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“A” students work (without solutions manual) ~ 10 problems/night. Alanah Fitch Flanner Hall 402 508-3119 Office Hours W – F 2-3 pm Module.

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Presentation on theme: "“A” students work (without solutions manual) ~ 10 problems/night. Alanah Fitch Flanner Hall 402 508-3119 Office Hours W – F 2-3 pm Module."— Presentation transcript:

1 “A” students work (without solutions manual) ~ 10 problems/night. Alanah Fitch Flanner Hall Office Hours W – F 2-3 pm Module #10: Covalent Bonding

2 FITCH Rules G1: Suzuki is Success G2. Slow me down G3. Scientific Knowledge is Referential G4. Watch out for Red Herrings G5. Chemists are Lazy C1. It’s all about charge C2. Everybody wants to “be like Mike” (grp.18) C3. Size Matters C4. Still Waters Run Deep C5. Alpha Dogs eat first Chemistry General

3 Properties and Measurements Property Unit Reference State Size msize of earth Volumecm 3 m Weightgrammass of 1 cm 3 water at specified Temp (and Pressure) Temperature o C, Kboiling, freezing of water (specified Pressure) x gamu(mass of 1C-12 atom)/12 quantitymoleatomic mass of an element in grams Pressureatm, mm Hgearth’s atmosphere at sea level Energy: Thermal BTU1 lb water 1 o F calorie1 g water 1 o C KineticJ2kg mass moving at 1m/s Energy, of electronsenergy of electron in a vacuum ElectronegativityF

4 1.Bonding = sharing –electrons between repulsive + nuclei 2.Lewis Dot structures help us visualize sharing of electrons Octets Double and triple bonds Resonance structures and No Clean Socks 3.Formal Charge to help distinguish between alternatives 4.Violations of the Octet Rule 2 electrons >8 electrons 5.Using electrons to predict the SHAPE of the molecules VESPR Effect of unpaired electrons on the central atom on molecular shape Effect of Multiple bonds How to deal with “no central atom” 6.Bond polarity Effect of electronegativity difference between atoms in bond Effect of molecular shape How to symbolize bond polarity 7.Discrepancies between Electron Orbital Diagrams and VESPR Introducing orbital hybridization, s, sp, sp 2, sp 3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

5 Covalent bonding Patterns in abundance suggest a.periodicity b.preferred electronic configuration of elements Leading to the Rule: “Everybody wants to be “Like Mike” a.Ions: Groups 16 and 17 gain electrons; Groups 1 and 2 lose b.Other atoms share electrons to have eight electrons = COVALENT BONDING

6 Covalent Bonding – getting to a noble gas electron configuration by sharing electrons Repulsion of two hydrogen atoms with their Proton core Bring two elements close together 2.When very close the positive nuclei repel each other

7 Repulsion of two hydrogen atoms with their proton core Repulsion is high close where Protons see each other + + e e Atoms which are far apart Do not even see each other There is no energy, repulsive Or attractive between the two Repulsive energy Attractive energy Repulsion is low where Electrons shield nucleus, and where Electrons can be stabilized by both Positive charges + e + + e e + e Electrons are the jelly and peanut butter between the slices of bread (protons)

8 “A” students work (without solutions manual) ~ 10 problems/night. Alanah Fitch Flanner Hall Office Hours W – F 2-3 pm

9 1.Bonding = sharing –electrons between repulsive + nuclei 2.Lewis Dot structures help us visualize sharing of electrons Octets Double and triple bonds Resonance structures and No Clean Socks 3.Formal Charge to help distinguish between alternatives 4.Violations of the Octet Rule 2 electrons >8 electrons 5.Using electrons to predict the SHAPE of the molecules VESPR Effect of unpaired electrons on the central atom on molecular shape Effect of Multiple bonds How to deal with “no central atom” 6.Bond polarity Effect of electronegativity difference between atoms in bond Effect of molecular shape How to symbolize bond polarity 7.Discrepancies between Electron Orbital Diagrams and VESPR Introducing orbital hybridization, s, sp, sp 2, sp 3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

10 Gilbert Newton Lewis ; Caltech Physical Chemist Lewis dot structure (electron dot Structure or diagram) are diagrams that Show the bonding between atoms of A molecule based on shared “valence” shell (outer shell) electrons and shows the Presence of any “lone pair” of electrons That may exist in the covalently bonded Molecule. Covalent Valence = outermost shell electrons of an atom Latin: valere – to be strong shared

11 When the two hydrogen atoms are together, the electron configuration Looks like? He When a hydrogen atom and a fluorine atom share electrons, the Electron configuration on fluorine looks like? The “inner” shell electrons do not Show in this diagram

12 When the two hydrogen atoms are together, the electron configuration Looks like? He When a hydrogen atom and a fluorine atom share electrons, the Electron configuration on fluorine looks like? Only the “outer-most” or valence shell electrons Show in this Lewis Dot Structure How many valence electrons?: = last number in group

13 When the two hydrogen atoms are together, the electron configuration Looks like? He When a hydrogen atom and a fluorine atom share electrons, the Electron configuration on fluorine looks like? The shared pair of Electrons = covalent bond The unshared pairs of electrons are “regions of high Charge density”

14 When a hydrogen atom and an oxygen atom share valence electrons plus an Extra electron, the electron configuration on hydrogen and oxygen look like? The single electron pair shaired between the two bonded atoms Is called a single bond It is drawn as a line. Lewis dot structure for hydroxide Invoking Rule: Chemists are Lazy the diagram above is too tedious to write out all the time make shared electrons (bond) a line Valence electrons on oxygen?Valence electrons on hydrogen?

15 When two hydrogen atoms and an oxygen atom share valence electrons, the electron configuration on hydrogen and oxygen look like? Two shared electron pairs = Two single bonds

16 When two hydrogen atoms and an oxygen atom share valence electrons, the electron configuration on hydrogen and oxygen look like? When three hydrogen atoms and a nitrogen atom share valence electrons, the electron configuration on hydrogen and nitrogen look like?

17 When two hydrogen atoms and an oxygen atom share valence electrons, the electron configuration on hydrogen and oxygen look like? When three hydrogen atoms and a nitrogen atom share valence electrons, the electron configuration on hydrogen and nitrogen look like? Valence shell of nitrogen? Three pairs of shared electrons = three single bonds

18 When four hydrogen atoms and two carbon atoms share valence electrons, the electron configuration on hydrogen and carbon look like? Two electron pairs shared is a Double bond Valence shell of carbon?

19 When two hydrogen atoms and two carbon atoms share valence electrons, the electron configuration on hydrogen and carbon look like? Three electron pairs shared is a Triple bond

20 Rules for Writing Lewis Dot Structures 1.Count the number of valence electrons (last number of group) of all atoms a.For an anion add the appropriate extra number of electrons b.For a cation subtract the appropriate extra number of electrons 2.Draw a molecular skeleton, joining by single bonds to the central atom. a.The central is usually the atom written first in the formula (N in NH 4 +, S in SO 2, and C in CCl 4 ). b. The terminal atoms are usually H, O. c.Halogens are always terminal atoms. 3.Determine the number of valence electrons still available for distribution after subtracting two electrons for each single bond. 4.Determine the number of electrons required to complete the octet a.H gets only two electrons b.Other exceptions to be noted below 5.Fill in the region required for the octet. 6.Make up deficit of electrons by creating double bonds a.C, N, O, S H can only have one bond because it can share only one Electron. Poor H. Halogens have lots of electrons but really do not like to share. Greedy halogens All they want is one more to make up the Mike configuration

21 Draw Lewis structures of a)Hypochlorite ion b)Methyl alcohol c)N 2 d)SO 2 Hypochlorite? Hypo – smallest number of oxygens OCl - Valence shell electrons? Skeleton O 6 +Cl7 +Negative charge 1 Total electrons14 -1Single bond (6 electrons for O,Cl)12 remaining0

22 Draw Lewis structures of a)Hypochlorite ion b)Methyl alcohol, CH 3 OH c)N 2 d)SO 2 Valence shell electrons? Skeleton Carbon is first in formula Hydrogen is always terminal Octets Carbon has its octet Hydrogen has its duet Oxygen requires 4 more electrons O 6 +C4 +4(H)4 Negative charge 0 Total electrons14 -5single bonds-10 remaining 4 -octet for oxygen-4 remaining0

23 Draw Lewis structures of a)Hypochlorite ion b)Methyl alcohol, CH 3 OH c)N 2 d)SO 2 Valence shell electrons? Skeleton OctetsEach nitrogen requires 6 more 2N10 Negative charge 0 Total electrons10 -1single bond-2 Remaining8 Octet completion-12 Difference-4 2N10 Negative charge 0 Total electrons10 -3 single bonds-6 Remaining4 We are short 4 electrons for the octet, The only way to get extra ones is to Share four more electrons = triple Bond. Place the remaining 4 electrons equally On the two equal nitrogens

24 Draw Lewis structures of a)Hypochlorite ion b)Methyl alcohol, CH 3 OH c)N 2 d)SO 2 Valence shell electrons? Skeleton, First atom in formula is central Octets We are short 2 electrons for the octet, The only way to get extra ones is to Share two more electrons = double bond. Place the remaining 12 electrons to fill octets 2O12 +1S 6 Negative charge 0 Total electrons18 -2(single bonds)-4 Remaining electrons14 Octet for S-4 2(Octet for each O)-12 Deficit?-2 2O12 +1S 6 Negative charge 0 Total electrons18 -3(single bonds)-6 12

25 We got this Lewis dot structure No reason not to write instead Which would lead to Is there any reason for us to Presume one of these is correct And not the other? No Grammar: double-headed arrow is used to separate resonance structures

26 1.Bonding = sharing –electrons between repulsive + nuclei 2.Lewis Dot structures help us visualize sharing of electrons Octets Double and triple bonds Resonance structures and No Clean Socks 3.Formal Charge to help distinguish between alternatives 4.Violations of the Octet Rule 2 electrons >8 electrons 5.Using electrons to predict the SHAPE of the molecules VESPR Effect of unpaired electrons on the central atom on molecular shape Effect of Multiple bonds How to deal with “no central atom” 6.Bond polarity Effect of electronegativity difference between atoms in bond Effect of molecular shape How to symbolize bond polarity 7.Discrepancies between Electron Orbital Diagrams and VESPR Introducing orbital hybridization, s, sp, sp 2, sp 3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

27 Remember our Low Charge Density; Spectator Polyatomic Anions? No Clean Socks NO 3 - N5N5 3(O)18 Charge 1 Total24 -single bonds-6 Remaining18 Octets (6x3 O +2 for N)-20 Deficit of 1 electron pair-2 N5N5 3(O)18 Charge 1 Total24 -Bonds-8 16 Charge is distributed over All three of the resonance Forms = one big fat marshmallow

28 Resonance 1.The “real” molecule is none of the three nitrates we drew but something intermediate to the three. 2.Resonance can be “assumed” or “predicted” when there are equally plausible Lewis dot structures. 3.Resonance forms differ only in the distribution of electrons and not in the arrangement of atoms.

29 Write three resonance forms for SO 3 Valence electrons 4(6)24 Sulfur central atom Three single bonds to the sulfur-3(2) -6 Remaining electrons18 2 electrons to complete S octet-2 3(6) electrons to complete O octets = Deficit of two electrons = double bond-2 Valence electrons 4(6)24 Sulfur central atom Four single bonds to the sulfur-3(2) -8 Remaining electrons16

30 1.Bonding = sharing –electrons between repulsive + nuclei 2.Lewis Dot structures help us visualize sharing of electrons Octets Double and triple bonds Resonance structures and No Clean Socks 3.Formal Charge to help distinguish between alternatives 4.Violations of the Octet Rule 2 electrons >8 electrons 5.Using electrons to predict the SHAPE of the molecules VESPR Effect of unpaired electrons on the central atom on molecular shape Effect of Multiple bonds How to deal with “no central atom” 6.Bond polarity Effect of electronegativity difference between atoms in bond Effect of molecular shape How to symbolize bond polarity 7.Discrepancies between Electron Orbital Diagrams and VESPR Introducing orbital hybridization, s, sp, sp 2, sp 3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

31 Formal Charge helps determine the correct Lewis Dot Structure X=number of valence e- in the free atom (last number of group) Y = number of unshared e- owned by the atom in the Lewis structure Z = number of bonding e- shared by the atom in the Lewis structure The correct Lewis dot structure is generally the one in which a.The formal charges are as close to zero as possible b.Negative charge is located on the more electronegative atom

32 X=number of valence e- in the free atom (last number of group) Y = number of unshared e- owned by the atom in the Lewis structure Z = number of bonding e- shared by the atom in the Lewis structure Which is correct?

33 1.Bonding = sharing –electrons between repulsive + nuclei 2.Lewis Dot structures help us visualize sharing of electrons Octets Double and triple bonds Resonance structures and No Clean Socks 3.Formal Charge to help distinguish between alternatives 4.Violations of the Octet Rule 2 electrons >8 electrons 5.Using electrons to predict the SHAPE of the molecules VESPR Effect of unpaired electrons on the central atom on molecular shape Effect of Multiple bonds How to deal with “no central atom” 6.Bond polarity Effect of electronegativity difference between atoms in bond Effect of molecular shape How to symbolize bond polarity 7.Discrepancies between Electron Orbital Diagrams and VESPR Introducing orbital hybridization, s, sp, sp 2, sp 3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

34 The Octet Rule is a form of Rule #C2- Everybody wants to be like a Noble gas In same fashion: not everybody can share enough Electrons to make up a perfect octet These guys will have 1, 2, and 3 bonds only

35 The Octet Rule is a form of Rule #C2- Everybody wants to be like a Noble gas In same fashion: not everybody can share enough Electrons to make up a perfect octet This guy may end up “holding the bag” Having an unpaired electron Because he is Not really Strong enough To always Get the lion’s Share of the Electrons in A covalent bond, Particularly With oxygen The presence of these unpaired electrons on these gases Gives rise to the many atmospheric reactions involved In ozone destruction and formation of smog. Para = paramour = love = similar orientation Dia – diatribe = against = opposite orientation The unpaired electron means The compound is paramagnetic

36 A singlet electron is also called a “free radical” Aging hydroxide Hydroxy radical Cellular membrane damage

37 After hydroxyl radical cleavage and denaturing gel electrophoresis, the gel patterns differ. The unfolded RNA molecule is cleaved uniformly, giving rise to a homogeneous ladder of bands on the gel (left, bottom). The gel pattern for the folded RNA (right, bottom), in contrast, shows several region where strand cleavage is inhibited, corresponding to the sites of low solvent accessibility in the folded structure (right, top). TD Tullius, JA Greenbaum, Curr Opin Chem Biol 2005, 9:127–134 (Figure 1) Commentary from grove.ufl.edu/~dmorgan/Articles/DER/Journal%20Club.ppt

38 The Octet Rule is a form of Rule #C2- Everybody wants to be like a Noble gas Some guys can take on more electrons because they Make use of their d orbitals these guys Have d orbitals That allow them To have more Than 8 electrons 3d 4s 3p

39 Draw the Lewis structure of XeF 4 8+4(7)=36 electrons 4bonds = 8 electrons Remainder = 28 electrons Octets: 4(6) for F = 24 Remainder to Xe = 4

40 Draw the Lewis structure of XeF 2 8+2(7)=22 electrons 2bonds = 4 electrons Remainder = 18 electrons Octets: 2(6) for F = 12 Remainder to Xe = 6

41 “A” students work (without solutions manual) ~ 10 problems/night. Alanah Fitch Flanner Hall Office Hours W – F 2-3 pm

42 1.Bonding = sharing –electrons between repulsive + nuclei 2.Lewis Dot structures help us visualize sharing of electrons Octets Double and triple bonds Resonance structures and No Clean Socks 3.Formal Charge to help distinguish between alternatives 4.Violations of the Octet Rule 2 electrons >8 electrons 5.Using electrons to predict the SHAPE of the molecules VESPR Effect of unpaired electrons on the central atom on molecular shape Effect of Multiple bonds How to deal with “no central atom” 6.Bond polarity Effect of electronegativity difference between atoms in bond Effect of molecular shape How to symbolize bond polarity 7.Discrepancies between Electron Orbital Diagrams and VESPR Introducing orbital hybridization, s, sp, sp 2, sp 3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

43 Can now predict SHAPE of molecule determines: how two molecules orient themselves for reaction together. Can they dock? And actually do work together in three dimensional space?

44 V alence S hell E lectron Pair R epulsion Valence electron pairs surrounding an atom repel one another. Consequently, the orbitals containing those electron pairs are oriented to be as far apart as possible E el d Rule C1 = It’s all about charge

45 Geometries of AX 2 -AX 6 molecules AX Repulsion of valence shell electrons in Bonds and on F push F apart to a 180 o orientation X

46 Geometries of AX 2 -AX 6 molecules What orientation would put electrons as far apart as possible? 120 o degrees apart in a “circle”

47 Draw Lewis structures of a)CH 4 Valence shell electrons? Skeleton Carbon is first in formula Hydrogen is always terminal +C4 +4(H)4 Negative charge 0 Total electrons8 -4single bonds-8 remaining 0 Octets Carbon has its octet Hydrogen has its duet

48 How can four bonds be organized in 3-D space to be farthest apart? A X X X X

49 A pyramid is a space figure with a square base and 4 triangle-shaped sides. (5 “faces”) A tetrahedron is a space figure and 4 triangle shaped faces. Dictionary: a four-sided solid; a Triangular pyramid

50 A pyramid is a space figure with a square base and 4 triangle-shaped sides. (5 “faces”) A tetrahedron is a space figure and 4 triangle shaped faces. Dictionary: Square base and sloping Sides rising to an apex Dictionary: a four-sided solid; a Triangular pyramid

51 A pyramid is a space figure with a square base and 4 triangle-shaped sides. (5 “faces”) A tetrahedron is a space figure and 4 triangle shaped faces. Dictionary 1: Square base and sloping Sides rising to an apex Dictionary: a four-sided solid; a Triangular pyramid Dictionary 2: A solid figure with a polygon base. The surface, or lateral faces, are triangles having a common vertex. In a regular pyramid the base is a regular polygon and the lateral faces are congruent triangles SIGNIFICANT AMBIGUITY In nomenclature!!!!

52 90 o Assume Bond distance=1 X 1,1- X 2,1- Why tetrahedron and not this orientation? (Square planar) 180 o X 3,1- X 4,1- Since our example has all Q the same A Calculate the repulsion experience By atom X 1 with charge -1

53 90 o 1 X 1,1- X 2,1- Why tetrahedron and not this orientation? X 3,1- X 4,1- d 12 d 14 d13 A Calculation suggests that the electrostatic charge repulsion Energy is proportional to for a “square planar” orientation Of four identically charged atoms

54 A d 14 A X1X1 X4X4 60 o Square planar orientation= Compare to tetrahedron This means: atom 1 experiences less charge Repulsion from 2, 3, and 4 when tetrahedral

55 Geometries of AX 2 -AX 6 molecules

56 How can five bonds be arranged in space to be as far apart as possible? Triangular pyramid Triangular bipyramid A=Central atom X= atoms

57 Geometries of AX 2 -AX 6 molecules octahedron SF 6 How can these six guys best position themselves away from each other?

58

59 Some of the molecules we constructed using Lewis Dot structures had UNSHARED PAIRS of electrons on the CENTRAL ATOM What effect will this have on the geometry?. Unshared electron pairs orient themselves pretty much the same as single bonds. The observed molecular geometry (invisible electrons) is very different AX 2 E

60 This geometry, with respect to electron pairs and bonds, is triangular planar (three guys trying to get out of each others way) But one of the “terminal atoms” is missing so the molecular geometry differs from triangular planar Actual degrees observed is slightly less than 120 o because unshared electron pair expands Molecular Geometry is “bent” <120

61 Molecular Geometry Triangular planar AX 2 E Triangular planar 120 o Bent AX 3

62 AX 3 E This geometry, with respect to electron pairs and bonds, is tetrahedral (four guys trying to get out of each others way) But one of the “terminal atoms” is missing so the molecular geometry differs from tetrahedral Triangular Pyramid Explains why amines like ammonia can “steal” a proton From water – high charge density from the lone pair

63 107.2 o o F F F Effect of lone pairs on substituents (non-central atoms)

64 Effect of F is NOT by geometry of it’s lone pairs BUT By it’s electronegativity which pulls electrons along the bond, lowers Density of electrons in the bondings area Allows N lone pair to expand Compressing the angle between F atoms o o F F F

65 AX 2 E 2 This geometry, with respect to electron pairs and bonds, is tetrahedral (four guys trying to get out of each others way) But two of the “terminal atoms” are missing so the molecular geometry differs from tetrahedral This shape is “bent” What happens when we have Two Unbonded electron pairs on the Central atom, A?

66 AX 2 E 2 AX 2 E Both are bent, but the angle is different. Depends upon the number of valence shell electron pairs

67 Molecular Geometry AX 4 Tetrahedron AX 3 E Tetrahedron<109.5 o Triangular pyramid AX 2 E 2 Tetrahedron<109.5 o Bent

68 Triangular pyramid Triangular bipyramid 5 ELECTRON PAIRS AX 2 E 3 XeF o Triangular bipyramid Molecular Orientation Linear AX 3 E 2 Triangular bipyramid 90 o 180 o ClF 3 T-shape Triangular bipyramid Why put the E at equator?

69 E = non-bonding Electron pair S = Shared (bonded) electron pair Comparing where the non-bonded electron pair will go Variations: Axial versus Equatorial orientation equatorialaxial 120 o 90 o E E Electrostatic repulsion is sum of all near neighbor repulsions 1 S S S E E 2 S S S

70 If we put the E at the axial orientation they minimize E-E repulsion; increase E-S repulsion If we put E at the equatorial orientation E-E repulsion exists, but we decrease the E-S repulsion E E Minimizes impact Of E on S

71 Triangular pyramid Triangular bipyramid 5 ELECTRON PAIRS AX 2 E 3 XeF o Triangular bipyramid Molecular Orientation Linear AX 3 E 2 Triangular bipyramid 90 o 180 o ClF 3 T-shape

72 octahedron SF 6 AX 4 E 2 octahedron 90 o 180 o Molecular Orientation octahedron XeF 2 Square planar Square pyramidal

73 Square planar lets it slide into the DNA grove Why electron configuration is important: controls shape of molecule dictates 3D interaction of molecules Anticancer Drug

74 MULTIPLE Bonds Has no effect on geometry Multiple bond acts as a single bond Compare BF 3 and SO 3 Triangular Planar

75 Compare Molecular Geometries for BeF 2 and CO 2 1. We already did BeF 2

76 Compare Molecular Geometries for BeF 2 and CO 2 1. We already did BeF 2 Valence shell electrons for CO 2 ? Skeleton Carbon is first in formula= central atom +C4 +2(O)4 Negative charge 0 Total electrons16 -2single bonds-4 remaining12 e required for octets-16 deficit = multiple bonds-4 +C4 +2(O)4 Negative charge 0 Total electrons16 -4bonds-8 Remaining8 Octet for C = 0 Octet for each O = 4

77 Compare Molecular Geometries for BeF 2 and CO 2 Both are linear

78 Figure out geometry with NO CENTRAL ATOM Consider each carbon separately AX 3 Geometry around the carbon = Triangular Planar

79 Figure out geometry with NO CENTRAL ATOM Consider each carbon separately AX 2 Geometry around the carbon =Linear

80 Number of Electron Domains Electron-domain geometry Predicted Bond Angles TetrahedralTrigonal Tetrahedral Planar o 120 o o

81 “A” students work (without solutions manual) ~ 10 problems/night. Alanah Fitch Flanner Hall Office Hours W – F 2-3 pm

82 1.Bonding = sharing –electrons between repulsive + nuclei 2.Lewis Dot structures help us visualize sharing of electrons Octets Double and triple bonds Resonance structures and No Clean Socks 3.Formal Charge to help distinguish between alternatives 4.Violations of the Octet Rule 2 electrons >8 electrons 5.Using electrons to predict the SHAPE of the molecules VESPR Effect of unpaired electrons on the central atom on molecular shape Effect of Multiple bonds How to deal with “no central atom” 6.Bond polarity Effect of electronegativity difference between atoms in bond Effect of molecular shape How to symbolize bond polarity 7.Discrepancies between Electron Orbital Diagrams and VESPR Introducing orbital hybridization, s, sp, sp 2, sp 3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

83 Polarity – distribution of electrons in the bond 1.Depends upon the difference in electronegativity of bonded atoms 2.If two atoms in the bond are identical =nonpolar 3.Otherwise all bonds are polar Bond Polarity

84 Electronegativity is a measure of Positive Nuclear charge density experienced by Electron on another atom, scaled to a maxium of 4 Decreasing size H–HH–H H–CH–C H–FH–F  E.N. = =0 nonpolar  E.N. = = 0.3 slightly polar  E.N. = 4-2.2=1.8 strongly polar

85 Depends on 1.bond polarity 2.molecular shape a.diatomic molecules are linear molecule polar if atoms differ b.Polyatomic molecules can have polar bonds and still be non-polar H – ClCl – Cl polarnonpolar Molecular Polarity Polar molecules line up in an electric field

86 F  Be  F Arrow indicates the direction in which electrons are biased - the negative pole Bonds are polar  E.N. = 4-1.6=2.4 Vectors cancel each other Non-polar molecule Bonds are polar  E.N. = =1.3 Net charge direction Polar molecule CCl 4 Bonds are polar  E.N. = =0.7 Non polar molecule No net charge direction

87 Non-polarpolar Very polar bonds (Cl = 3.2; C=2.5) Non polar Very polar bonds (F=4; Be=1.6) Changing One atom Means Vectors Don’t cancel Changing Central atom so That a lone Electron pair Bends the molecule Means Vectors Don’t cancel

88 “A” students work (without solutions manual) ~ 10 problems/night. Alanah Fitch Flanner Hall Office Hours W – F 2-3 pm

89 1.Bonding = sharing –electrons between repulsive + nuclei 2.Lewis Dot structures help us visualize sharing of electrons Octets Double and triple bonds Resonance structures and No Clean Socks 3.Formal Charge to help distinguish between alternatives 4.Violations of the Octet Rule 2 electrons >8 electrons 5.Using electrons to predict the SHAPE of the molecules VESPR Effect of unpaired electrons on the central atom on molecular shape Effect of Multiple bonds How to deal with “no central atom” 6.Bond polarity Effect of electronegativity difference between atoms in bond Effect of molecular shape How to symbolize bond polarity 7.Discrepancies between Electron Orbital Diagrams and VESPR Introducing orbital hybridization, s, sp, sp 2, sp 3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

90 VSEPR model suggests that once Be bonds to F the orbitals are “equivalent” and therefore are equidistant from each other. But the electron orbital diagram suggests otherwise that Be has paired electrons and would not make bonds at all. VSEPR AX 2 Linear Repulsion of valence shell electrons pushes F apart to a 180 o orientation The “fly in the ointment”

91 Orbital diagram of F and Be in BeF 2 sp Shared electrons Orbital diagram of isolated atoms F and Be Orbital diagram of isolated F and hybridized Be atoms 1.Be has no unpaired electrons available for bonding 2.Be promotes 1 “2s” electron to a “2p” orbital 3.= “sp” hybridization 4.The new “sp” electrons engage in bonding with the unpaired electrons on F

92 Formation of Hybrid Atomic Orbitals S + p orbitals = 2 “sp” hybridized orbitals = The number of hybridized orbitals formed = number of atomic orbitals mixed Energies of hybridized orbitals intermediate to the atomic orbitals mixed

93 sp What does the bond made from the atomic “sp” and “p” atomic orbitals look like? + Sigma bond Single bond

94 Orbital diagram of isolated atoms F and B sp 2 1.Boron needs to have three of its electrons shared with the three F electrons to create the three single bonds. 2.B mixes 2 “s2” and 1 “2p” electrons using 1 s and 2 “p” orbitals 3.3 sp 2 atomic orbitals with 1 e each formed 4.Electrons in these orbitals are shared with unpaired electrons on F to create single bonds (blue) VSEPR AX 3 Triangular Planar

95 What does an sp2 orbital look like?

96 1.C needs four energetically equivalent bonds 2.C mixes 1 “2s” and 3 “2p” atomic orbitals 3.4 sp 3 atomic orbitals formed 4.sp 3 orbitals used to create bonds sp 3 VSEPR AX 4 Tetrahedral Orbital diagram of isolated atoms H and C

97 sp 3 orbital on C + s orbital on H Form a sigma bond

98

99 So far we have considered VSEPR AX 2 Linear VSEPR AX 3 Triangular Planar VSEPR AX 4 Tetrahedral sp sp 2 sp 3 What about the guys with expanded octets? 2 orbitals 3 orbitals 4 orbitals

100 these guys Have d orbitals That allow them To have more Than 8 electrons 3d 4s 3p VSEPR AX 5 Triangular bipyramid

101 VSEPR AX 5 triangular bipyramid sp 3 d

102 Nothing new here – same as we got with VSEPR AX 2 AX 3 AX 4 AX 5 AX 6 VSEPR Note to myself from 2006 – F does not hybridize – has one unpaired electron Forming one bond

103 Atomic Orbitals Hybridize to Create Equivalent Orbitals on One atom for Bonding with another atom Hybridized Orbitals From two diff atoms combine to create Bonding orbitals 0kJ -kJ 180 o

104 Whenever we have a “single” bond we can assume that it has the sigma shape, resulting from hybridization between atomic orbitals We need to simply create additional bonds within the shape predicted by VSEPR Sigma bond Single bond For double and triple bonds, we do not need to create more equivalent bonds which can be moved as far apart as predicted by Valence Shell Electron Pair Repulsion. Pi bond Double bond around single bond Double and Triple Bonds

105 AX 3 Geometry around the carbon = Triangular Planar What does ethylene, C 2 H 4 look Like?

106

107 AX 2 Geometry around the carbon =Linear What does acetylene, C 2 H 2, look like?

108 Example problem: What is the hybridization of nitrogen in a) NO 3 -, b) NH 2 Cl, c) N 2, d) and N 2 O? 1.Make the Lewis dot structure 2.Count sigma bonds on central atom 3.Count the unbonded electron pairs on central atom 4.Sum the 2 (n) 5.Number of hybridized orbitals =n Cmpdσ Bonds NO Cmpdσ Bonds Unbonded electron pairs on N NO Cmpdσ Bonds Unbonded electron pairs on N n NO Cmpdσ Bonds Unbonded electron pairs on N nHybridization on N NO sp 2 Cmpdσ Bonds Unbonded electron pairs on N nHybridization on N NO sp 2 NH 2 Cl314sp 3 Cmpdσ Bonds Unbonded electron pairs on N nHybridization on N NO sp 2 NH 2 Cl314sp 3 N2N2 112sp Cmpdσ Bonds Unbonded electron pairs on N nHybridization on N NO sp 2 NH 2 Cl314sp 3 N2N2 112sp N 2 O a112sp N 2 O b Cmpdσ Bonds Unbonded electron pairs on N nHybridization on N NO sp 2 NH 2 Cl314sp 3 N2N2 112sp N 2 O a112sp N 2 O b202sp

109 RESONANCE

110 Benzene, C 6 H 6, is a very common compound Usually don’t show Hydrogens or Carbons Sigma bonds P orbitals Pi (double) bonds Benzene resonance Charge delocalization

111 Delocalization/Resonance Structures Nitrate is a is not charge dense Poor Nitrate 1 charge Large radius Low Charge density BIG KEY POINT!!!!!! Nitrate: no clean socks

112 1.Bonding = sharing –electrons between repulsive + nuclei 2.Lewis Dot structures help us visualize sharing of electrons Octets Double and triple bonds Resonance structures and No Clean Socks 3.Formal Charge to help distinguish between alternatives 4.Violations of the Octet Rule 2 electrons >8 electrons 5.Using electrons to predict the SHAPE of the molecules VESPR Effect of unpaired electrons on the central atom on molecular shape Effect of Multiple bonds How to deal with “no central atom” 6.Bond polarity Effect of electronegativity difference between atoms in bond Effect of molecular shape How to symbolize bond polarity 7.Discrepancies between Electron Orbital Diagrams and VESPR Introducing orbital hybridization, s, sp, sp 2, sp 3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

113 “A” students work (without solutions manual) ~ 10 problems/night. Alanah Fitch Flanner Hall Office Hours W – F 2-3 pm


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