Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 305

30 Molecular Shapes AB 2 Linear AB 3 Trigonal planar AB 4 Tetrahedral AB 5 Trigonal bipyramidal AB 6 Octahedral AB 2 E Angular or Bent AB 3 E Trigonal pyramidal AB 2 E 2 Angular or Bent AB 4 E Irregular tetrahedral (see saw) AB 3 E 2 T-shaped AB 2 E 3 Linear AB 5 E Square pyramidal AB 4 E 2 Square planar

31 Geometry of Covalent Molecules AB n, and AB n E m AB 2 AB 2 E AB 2 E 2 AB 2 E 3 AB 3 AB 3 E AB 3 E 2 AB 4 AB 4 E AB 4 E 2 AB 5 AB 5 E AB Linear Trigonal planar Tetrahedral Trigonal bipyramidal Trigonal planar Tetrahedral Triangular bipyramidal Tetrahedral Triangular bipyramidal Octahedral Triangular bipyramidal Octahedral Linear Angular, or bent Linear Trigonal planar Triangular pyramidal T-shaped Tetrahedral Irregular tetrahedral (or “see-saw”) Square planar Triangular bipyramidal Square pyramidal Octahedral CdBr 2 SnCl 2, PbI 2 OH 2, OF 2, SCl 2, TeI 2 XeF 2 BCl 3, BF 3, GaI 3 NH 3, NF 3, PCl 3, AsBr 3 ClF 3, BrF 3 CH 4, SiCl 4, SnBr 4, ZrI 4 SF 4, SeCl 4, TeBr 4 XeF 4 PF 5, PCl 5 (g), SbF 5 ClF 3, BrF 3, IF 5 SF 6, SeF 6, Te(OH) 6, MoF 6 Type Formula Shared Electron Pairs Unshared Electron Pairs Ideal Geometry Observed Molecular ShapeExamples Bailar, Moeller, Kleinberg, Guss, Castellion, Metz, Chemistry, 1984, page 317.

32 Predicting the Geometry of Molecules Lewis electron-pair approach predicts number and types of bonds between the atoms in a substance and indicates which atoms have lone pairs of electrons but gives no information about the actual arrangement of atoms in space Valence-shell electron-pair repulsion (VSEPR) model predicts the shapes of many molecules and polyatomic ions but provides no information about bond lengths or the presence of multiple bonds

33 Introduction to Lewis Structures Lewis dot symbols 1. Used for predicting the number of bonds formed by most elements in their compounds 2. Consists of the chemical symbol for an element surrounded by dots that represent its valence electrons 3. A single electron is represented as a single dot

34 Lewis Structures 1) Count up total number of valence electrons 2) Connect all atoms with single bonds - “multiple” atoms usually on outside - “single” atoms usually in center; C always in center, H always on outside. 3) Complete octets on exterior atoms (not H, though) 4) Check - valence electrons math with Step 1 - all atoms (except H) have an octet;if not, try multiple bonds - any extra electrons?Put on central atom

35 Molecules with Expanded Valence Shells Atoms that have expanded octets have AB 5 (trigonal bipyramidal) or AB 6 (octahedral) electron domain geometries. Trigonal bipyramidal structures have a plane containing three electron pairs. The fourth and fifth electron pairs are located above and below this plane. In this structure two trigonal pyramids share a base. For octahedral structures, there is a plane containing four electron pairs. Similarly, the fifth and sixth electron pairs are located above and below this plane. Two square pyramids share a base. F F F P F F F F F S F F F

36 Trigonal Bipyramid F F F P F F The three electron pairs in the plane are called equatorial. The two electron pairs above and below this plane are called axial. The axial electron pairs are 180 o apart and 90 o from to the equatorial electrons. The equatorial electron pairs are 120 o apart. To minimize electron-electron repulsions, nonbonding pairs are always placed in equatorial positions, and bonding pairs in either axial or equatorial positions.

37 Octahedron The four electron pairs in the plane are 90 o to each other. The remaining two electron pairs are 180 o apart and 90 o from the electrons in the plane. Because of the symmetry of the system, each position is equivalent. The equatorial electron pairs are 120 o apart. If we have five bonding pairs and one nonbonding pair, it doesn’t matter where the nonbonding pair is placed.  The molecular geometry is square pyramidal. If two nonbonding pairs are present, the repulsions are minimized by pointing them toward opposite sides of the octahedron.  The molecular geometry is square planar. F F F S F F F F F Xe FF

38 Electron-Domain Geometries Number of Electron Domains Arrangement of Electron Domains Electron-Domain Geometry Predicted Bond Angles Linear Trigonal planar Tetrahedral Trigonal- bipyramidal Octahedral 180 o 120 o o 120 o 90 o A BeBe BeBe BeBe BaBa BaBa B B B B B B A B B A B B A B B B BBA

39 Number of electron domains Electron-domain geometry Predicted bond angles Tetrahedral Trigonal planar Tetrahedral o 120 o o CCOHH H HO 434 Acetic Acid, CH 3 COOH Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 314 Hybridization of central atomsp 3 sp 2 none

40 Intermolecular Forces Ion-ion (ionic bonds) Ion-dipole Dipole-dipole Hydrogen bonding London dispersion forces  +  − − + −+

41 London Dispersion Forces London dispersion forces are created when on molecule with a temporarily dipole causes another to become temporarily polar. ++ −− ++ −− ++ −−

42 Molecular Polarity Molecular Structure Courtesy Christy Johannesson

43 Electronegativity ++ –– 00 00 HClHH

44 Na Ionic vs. Covalent Ionic compounds form repeating units. Covalent compounds form distinct molecules. Consider adding to NaCl(s) vs. H 2 O(s): H O H Cl H O H H O H NaCl: atoms of Cl and Na can add individually forming a compound with million of atoms. H 2 O: O and H cannot add individually, instead molecules of H 2 O form the basic unit. Na Cl Na

45 Holding it together Q:Consider a glass of water. Why do molecules of water stay together? A:There must be attractive forces. Intramolecular forces occur between atoms Intermolecular forces occur between molecules Intermolecular forces are not considered in ionic bonding because there are no molecules. The type of intramolecular bond determines the type of intermolecular force. Intramolecular forces are much stronger

46 I’m not stealing, I’m sharing unequally We described ionic bonds as stealing electrons In fact, all bonds share – equally or unequally. Note how bonding electrons spend their time: Bonding electrons are shared in each compound, but are NOT always shared equally. The greek symbol  indicates “partial charge ”. H2H2 HClLiCl ++ –– 00 00 +– covalent (non-polar) polar covalent ionic HH H Cl [Li] + [ Cl ] –

47 Dipole Moment Direction of the polar bond in a molecule. Arrow points toward the more electronegative atom. H Cl ++ -- Courtesy Christy Johannesson

48 Dipole-induced dipole attraction The attraction between a dipole and an induced dipole.

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50 Oxygen, O 2

51 Nonpolar

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55 Water, H 2 O

56 -- -- ++ ++

57 -- -- ++ ++

58 -- -- ++ ++

59 -- -- ++ ++ ++ ++ -- --

60 -- -- ++ ++ ++ ++ -- --

61 ++ ++ -- -- -- -- ++ ++

62 ++ ++ -- -- -- -- ++ ++

63 ++ ++ -- -- -- -- ++ ++

64 ++ ++ -- -- ++ ++ induced dipole Dipole -- --

65 ++ ++ -- -- -- -- ++ ++

66 ++ ++ -- -- -- -- ++ ++

67 ++ ++ -- -- -- -- ++ ++

68 -- -- ++ ++ ++ ++ -- --

69 -- -- ++ ++ ++ ++ -- --

70 -- -- ++ ++

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78 ++ ++ -- -- ++ ++ -- --

79 ++ ++ -- -- ++ ++ -- -- induced dipole Dipole

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83 Polar Nonpolar Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

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85 Determining Molecular Polarity Depends on: –dipole moments –molecular shape Courtesy Christy Johannesson ++ –– ++ –– ++ –– ++ –– HCl ++ ––

86 BF 3 F F F B Determining Molecular Polarity Nonpolar Molecules –Dipole moments are symmetrical and cancel out. Courtesy Christy Johannesson

87 H2OH2O H H O Determining Molecular Polarity Polar Molecules –Dipole moments are asymmetrical and don’t cancel. net dipole moment Courtesy Christy Johannesson

88 CHCl 3 H Cl Determining Molecular Polarity Therefore, polar molecules have... –asymmetrical shape (lone pairs) or –asymmetrical atoms net dipole moment Courtesy Christy Johannesson

89 Dipole Moment Nonpolar Polar.. HH O C OO Bond dipoles Overall dipole moment = 0 Bond dipoles Overall dipole moment The overall dipole moment of a molecule is the sum of its bond dipoles. In CO 2 the bond dipoles are equal in magnitude but exactly opposite each other. The overall dipole moment is zero. In H 2 O the bond dipoles are also equal in magnitude but do not exactly oppose each other. The molecule has a nonzero overall dipole moment. Coulomb’s law  = Q r Dipole moment,  Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 315

90 .. Polar Bonds H Cl Polar A molecule has a zero dipole moment because their dipoles cancel one another. HH O Polar FF B F Nonpolar H H H N Polar Nonpolar FF Cl F F F Xe FF Cl C NonpolarPolar Cl H C H H

91 HF HCl HBr HI Mark Wirtz, Edward Ehrat, David L. Cedeno* How is the electron density distributed in these different molecules? Based on your comparison of the electron density distributions, which molecule should have the most polar bond, and which one the least polar? Arrange the molecules in increasing order of polarity.

92 CH 3 Cl CH 2 Cl 2 CHCl 3 CCl 4 Mark Wirtz, Edward Ehrat, David L. Cedeno* Describe how is the electron density distributed in these different molecules? Based on your comparison of the electron density distributions, which molecule(s) should be the most polar, and which one(s) the least polar? Arrange the molecules in increasing order of polarity.

93 Benzene NO 3 - Nitrobenzene Mark Wirtz, Edward Ehrat, David L. Cedeno*

94 2s 2p (x, y, z) carbon Mark Wirtz, Edward Ehrat, David L. Cedeno*

95 How does H 2 form? ++ The nuclei repel But they are attracted to electrons They share the electrons

96 Hydrogen Bond Formation 0.74 A H – H distance Energy (KJ/mol) Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 318 no interaction increased attraction balanced attraction & repulsion increased repulsion Potential Energy Diagram - Attraction vs. Repulsion (internuclear distance)

97 Covalent bonds Nonmetals hold onto their valence electrons. They can’t give away electrons to bond. Still want noble gas configuration. Get it by sharing valence electrons with each other. By sharing both atoms get to count the electrons toward noble gas configuration. 1s 2 2s 2 2p 6 3s 2 3p 6 …eight valence electrons (stable octet)

98 Covalent bonding Fluorine has seven valence electrons F A second atom also has seven F By sharing electrons …both end with full orbitals 8 Valence electrons 8 Valence electrons

99 Single Covalent Bond A sharing of two valence electrons. Only nonmetals and Hydrogen. Different from an ionic bond because they actually form molecules. Two specific atoms are joined. In an ionic solid you can’t tell which atom the electrons moved from or to.

100 Sigma bonding orbitals From s orbitals on separate atoms ++ s orbital +++ Sigma bonding molecular orbital

101 Sigma bonding orbitals From p orbitals on separate atoms p orbital Sigma bonding molecular orbital  

102 Pi bonding orbitals P orbitals on separate atoms      Pi bonding molecular orbital

103 Sigma and pi bonds All single bonds are sigma bonds A double bond is one sigma and one pi bond A triple bond is one sigma and two pi bonds.

104 Atomic Orbitals and Bonding Bonds between atoms are formed by electron pairs in overlapping atomic orbitals 1s : E Example: H 2 (H-H) –Use 1s orbitals for bonding Example: H 2 O –From VSEPR: bent, 104.5° angle between H atoms –Use two 2p orbitals for bonding? E 2s 2p 90° How do we explain the structure predicted by VSEPR using atomic orbitals? 2p 1s

105 Overlapping Orbitals Draw orbital diagrams for F + F, H + O, Li + F 1s1s2s2s2p2p 1s1s2s2s2p2p 1s1s2s2s2p2p 1s1s 1s1s F2F2 H2OH2O 1s1s2s2s1s1s2s2s2p2p LiF is ionic (metal + non-metal) F Li electron transfer 1+ 1-

106 e-e- 3p+3p+ lithium atom Li e-e- loss of one valence electron e-e- e-e- lithium ion Li + 3p+3p+ e-e- e-e- 9p+9p+ fluorine atom F e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- gain of one valence electron fluoride ion F 1- 10p + e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e-

107 Formation of Cation 3p+3p+ lithium atom Li e-e- loss of one valence electron e-e- e-e- lithium ion Li + 3p+3p+ e-e- e-e- e-e-

108 Formation of Anion 9p+9p+ fluorine atom F e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- gain of one valence electron fluoride ion F 1- 10p + e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e-

109 Formation of Ionic Bond fluoride ion F 1- 9p+9p+ e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- lithium ion Li + 3p+3p+ e-e- e-e-

110 Be H H BeH 2 s p First, the formation of BeH 2 using pure s and p orbitals. The formation of BeH 2 using hybridized orbitals. atomic orbitals Be sp H H sp atomic orbitals hybrid orbitals No overlap = no bond! spp Be HH All hybridized bonds have equal strength and have orbitals with identical energies. BeH 2 Be Be = 1s 2 2s 2

111 Hybrid Orbitals Ground-state Be atom 1s1s2s2s2p2p1s1s2s2s2p2p Be atom with one electron “promoted” s pxpx pypy pzpz sp hybrid orbitals Energy hybridize s orbital p orbital two sp hybrid orbitals sp hybrid orbitals shown together (large lobes only) 1s1ssp2p2p Be atom of BeH 2 orbital diagram HH Be n = 1 n = 2

112 Hybrid Orbitals 2s2s2p2p Ground-state B atom s pxpx pypy pzpz Energy sp 2 2p2p B atom of BH 3 orbital diagram hybridize s orbital 2s2s2p2p B atom with one electron “promoted” sp 2 hybrid orbitals p orbitals sp 2 hybrid orbitals shown together (large lobes only) three sp s hybrid orbitals H H H B

113 Hybridization …the blending of orbitals Valence bond theory is based on two assumptions: 1. The strength of a covalent bond is proportional to the amount of overlap between atomic orbitals; the greater the overlap, the more stable the bond. 2. An atom can use different combinations of atomic orbitals to maximize the overlap of orbitals used by bonded atoms.

114 We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds. Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it. Lets look at a molecule of methane, CH 4.

115 What is the expected orbital notation of carbon in its ground state? (Hint: How many unpaired electrons does this carbon atom have available for bonding?) Can you see a problem with this? Carbon ground state configuration 1s 2s 2p You should conclude that carbon only has TWO electrons available for bonding. That is not enough! How does carbon overcome this problem so that it may form four bonds?

116 The first thought that chemists had was that carbon promotes one of its 2s electrons… …to the empty 2p orbital. Carbon’s Empty Orbital 1s 2s 2p 1s 2s 2p 1s 2s 2p Non-hybridized orbitalhybridized orbital

117 However, they quickly recognized a problem with such an arrangement… Three of the carbon-hydrogen bonds would involve an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom. A Problem Arises 1s 2s 2p 1s Unequal bond energy But what about the fourth bond…?

118 The fourth bond is between a 2s electron from the carbon and the lone 1s hydrogen electron. less Such a bond would have slightly less energy than the other bonds in a methane molecule. Unequal bond energy #2 1s 2s 2p 1s

119 This bond would be slightly different in character than the other three bonds in methane. This difference would be measurable to a chemist by determining the bond length and bond energy. But is this what they observe?

120 The simple answer is, “No”. Chemists have proposed an explanation – they call it hybridization. Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy. Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane. Enter Hybridization

121 In the case of methane, they call the hybridization sp 3, meaning that an s orbital is combined with three p orbitals to create four equal hybrid orbitals. These new orbitals have slightly MORE energy than the 2s orbital… … and slightly LESS energy than the 2p orbitals. sp 3 Hybrid Orbitals

122 s pxpx pypy pzpz Carbon 1s 2 2s 2 2p 2 Carbon could only make two bonds if no hybridization occurs. However, carbon can make four equivalent bonds. sp 3 hybrid orbitals Energy sp 3 C atom of CH 4 orbital diagram B A B B B Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 321

123 Hybridization of s and p Orbitals The combination of an ns and an np orbital gives rise to two equivalent sp hybrids oriented at 180º. Combination of an ns and two or three np orbitals produces three equivalent sp 2 hybrids or four equivalent sp 3 hybrids. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

124 Hybridization of s and p Orbitals Both promotion and hybridization require an input of energy; the overall process of forming a compound with hybrid orbitals will be energetically favorable only if the amount of energy released by the formation of covalent bonds is greater than the amount of energy used to form the hybrid orbitals. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

125 Hybridization Involving d Orbitals 3s 3p 3d promote five sp 3 d orbitals 3d3d F F F P F F A BeBe BeBe BeBe BaBa BaBa Trigonal bipyramidal hybridize degenerate orbitals (all EQUAL) unhybridized P atom P = [Ne]3s 2 3p 3 vacant d orbitals

126 Pure atomic orbitals of central atom Hybridization of the central atom Number of hybrid orbitals Shape of hybrid orbitals s,p s,p,p s,p,p,p s,p,p,p,d s,p,p,p,d,d sp sp 2 sp 3 sp 3 d sp 3 d Linear Trigonal Planar Tetrahedral Trigonal Bipyramidal Octahedral Hybridization Animation, by Raymond Chang

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128 Bonding Single bonds –Overlap of bonding orbitals on bond axis –Termed “sigma” or σ bonds Double bonds –Sharing of electrons between 2 p orbitals perpendicular to the bonding atoms – Termed “pi” or π bonds 2p Bond Axis of σ bond One π bond

129 Multiple Bonds 2s 2p 2s 2p sp 2 2p promotehybridize CC H HH H C 2 H 4, ethene one  bond and one  bond H H C C H H      H H C C H H Two lobes of one  bond Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page

130 Multiple Bonds 2s 2p 2s 2p sp 2 2p promotehybridize CC H HH H C 2 H 4, ethene one  bond and one  bond H H C C H H      H H C C H H Two lobes of one  bond Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page CC H H sp 2 H H pp pp

131  bond Internuclear axis pp

132 Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 326 Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

133  bonds H C H C H C H C H C H C Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329 C 6 H 6 = benzene

134 2p atomic orbitals Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329

135  bonds and  bonds H C H C H C H C H C H C Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329

136  bonds H C H C H C H C H C H C H C H C Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329

137  bonds H C H C H C H C H C H C H C H C Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329

138 NN O O O O dinitrogen tetraoxide N2O4N2O4 N O O N O O h 2 NO 2 nitrogen dioxide (free radical) colorless red-brown 

139 1s1s1s1s 1s1s 1s1s 1s1s1s1s 1s1s 1s1s H 2 molecule He 2 molecule H atom He atom Energy Energy-level diagram for (a) the H 2 molecule and (b) the hypothetical He 2 molecule (a) (b) Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 332

140 Bond Order Bond order = ½ (# or bonding electrons - # of antibonding electrons) A bond order of 1 represents a single bond, A bond order of 2 represents a double bond, A bond order of 3 represents a triple bond. A bond order of 0 means no bond exists. Because MO theory also treats molecules with an odd number of electrons, Bond orders of 1/2, 3/2, or 5/2 are possible.

141 1s21s2 1s21s2 1s1s 1s1s 2s12s1 2s12s1 2s2s 2s2s Energy-level diagram for the Li 2 molecule Li = 1s 2 2s 1 Li Li 2 Energy Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 334

142 2s  2s   2s 2p  2p   2p  2p   2p Energy-level diagram for molecular orbitals of second-row homonuclear diatomic molecules. Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 337

143 Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 338

144  2s   2s  2p  2p Energy of molecular orbitals O 2, F 2, Ne 2 B 2, C 2, N 2 Increasing 2s – 2p interaction Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 338

145  2s  2s  2p  2p  2p  2p Large 2s – 2p interaction Small 2s – 2p interaction  2s  2s  2p  2p  2p  2p C2C2 N2N2 B2B2 F2F2 Ne 2 O2O2 Bond order Bond enthalpy (kJ/mol) Bond length (angstrom) Magnetic behavior Paramagnetic Diamagnetic Diamagnetic Paramagnetic Diamagnetic _____ Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 339

146  2s   2s  2px   2px  2py   2py  2p   2p C2C2 Arrange the atomic and molecular orbitals in order of increasing energy. How many orbitals are per molecule? Can you distinguish the bonding from the antibonding MOs? Mark Wirtz, Edward Ehrat, David L. Cedeno*

147 Magnetic Properties of a Sample PARAMAGNETISM – molecules with one or more unpaired electrons are attracted into a magnetic field. (appears to weigh MORE in a magnetic field) DIAMAGNETISM – substances with no unpaired electrons are weakly repelled from a magnetic field. (appears to weigh LESS in a magnetic field)

148 Experiment for determining the magnetic properties of a sample The sample is first weighed in the absence of a magnetic field. When a field is applied, a diamagnetic sample tends to move out of the field and appears to have a lower mass. A paramagnetic sample is drawn into the field and thus appears to gain mass. Paramagnetism is a much stronger effect than is diamagnetism. sample NS NS NS Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 339

149 Experiment for determining the magnetic properties of a sample NS NS The sample is first weighed in the absence of a magnetic field. When a field is applied, a diamagnetic sample tends to move out of the field and appears to have a lower mass. A paramagnetic sample is drawn into the field and thus appears to gain mass. Paramagnetism is a much stronger effect than is diamagnetism. sample Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 339

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152 lone Pair single bond double bond triple bond Electron Domains Determine the shape of the BCl 3 molecule: B Cl :: : : : :: : : There are 3 electron domains about the central atom: no lone pairs and three single bonds. Three electron domains arrange themselves in a trigonal plane, with 120 o angles. We predict a trigonal planar geometry. : B Cl :: : : : :: : Electron-domain geometry: trigonal planar Molecular geometry (shape): trigonal planar

153 One s orbital Two p orbitals Three sp 2 hybrid orbitals sp 2 hybrid orbitals shown together (large lobes only) Hybridize

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156 IIIIII Courtesy Christy Johannesson Molecular Geometry

157 VSEPR Theory Valence Shell Electron Pair Repulsion Theory Electron pairs orient themselves in order to minimize repulsive forces. Courtesy Christy Johannesson

158 VSEPR Theory Types of e - Pairs  Bonding pairs - form bonds  Lone pairs - nonbonding electrons Lone pairs repel more strongly than bonding pairs!!! Courtesy Christy Johannesson

159 VSEPR Theory Lone pairs reduce the bond angle between atoms. Courtesy Christy Johannesson bond angle

160 Draw the Lewis Diagram. Tally up e - pairs on central atom.  double/triple bonds = ONE pair Shape is determined by the # of bonding pairs and lone pairs. Know the 8 common shapes & their bond angles! Determining Molecular Shape Courtesy Christy Johannesson

161 Common Molecular Shapes 2 total 2 bond 0 lone LINEAR 180° BeH 2 Courtesy Christy Johannesson BBA

162 3 total 3 bond 0 lone TRIGONAL PLANAR 120° BF 3 Common Molecular Shapes Courtesy Christy Johannesson B B A B

163 Common Molecular Shapes 3 total 2 bond 1 lone BENT <120° SO 2 Courtesy Christy Johannesson

164 4 total 4 bond 0 lone TETRAHEDRAL 109.5° CH 4 Common Molecular Shapes Courtesy Christy Johannesson B A B B B

165 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107° NH 3 Common Molecular Shapes Courtesy Christy Johannesson

166 Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

167 Ammonia, NH 3

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180 Triangular pyramidal Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

181 4 total 2 bond 2 lone BENT 104.5° H2OH2O Common Molecular Shapes Courtesy Christy Johannesson H x H x O H x H x O

182 5 total 5 bond 0 lone TRIGONAL BIPYRAMIDAL 120°/90° PCl 5 Common Molecular Shapes Courtesy Christy Johannesson A BeBe BeBe BeBe BaBa BaBa

183 6 total 6 bond 0 lone OCTAHEDRAL 90° SF 6 Common Molecular Shapes Courtesy Christy Johannesson B B B B B B A

184 PF 3 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107° F P F F Examples Courtesy Christy Johannesson

185 CO 2 O C O 2 total 2 bond 0 lone LINEAR 180° Examples Courtesy Christy Johannesson

186 Introduction to Bonding Courtesy Christy Johannesson

187 l Chemical bond — the force that holds atoms together in a chemical compound l Covalent bonding — electrons are shared between atoms in a molecule or polyatomic ion l Ionic bonding — positively and negatively charged ions are held together by electrostatic forces l Ionic compounds — dissolve in water to form aqueous solutions that conduct electricity l Covalent compounds — dissolve to form solutions that do not conduct electricity

188 Vocabulary lClChemical Bond –a–attractive force between atoms or ions that binds them together as a unit –b–bonds form in order to… decrease potential energy (PE) increase stability Courtesy Christy Johannesson

189 Vocabulary CHEMICAL FORMULA molecular formula unit IONICCOVALENT CO 2 NaCl Courtesy Christy Johannesson

190 Vocabulary COMPOUND ternary compound binary compound 2 elements more than 2 elements NaNO 3 NaCl Courtesy Christy Johannesson

191 Vocabulary ION polyatomic Ion monatomic Ion 1 atom 2 or more atoms NO 3 - Na + Courtesy Christy Johannesson

192 IONIC COVALENT Bond Formation Type of Structure Solubility in Water Electrical Conductivity Other Properties e - are transferred from metal to nonmetal high yes (solution or liquid) yes e - are shared between two nonmetals low no usually not Melting Point crystal lattice true molecules Types of Bonds Physical State solid liquid or gas odorous Courtesy Christy Johannesson

193 “electron sea” METALLIC Bond Formation Type of Structure Solubility in Water Electrical Conductivity Other Properties Melting Point Types of Bonds Physical State e - are delocalized among metal atoms very high yes (any form) no malleable, ductile, lustrous solid Courtesy Christy Johannesson

194 Ionic Bonding - Crystal Lattice Types of Bonds Table salt

195 Ionic Bonding - Crystal Lattice Types of Bonds Table salt

196 Lattice Energies in Ionic Solids Ionic compounds 1. Usually rigid, brittle, crystalline substances with flat surfaces that intersect at characteristic angles 2. Not easily deformed 3. Melt at relatively high temperatures 4. Properties result from the regular arrangement of the ions in the crystalline lattice and from the strong electrostatic attractive forces between ions with opposite charges

197 Covalent Bonding - True Molecules Types of Bonds Diatomic Molecule Courtesy Christy Johannesson

198 Metallic Bonding - “Electron Sea” Types of Bonds

199 Bond Polarity l Most bonds are a blend of ionic and covalent characteristics. l Difference in electronegativity determines bond type. Courtesy Christy Johannesson Ionic Polar-covalent Nonpolar-covalent % 50% 5% 0% Difference in electronegativities Percentage ionic character

200 Types of Chemical Bonds Copyright © 2006 Pearson Education Inc., publishing as Benjamin Cummings

201 Bond Polarity l Electronegativity –Attraction an atom has for a shared pair of electrons. –higher e - neg atom   - –lower e - neg atom   + Courtesy Christy Johannesson

202 Bond Polarity l Electronegativity Trend –Increases up and to the right. Courtesy Christy Johannesson

203 Ionic bonding: Li + Cl Ionic bonding (stealing/transfer of electrons) can be represented in three different ways Li + Cl  [Li] + [Cl] – 1e - 3p + 4n 0 2e - 17p + 18n 0 8e - 8e - 2e 3p + 4n 0 2e - 1e - 17p + 18n 0 7e - 8e - 2e - Li Cl [ Cl ] – [Li] + lithium atomchlorine atomlithium ionchlorine ionchloride ion

204 Ionic bonding: Mg + O Mg + O  [Mg] 2+ [O] 2– 12p + 12n 0 2e - 8e - 2e - 1e - [ O ] 2– [Mg] 2+ 6e - 2e - 8n 0 8p + 1e - 8e - 2e - 8n 0 8p + 12p + 12n 0 2e - 8e - O Mg

205 F 4.0 Ar -- Kr 3.0 Xe 2.6 Rn 2.4 Bond Polarity l Electronegativity Trend –Increases up and to the right Be 1.5 Al 1.5 Si 1.8 Ti 1.5 V 1.6 Cr 1.6 Mn 1.5 Fe 1.8 Co 1.8 Ni 1.8 Cu 1.9 Zn 1.7 Ga 1.6 Ge 1.8 Nb 1.6 Mo 1.8 Tc 1.9 Ag 1.9 Cd 1.7 In 1.7 Sn 1.8 Sb 1.9 Ta 1.5 W 1.7 Re 1.9 Hg 1.9 Tl 1.8 Pb 1.8 Bi 1.9 N 3.0 O 3.5 F 4.0 Cl 3.0 C 2.5 S 2.5 Br 2.8 I 2.5 Na 0.9 K 0.8 Rb 0.8 Cs 0.7 Ba 0.9 Fr 0.7 Ra 0.9 H 2.1 B 2.0 P 2.1 As 2.0 Se 2.4 Ru 2.2 Rh 2.2 Pd 2.2 Te 2.1 Os 2.2 Ir 2.2 Pt 2.2 Au 2.4 Po 2.0 At 2.2 Period Actinides: Li 1.0 Ca 1.0 Sc 1.3 Sr 1.0 Y 1.2 Zr 1.4 Hf 1.3 Mg 1.2 La 1.1 Ac 1.1 Lanthanides:     Ne -- He --

206 F Ar Kr Xe Rn 2.4 Bond Polarity l Electronegativity Trend –Increases up and to the right Be Al Si Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge Nb Mo Tc Ag Cd In Sn Sb Ta W Re Hg Tl Pb Bi N O F Cl C S Br I Na K Rb Cs Ba Fr Ra H B P As Se Ru Rh Pd Te Os Ir Pt Au Po At 2.2 Period Actinides: Li Ca Sc Sr Y Zr Hf Mg La Ac 1.1 Lanthanides:     Ne -- 2 He --

207 F 4.0 Ar -- Kr 3.0 Xe 2.6 Rn 2.4 Bond Polarity l Electronegativity Trend –Increases up and to the right. Be 1.5 Al 1.5 Si 1.8 Ti 1.5 V 1.6 Cr 1.6 Mn 1.5 Fe 1.8 Co 1.8 Ni 1.8 Cu 1.9 Zn 1.7 Ga 1.6 Ge 1.8 Nb 1.6 Mo 1.8 Tc 1.9 Ag 1.9 Cd 1.7 In 1.7 Sn 1.8 Sb 1.9 Ta 1.5 W 1.7 Re 1.9 Hg 1.9 Tl 1.8 Pb 1.8 Bi 1.9 N 3.0 O 3.5 F 4.0 Cl 3.0 C 2.5 S 2.5 Br 2.8 I 2.5 Na 0.9 K 0.8 Rb 0.8 Cs 0.7 Ba 0.9 Fr 0.7 Ra 0.9 H 2.1 B 2.0 P 2.1 As 2.0 Se 2.4 Ru 2.2 Rh 2.2 Pd 2.2 Te 2.1 Os 2.2 Ir 2.2 Pt 2.2 Au 2.4 Po 2.0 At 2.2 Actinides: Li 1.0 Ca 1.0 Sc 1.3 Sr 1.0 Y 1.2 Zr 1.4 Hf 1.3 Mg 1.2 La 1.1 Ac 1.1 Lanthanides:     Ne -- He --

208 Bond Polarity l Electronegativity Trend –Increases up and to the right A 2A 3B 4B 5B 6B 7B 1B 2B 3A 4A 5A 6A 7A 8A 8B

209 l Nonpolar Covalent Bond –electrons are shared equally –symmetrical electron density –usually identical atoms Bond Polarity

210 ++ -- l Polar Covalent Bond –electrons are shared unequally –asymmetrical e - density –results in partial charges (dipole) Courtesy Christy Johannesson

211 l Nonpolar l Polar l Ionic Bond Polarity Courtesy Christy Johannesson

212 Bond Polarity Examples: l Cl 2 l HCl l NaCl = 0.0 Nonpolar = 0.9 Polar = 2.1 Ionic Polar-covalent Nonpolar-covalent % 50% 5% 0% Difference in electronegativities Percentage ionic character

213 Mg Write the electron dot diagram for l Na l Mg lClC lOlO lFlF l Ne l He 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 1s22s22p21s22s22p2 1s22s22p41s22s22p4 1s22s22p51s22s22p5 1s22s22p61s22s22p6 1s21s2 Na C O F He Ne

214 Ionic Bonding NaCl transfer of electron + - NaCl

215 Ca +2 P -3 Ca +2 P lAlAll the electrons must be accounted for! Ionic Bonding Ca -3

216 Ionic Bonding Ca 3 P 2 Formula Unit Ca 2+ P 3- Ca 2+ P 3- Ca 2+ P 3- Ca 2+

217 Metals are Malleable lHlHammered into shape (bend). lDlDuctile - drawn into wires. lElElectrons allow atoms to slide by

218 Ionic solids are brittle Force Strong repulsion breaks crystal apart.

219 How does H 2 form? ++ The nuclei repel But they are attracted to electrons They share the electrons

220 Hydrogen Bond Formation 0.74 A H – H distance Energy (KJ/mol) Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 318 no interaction increased attraction balanced attraction & repulsion increased repulsion Potential Energy Diagram - Attraction vs. Repulsion (internuclear distance)

221 Covalent bonds l Nonmetals hold onto their valence electrons. l They can’t give away electrons to bond. l Still want noble gas configuration. l Get it by sharing valence electrons with each other. l By sharing both atoms get to count the electrons toward noble gas configuration.

222 Covalent bonding FF lFlFluorine has seven valence electrons lAlA second F atom also has seven lBlBy sharing electrons lBlBoth end with full orbitals (stable octets) 8 Valence electrons

223 Single Covalent Bond l A sharing of two valence electrons. l Only nonmetals and Hydrogen. l Different from an ionic bond because they actually form molecules. l Two specific atoms are joined. l In an ionic solid you can’t tell which atom the electrons moved from or to.

224 How to show how they formed l It’s like a jigsaw puzzle. l I have to tell you what the final formula is. l You put the pieces together to end up with the right formula. l For example - show how water is formed with covalent bonds.

225 Water H O Each hydrogen has 1 valence electron Each hydrogen wants 1 more The oxygen has 6 valence electrons The oxygen wants 2 more They share to make each other happy

226 l Put the pieces together l The first hydrogen is happy l The oxygen still wants one more HO

227 Water l The second hydrogen attaches l Every atom has full energy levels l A pair of electrons is a single bond HO H H HO

228 Lewis Structures 1) Count up total number of valence electrons 2) Connect all atoms with single bonds - “multiple” atoms usually on outside - “single” atoms usually in center; C always in center, H always on outside. 3) Complete octets on exterior atoms (not H, though) 4) Check - valence electrons math with Step 1 - all atoms (except H) have an octet;if not, try multiple bonds - any extra electrons?Put on central atom

229 Multiple Bonds l Sometimes atoms share more than one pair of valence electrons. l A double bond is when atoms share two pair (4) of electrons. l A triple bond is when atoms share three pair (6) of electrons.

230 Carbon dioxide l CO 2 - Carbon is central atom ( I have to tell you) l Carbon has 4 valence electrons l Wants 4 more l Oxygen has 6 valence electrons l Wants 2 more OC

231 Carbon dioxide l Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short O C

232 Carbon dioxide l Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short O C O

233 l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond 8 valence electrons Carbon dioxide O CO

234 Water H O Each hydrogen has 1 valence electron Each hydrogen wants 1 more The oxygen has 6 valence electrons The oxygen wants 2 more They share to make each other happy

235 Put the pieces together The first hydrogen is happy The oxygen still wants one more HO

236 Water The second hydrogen attaches Every atom has full energy levels A pair of electrons is a single bond HO H H HO

237 Lewis Structures 1) Count up total number of valence electrons 2) Connect all atoms with single bonds - “multiple” atoms usually on outside - “single” atoms usually in center; C always in center, H always on outside. 3) Complete octets on exterior atoms (not H, though) 4) Check - valence electrons math with Step 1 - all atoms (except H) have an octet;if not, try multiple bonds - any extra electrons?Put on central atom

238 Multiple Bonds Sometimes atoms share more than one pair of valence electrons. A double bond is when atoms share two pair (4) of electrons. A triple bond is when atoms share three pair (6) of electrons.

239 Formation of Multiple Covalent Bonds By combining more than one unpaired electron at a time, a double bond is formed. Both oxygen atoms end up with eight valence electrons. O x x x x O x x x x x x O x x O

240 Carbon dioxide CO 2 - Carbon is central atom ( I have to tell you) Carbon has 4 valence electrons Wants 4 more Oxygen has 6 valence electrons Wants 2 more OC

241 Carbon dioxide Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short O C

242 Carbon dioxide l Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short O C O

243 l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond 8 valence electrons Carbon dioxide O CO

244 How to draw them l Add up all the valence electrons. l Count up the total number of electrons to make all atoms happy. l Subtract. l Divide by 2 l Tells you how many bonds - draw them. l Fill in the rest of the valence electrons to fill atoms up.

245 Examples l NH 3 l N - has 5 valence electrons wants 8 l H - has 1 valence electrons wants 2 l NH 3 has 5+3(1) = 8 l NH 3 wants 8+3(2) = 14 l (14-8)/2= 3 bonds l 4 atoms with 3 bonds N H

246 NHH H Examples l Draw in the bonds l All 8 electrons are accounted for l Everything is full

247 Examples l HCN C is central atom l N - has 5 valence electrons wants 8 l C - has 4 valence electrons wants 8 l H - has 1 valence electrons wants 2 l HCN has = 10 l HCN wants = 18 l ( ) / 2= 4 bonds l 3 atoms with 4 bonds -will require multiple bonds - not to H

248 HCN l Put in single bonds l Need 2 more bonds l Must go between C and N NHC

249 HCN l Put in single bonds l Need 2 more bonds l Must go between C and N l Uses 8 electrons - 2 more to add NHC

250 HCN l Put in single bonds l Need 2 more bonds l Must go between C and N l Uses 8 electrons - 2 more to add l Must go on N to fill octet NHC

251 Another way of indicating bonds l Often use a line to indicate a bond l Called a structural formula l Each line is 2 valence electrons HHO = HHO

252 Structural Examples H CN C O H H l C has 8 electrons because each line is 2 electrons l Ditto for N l Ditto for C here l Ditto for O

253 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide l CO OC

254 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide l CO OC

255 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide l CO OC

256 How do we know if l Have to draw the diagram and see what happens. l Often happens with polyatomic ions and acids.

257 Resonance l When more than one dot diagram with the same connections are possible. l NO 2 - l Which one is it? l Does it go back and forth. l It is a mixture of both, like a mule. l NO 3 -

258 VSEPR l Valence Shell Electron Pair Repulsion. l Predicts three dimensional geometry of molecules. l Name tells you the theory. l Valence shell - outside electrons. l Electron Pair repulsion - electron pairs try to get as far away as possible. l Can determine the angles of bonds.

259 VSEPR l Based on the number of pairs of valence electrons both bonded and unbonded. l Unbonded pair are called lone pair. l CH 4 - draw the structural formula l Has 4 + 4(1) = 8 l wants 8 + 4(2) = 16 l (16-8)/2 = 4 bonds

260 VSEPR l Single bonds fill all atoms. l There are 4 pairs of electrons pushing away. l The furthest they can get away is 109.5º. CHH H H

261 4 atoms bonded l Basic shape is tetrahedral. l A pyramid with a triangular base. l Same shape for everything with 4 pairs. C HH H H 109.5º

262 3 bonded - 1 lone pair NHH H N HH H <109.5º l Still basic tetrahedral but you can’t see the electron pair. l Shape is called trigonal pyramidal.

263 2 bonded - 2 lone pair OH H O H H <109.5º l Still basic tetrahedral but you can’t see the 2 lone pair. l Shape is called bent.

264 3 atoms no lone pair C H H O l The farthest you can the electron pair apart is 120º

265 3 atoms no lone pair C H H O l The farthest you can the electron pair apart is 120º. l Shape is flat and called trigonal planar. C H HO 120º

266 2 atoms no lone pair l With three atoms the farthest they can get apart is 180º. l Shape called linear. C O O 180º

267 Hybrid Orbitals Combines bonding with geometry

268 Hybridization l The mixing of several atomic orbitals to form the same number of hybrid orbitals. l All the hybrid orbitals that form are the same (degenerate = equal energy). l sp 3 - one s and three p orbitals mix to form four sp 3 orbitals. l sp 2 - one s and two p orbitals mix to form three sp 2 orbitals leaving one p orbital. l sp - one s and one p orbitals mix to form four sp orbitals leaving two p orbitals.

269 Hybridization l We blend the s and p-orbitals of the valence electrons and end up with the tetrahedral geometry. We combine one s orbital and three p-orbitals. l sp 3 hybridization has tetrahedral geometry.

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272 sp 3 geometry 109.5º l This leads to tetrahedral shape. l Every molecule with a total of 4 atoms and lone pair is sp 3 hybridized. l Gives us trigonal pyramidal and bent shapes also.

273 How we get to hybridization l We know the geometry from experiment. l We know the orbitals of the atom hybridizing atomic orbitals can explain the geometry. l So if the geometry requires a tetrahedral shape, it is sp 3 hybridized. –This includes bent and trigonal pyramidal molecules because one of the sp 3 lobes holds the lone pair.

274 sp 2 hybridization l C 2 H 4 l double bond acts as one pair l trigonal planar l Have to end up with three blended orbitals –use one s and two p orbitals to make three sp 2 orbitals. –leaves one p orbital perpendicular

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277 Where is the P orbital? l Perpendicular l The overlap of orbitals makes a sigma bond (  bond)

278 Two types of Bonds l Sigma bonds from overlap of orbitals between the atoms Pi bond (  bond) above and below atoms l Between adjacent p orbitals. l The two bonds of a double bond

279 CC H H H H

280 H H H B sp 2 hybridization l when three things come off atom l trigonal planar l 120º one  bond B B A B trigonal planar hybridize s orbital p orbitals three sp s hybrid orbitals

281 What about two l when two things come off l one s orbital and one p orbital hybridize l linear

282 sp hybridization l end up with two lobes 180º apart. l p orbitals are at right angles makes room for two  bonds and two sigma bonds. l a triple bond or two double bonds

283 CO 2 C can make two  and two  O can make one  and one  COO

284 N2N2

285 N2N2

286 Polar Bonds l When the atoms in a bond are the same, the electrons are shared equally. l This is a nonpolar covalent bond. l When two different atoms are connected, the atoms may not be shared equally. l This is a polar covalent bond. l How do we measure how strong the atoms pull on electrons?

287 Electronegativity l A measure of how strongly the atoms attract electrons in a bond. l The bigger the electronegativity difference the more polar the bond Covalent nonpolar Covalent moderately polar Covalent polar >2.0 Ionic

288 How to show a bond is polar l Isn’t a whole charge just a partial charge  means a partially positive  means a partially negative l The Cl pulls harder on the electrons l The electrons spend more time near the Cl H Cl  

289 Polar Molecules Molecules with ‘ends’

290 Polar Molecules l Molecules with a positive and a negative end l Requires two things to be true ¬ The molecule must contain polar bonds This can be determined from differences in electronegativity. ­ Symmetry can not cancel out the effects of the polar bonds. Must determine geometry first.

291 Is it polar?.. H Cl Polar HH O FF B F Nonpolar H H H N Polar Nonpolar FF Cl F F F Xe FF NonpolarPolar Cl C H C H H XeF 4 CCl 4 CH 3 Cl HClH2OH2O BF 3 NH 3

292 Bond Dissociation Energy l The energy required to break a bond l C - H kJ C + H l We get the Bond dissociation energy back when the atoms are put back together If we add up the BDE of the reactants and subtract the BDE of the products we can determine the energy of the reaction (  H)

293 Find the energy change for the reaction l CH 4 + 2O 2 CO 2 + 2H 2 O l For the reactants we need to break 4 C-H bonds at 393 kJ/mol and 2 O=O bonds at 495 kJ/mol= 2562 kJ/mol l For the products we form 2 C=O at 736 kJ/mol and 4 O-H bonds at 464 kJ/mol l = 3328 kJ/mol l reactants - products = = -766kJ

294 Intermolecular Forces What holds molecules to each other?

295 Intermolecular Forces l They are what make solid and liquid molecular compounds possible. l The weakest are called van derWaal’s forces - there are two kinds l Dispersion forces l Dipole Interactions –depend on the number of electrons –more electrons stronger forces –bigger molecules

296 l Depend on the number of electrons l More electrons stronger forces l Bigger molecules more electrons fluorine (F 2 ) is a gas bromine (Br 2 ) is a liquid iodine (I 2 ) is a solid Dipole interactions

297 l Occur when polar molecules are attracted to each other. l Slightly stronger than dispersion forces. l Opposites attract but not completely hooked like in ionic solids.

298 Dipole interactions l Occur when polar molecules are attracted to each other. l Slightly stronger than dispersion forces. l Opposites attract but not completely hooked like in ionic solids. HFHF  HFHF 

299 Dipole Interactions    

300 Hydrogen bonding l Are the attractive force caused by hydrogen bonded to F, O, or N. l F, O, and N are very electronegative so it is a very strong dipole. l The hydrogen partially share with the lone pair in the molecule next to it. l The strongest of the intermolecular forces.

301 Hydrogen Bonding H H O ++ -- ++ H H O ++ -- ++

302 Hydrogen bonding H H O H H O H H O H H O H H O H H O H H O

303 Resources - Bonding Objectives Episode 8 Episode 8 – Chemical Bonds Episode 9 Episode 9 – Molecular Architecture


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