FITCH Rules “A” students work (without solutions manual) G1: Suzuki is Success ~ 10 problems/night. G2. Slow me down G3. Scientific Knowledge is Referential General Alanah Fitch G4. Watch out for Red Herrings Flanner Hall 402 G5. Chemists are Lazy 508-3119 [email protected] C1. It’s all about charge

Office Hours W – F 2-3 pm C2. Everybody wants to “be like Mike” (grp.18) C3. Size Matters

Chemistry C4. Still Waters Run Deep C5. Alpha Dogs eat first

Properties and Measurements 1. Bonding = sharing – between repulsive + nuclei Property Unit Reference State 2. Lewis Dot structures help us visualize sharing of electrons Size m size of earth Octets Volume cm3 m Double and triple bonds Weight gram mass of 1 cm3 water at specified Temp structures and No Clean Socks (and Pressure) 3. Formal Charge to help distinguish between alternatives Temperature oC, K boiling, freezing of water (specified 4. Violations of the Pressure) 2 electrons 1.66053873x10-24gamu (mass of 1C-12 )/12 >8 electrons quantity atomic mass of an element in grams 5. Using electrons to predict the SHAPE of the molecules Pressure atm, mm Hg earth’s atmosphere at sea level VESPR Energy: Thermal BTU 1 lb water 1 oF Effect of unpaired electrons on the central atom on molecular shape calorie 1 g water 1 oC Effect of Multiple bonds Kinetic J 2kg mass moving at 1m/s How to deal with “no central atom” Energy, of electrons energy of in a vacuum 6. Bond polarity F Effect of electronegativity difference between in bond Effect of molecular shape How to symbolize bond polarity 7. Discrepancies between Electron Orbital Diagrams and VESPR Introducing orbital hybridization, s, sp, sp2, sp3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

1 Covalent Bonding – getting to a by Covalent bonding sharing electrons 1. Bring two elements close together Patterns in abundance suggest 2. When very close the positive nuclei repel each other a. periodicity b. preferred electronic configuration of elements Leading to the Rule: “Everybody wants to be “Like Mike” a. : Groups 16 and 17 gain electrons; Groups 1 and 2 lose b. Other atoms share electrons to have eight electrons = COVALENT BONDING

Repulsion of two atoms with their Proton core + +

Repulsion of two hydrogen atoms with their proton core

Repulsion is high close where e Protons see each other “A” students work Atoms which are far apart (without solutions manual) + Repulsion is low where Electrons shield nucleus, and where Do not even see each other 10 problems/night. There is no energy, repulsive ~ e Electrons can be stabilized by both Positive charges Or attractive between the two +

e e Alanah Fitch Flanner Hall 402 Repulsive+ energy + 508-3119 e e [email protected] + + Office Hours W – F 2-3 pm

Electrons are the jelly and peanut butter between the slices of bread Attractive energy (protons)

2 1. Bonding = sharing –electrons between repulsive + nuclei 2. Lewis Dot structures help us visualize sharing of electrons Lewis dot structure (electron dot Octets Structure or diagram) are diagrams that Double and triple bonds Show the bonding between atoms of Resonance structures and No Clean Socks A molecule based on shared “” shell 3. Formal Charge to help distinguish between alternatives (outer shell) electrons and shows the 4. Violations of the Octet Rule Presence of any “” of electrons 2 electrons That may exist in the covalently bonded >8 electrons Molecule. 5. Using electrons to predict the SHAPE of the molecules VESPR Gilbert Newton Lewis Effect of unpaired electrons on the central atom on molecular shape Covalent 1875-1946; Caltech Effect of Multiple bonds Physical Chemist How to deal with “no central atom” 6. Bond polarity Effect of electronegativity difference between atoms in bond Valence = outermost shell electrons of an atom Effect of molecular shape shared How to symbolize bond polarity 7. Discrepancies between Electron Orbital Diagrams and VESPR Latin: valere – to be strong Introducing orbital hybridization, s, sp, sp2, sp3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

When the two hydrogen atoms are together, the electron configuration When the two hydrogen atoms are together, the electron configuration Looks like? Looks like?

2 2 He 1sHe= [] He 1sHe= []

When a hydrogen atom and a fluorine atom share electrons, the When a hydrogen atom and a fluorine atom share electrons, the Electron configuration on fluorine looks like? Electron configuration on fluorine looks like?

122ssp22 6= [] Ne 122ssp22 6= [] Ne

2 The “inner” shell 1 s electrons do not Only the “outer-most” or valence shell Show in this diagram electrons Show in this Lewis Dot Structure How many valence electrons?: = last number in group

3 When a hydrogen atom and an atom share valence electrons plus an When the two hydrogen atoms are together, the electron configuration Extra electron, the electron configuration on hydrogen and oxygen look like? Looks like? •• Valence electrons on oxygen? • Valence electrons on hydrogen? H • O • He 2 •• 1sHe= [] − •• ⎡ •• ⎤ • • • 2 • ++→OH• • ⎢• OH• ⎥ 1sHe= [] When a hydrogen atom and a fluorine atom share electrons, the •• ⎣⎢ •• ⎦⎥ 22 6 Electron configuration on fluorine looks like? 122ssp= [] Ne

Invoking Rule 2: Chemists are Lazy 22 6 the diagram above is too tedious to write out all the time 122ssp= [] Ne make shared electrons (bond) a line

Lewis dot structure for hydroxide

The shared pair of Electrons = covalent The single shaired between the two bonded atoms bond Is called a single bond The unshared pairs of electrons are “regions of high Charge density” It is drawn as a line.

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

•• •• •• •• 2 2 • • • 1sHe= [] • • • 1sHe= [] OHH• ++→•• HOH• • OHH• ++→•• HOH• • •• •• •• •• 22 6 22 6 122ssp= [] Ne 122ssp= [] Ne

Two shared electron pairs = Two single bonds When three hydrogen atoms and a atom share valence electrons, the electron configuration on hydrogen and nitrogen look like?

4 When four hydrogen atoms and two atoms share valence electrons, the When two hydrogen atoms and an oxygen atom share valence electrons, the electron configuration on hydrogen and carbon look like? electron configuration on hydrogen and oxygen look like? • Valence shell of carbon? •• •• 2 •C• • • • 1sHe= [] OHH• ++→•• HOH• • • •• •• 22 6 H H 122ssp= [] Ne • • •• •• • • • • •CH••••+++ H CHH•+ • +→ • HC• • • C• H • • 2 When three hydrogen atoms and a nitrogen atom share valence electrons, the 1sHe= [] Two electron pairs shared is a electron configuration on hydrogen and nitrogen look like? Double bond 22 6 Valence shell of nitrogen? •• 122ssp= [] Ne • N• • Three pairs of shared electrons = three single bonds •• •• • • 2 • NHHH••••+++→ HN• • H 1sHe= [] • •• 22 6 H 122ssp= [] Ne

Rules for Writing Lewis Dot Structures When two hydrogen atoms and two carbon atoms share valence electrons, the 1. Count the number of valence electrons (lat number of group) of all atoms electron configuration on hydrogen and carbon look like? 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 NH4 , S • ••••+ • • • • • • in SO , and C in CCl ). CH++ CH → HCCH• • • • • 2 4 • • b. The terminal atoms are usually H, O. c. are always terminal atoms. 1sHe2 = [] 3. Determine the number of valence electrons still available for distribution after Three electron pairs shared is a subtracting two electrons for each single bond. 22 6 Triple bond 4. Determine the number of electrons required to complete the octet 122ssp= [] Ne 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 H • Electron. Poor H. Halogens have lots of electrons but really do not like to share. •• Greedy halogens • • • F All they want is one more to make up the Mike configuration ••

5 Draw Lewis structures of Draw Lewis structures of Skeleton a) Hypochlorite Valence shell electrons? a) Hypochlorite ion b) Methyl alcohol b) Methyl alcohol, CH3OH Carbon is first in formula O 6 c) N2 c) N2 Hydrogen is always terminal d) SO2 +Cl 7 d) SO2 +Negative charge 1 H Total electrons 14 Valence shell electrons? Hypochlorite? -1Single bond -2 12 HC−−−OH Hypo – smallest number of -2(6 electrons for O,Cl) 12 O 6 remaining 0 +C 4 H OCl- +4(H) 4 Negative charge 0 Total electrons 14 Octets -5single bonds -10 Carbon has its octet remaining 4 Hydrogen has its duet -octet for oxygen -4 Oxygen requires 4 more electrons remaining 0 H − •• ⎡ •• ••⎤ HC−−−OH Skeleton • • OCl− ⎢• OCl− ⎥• •• •• •• ⎣⎢ ⎦⎥ H

Skeleton, First atom in formula is central Draw Lewis structures of Draw Lewis structures of a) Hypochlorite ion a) Hypochlorite ion OSO−− b) Methyl alcohol, CH3OH b) Methyl alcohol, CH3OH c) N c) N 2 2 Octets d) SO2 Skeleton NN− d) SO2 We are short 2 electrons for the octet, Octets Each nitrogen requires 6 more Valence shell electrons? Valence shell electrons? The only way to get extra ones is to Share two more electrons = double bond. 2N 10 2O 12 Negative charge 0 +1S 6 OS−= O Total electrons 10 Negative charge 0 -1single bond -2 We are short 4 electrons for the octet, Total electrons 18 Place the remaining 12 electrons to fill Remaining 8 The only way to get extra ones is to -2(single bonds) -4 Octet completion -12 Share four more electrons = triple Remaining electrons 14 octets Difference -4 Bond. Octet for S -4 2(Octet for each O) -12 •• •• •• Deficit? -2 • 2N 10 NN≡ • OS−= O Negative charge 0 •• •• Total electrons 10 Place the remaining 4 electrons equally 2O 12 -3 single bonds -6 On the two equal +1S 6 Remaining 4 Negative charge 0 Total electrons 18 • • • NN≡ • -3(single bonds) -6 12

6 We got this Lewis dot structure 1. Bonding = sharing –electrons between repulsive + nuclei 2. Lewis Dot structures help us visualize sharing of electrons OS−= O No reason not to write instead OSO=− Octets Double and triple bonds •• •• •• Which would lead to •• •• •• Resonance structures and No Clean Socks • OS−= O • • OSO=−• 3. Formal Charge to help distinguish between alternatives •• •• •• •• 4. Violations of the Octet Rule Is there any reason for us to 2 electrons Presume one of these is correct >8 electrons And not the other? 5. Using electrons to predict the SHAPE of the molecules VESPR No 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 • OS−= O O=− SO• •• •• •• •• Effect of molecular shape How to symbolize bond polarity 7. Discrepancies between Electron Orbital Diagrams and VESPR Grammar: double-headed arrow is used to separate resonance 2 3 structures Introducing orbital hybridization, s, sp, sp , sp Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

Remember our Low Charge Density; Spectator Polyatomic Anions? Resonance No Clean Socks 1. The “real” molecule is non of the three nitrates we drew but something - intermediate to the three. NO3 − 2. Resonance can be “assumed” or “predicted” when there are equally plausible − •• N5⎡ •• ⎤ ⎡ ⎤ Lewis dot structures. • • • O• 3(O) 18 ⎢ • O• ⎥ ⎢ • • ⎥ ⎢ ⎥ ⎢ ⎥ 3. Resonance forms differ only in the distribution of electrons and not in the Charge 1 •• •• Total 24 ⎢ •• ••⎥ ⎢ ⎥ arrangement of atoms. • ⎢ON=−O• ⎥ -single bonds -6 ⎢• ON−=O⎥ • ⎢ •• ••⎥ ⎢•• •• ⎥ Remaining 18 ⎢ ⎥ Octets (6x3 O +2 for N) -20 ⎢ ⎥ ⎢ ⎥ ⎢ − ⎥ Deficit of 1 electron pair -2 ⎡ ⎢ ⎤ ⎥ ⎣⎢ ⎦⎥ • • ⎣ ⎦ ⎢ • O• ⎥ ⎢ ⎥ O Charge is distributed over ⎢ •• •• ⎥ • • All three of the resonance ⎢• ON−−O• ⎥ •• •• ON−−O Forms = one big fat marshmallow ⎢ ⎥ ⎢ ⎥ ⎢ ⎥ ⎣⎢ ⎦⎥

7 1. Bonding = sharing –electrons between repulsive + nuclei Write three resonance forms for SO 3 2. Lewis Dot structures help us visualize sharing of electrons Octets Valence electrons 4(6) 24 O Double and triple bonds Sulfur central atom Resonance structures and No Clean Socks Three single bonds to the sulfur -3(2) -6 3. Formal Charge to help distinguish between alternatives Remaining electrons 18 −− 4. Violations of the Octet Rule 2 electrons to complete S octet -2 OSO 3(6) electrons to complete O octets = 18 -18 2 electrons Deficit of two electrons = double bond -2 >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 •• •• • O• • O• • O• How to deal with “no central atom” • • • • • • 6. Bond polarity •• •• •• •• •• •• Effect of electronegativity difference between atoms in bond • OS−=O • • • Effect of molecular shape • OS=−O• • OS−−O• •• •• •• •• •• •• How to symbolize bond polarity 7. Discrepancies between Electron Orbital Diagrams and VESPR Introducing orbital hybridization, s, sp, sp2, sp3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

X=number of valence e- in the free atom (last number of group) Formal Charge helps determine the correct Lewis Dot Structure Y = number of unshared e- owned by the atom in the ⎡ Z ⎤ formal charg e≡=−+ Cf X⎢ Y ⎥ ⎡ Z ⎤ ⎣ 2⎦ formal charg e≡=−+ Cf X⎢ Y ⎥ Z = number of bonding e- shared by the atom in the Lewis structure ⎣ 2⎦ X=number of valence e- in the free atom (last number of group) ≡ = ⎡ 6 ⎤ formalformalformal ch ch chargargarg e e e≡=−+≡ C C CfOfO =66262− [ 1 =+ H fO ⎣⎢ 2 ⎦⎥ Y = number of unshared e- owned by the atom in the Lewis structure •• HC− −−O H formal charg e≡=−≡ C = 424 [ ⎡ 6 ⎤ formal charg e≡=−+ C fC 42 1 =− Z = number of bonding e- shared by the atom in the Lewis structure •• fC ⎣⎢ 2 ⎦⎥ The correct Lewis dot structure is generally the one in which H

a. The formal charges are as close to zero as possible Which is correct? b. Negative charge is located on the more electronegative atom H •• ⎡ 4 ⎤ formal charg e≡=−+≡ C fO = 646 − [ 0 = HC−−−OH fO ⎣⎢ ⎦⎥ •• 2 H 8 ≡ = − ⎡ ⎤ formal charg e≡=−+≡=− C fC 404 [ 0 = fC ⎣⎢2 ⎦⎥

8 1. Bonding = sharing –electrons between repulsive + nuclei 2. Lewis Dot structures help us visualize sharing of electrons The Octet Rule is a form of Rule #C2- Everybody wants to be like a Noble gas Octets In same fashion: not everybody can share enough Double and triple bonds Electrons to make up a perfect octet Resonance structures and No Clean Socks 3. Formal Charge to help distinguish between alternatives These guys will have 1, 4. Violations of the Octet Rule 2, H • 2 electrons and 3 bonds only >8 electrons • 5. Using electrons to predict the SHAPE of the molecules • B• VESPR • • Effect of unpaired electrons on the central atom on molecular shape ••Be Effect of Multiple bonds How to deal with “no central atom” •• 6. Bond polarity • • Effect of electronegativity difference between atoms in bond •• •• F • • • • Effect of molecular shape • FBeF−−• How to symbolize bond polarity •• •• •• •• • • 7. Discrepancies between Electron Orbital Diagrams and VESPR • FB−−F• Introducing orbital hybridization, s, sp, sp2, sp3 •• •• Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

The Octet Rule is a form of Rule #C2- Everybody wants to be like a Noble gas 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 Some guys can take on more electrons because they Electrons to make up a perfect octet Make use of their d orbitals This guy may end up “holding the bag” Having an

• Because he is • • Not really • N • Strong enough To always • • Get the lion’s Share of the The presence of these unpaired electrons on these gases Electrons in these guys Gives rise to the many atmospheric reactions involved A , 3p Have d orbitals In ozone destruction and formation of smog. Particularly With oxygen That allow them 4s •• 3d To have more • •• • •• •• • •• Than 8 electrons • ON−= O • NO= • ONO=−• •• •• •• •• •• •• Para = paramour = love = similar orientation Dia – diatribe = against = opposite orientation

9 •• Draw the Lewis structure of XeF2 Draw the Lewis structure of XeF4 F • • • F • 8+4(7) =36 electrons 4bonds = 8 electrons Remainder = 28 electrons FXe−−F •• •• Octets: 4(6) for F = 24 • • • FXe− − F• Remainder to Xe = 4 •• •• 8+2(7) =22 electrons F 2bonds = 4 electrons • F • Remainder = 18 electrons • • • • •• Octets: 2(6) for F = 12 FXeF−− • F • Remainder to Xe = 6 • • •• •• • • • FXeF−−• •• •• •• •• • • • FXe− − F• •• •• •• •• •• • • • FXeF−−• • F • •• •• •• • • ••

1. Bonding = sharing –electrons between repulsive + nuclei “A” students work 2. Lewis Dot structures help us visualize sharing of electrons Octets (without solutions manual) Double and triple bonds ~ 10 problems/night. Resonance structures and No Clean Socks 3. Formal Charge to help distinguish between alternatives 4. Violations of the Octet Rule 2 electrons Alanah Fitch >8 electrons Flanner Hall 402 5. Using electrons to predict the SHAPE of the molecules 508-3119 [email protected] VESPR Effect of unpaired electrons on the central atom on molecular shape Office Hours W – F 2-3 pm 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, sp2, sp3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

10 kQ12 Q E = Valence Rule C1 = It’s all about charge el d Can now predict SHAPE of molecule 9 − Shell = 899. xJm 10 k Electron 2 determines: how two molecules orient themselves ()Coulomb Pair for reaction together. Eel Repulsion

Can they dock? And actually do work d together in three dimensional space?

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

Geometries of AX2-AX6 molecules Geometries of AX2-AX6 molecules

•• •• X A X • • • FBeF−−• •• •• •• •• • FBeF−−• •• • • • • •• •• • F• Repulsion of valence What •• •• shell electrons pushes Fe orientation would • • o put electrons as far • FB−−F• apart to a 180 orientation •• •• apart as possible?

120o degrees apart in a “circle”

11 H Draw Lewis structures of How can four a) CH4 bonds be organized in HC−−H Valence shell electrons? 3-D space to H +C 4 be farthest apart? +4(H) 4 Negative charge 0 Total electrons 8 -4single bonds -8 remaining 0

Skeleton Carbon is first in formula Hydrogen is always terminal H Octets Carbon has its octet HC−−H Hydrogen has its duet H

A pyramid is a space figure with a square base and 4 triangle-shaped sides. A pyramid is a space figure with a square base and 4 triangle-shaped sides. (5 “faces”) (5 “faces”)

Dictionary: Square base and sloping Sides rising to an apex 3 3 2 4 2 4 5 1 1 5

A tetrahedron is a space figure and 4 triangle shaped faces. A tetrahedron is a space figure and 4 triangle shaped faces. Dictionary: a four-sided solid; a Dictionary: a four-sided solid; a Triangular pyramid Triangular pyramid

3 3 4 2 2 4 1 1

12 Why tetrahedron and not this orientation? (Square planar) kQ Q A pyramid is a space figure with a square base and 4 triangle-shaped sides. 12 Calculate the Eel = (5 “faces”) d repulsion Dictionary 1: Square base and sloping 1,1- 9 − experience = 899. xJm 10 Sides rising to an apex k By atom 1 90 o 2 ()Coulomb 3 Dictionary 2: A solid figure with a Bond distance=1 2 4 polygon base. The surface, or 180 o 5 lateral faces, are triangles having a 4,1- A 2,1- 1 common vertex. In a regular pyramid the base is a regular polygon and the lateral faces are ⎡QQ12 QQ13 QQ14 ⎤ congruent triangles Ekel,,1 total =++⎢ ⎥ ⎣d12 d13 d14 ⎦ A tetrahedron is a space figure and 4 triangle shaped faces. 3,1- Dictionary: a four-sided solid; a Triangular pyramid 111⎡ Since⎤ our example has all Q the same EkQQel,,11 total=++() i ⎢ ⎥ 3 ddd12⎣ 13 14 ⎦ 2 4 SIGNIFICANT AMBIGUITY In nomenclature!!!! 1 ⎡ 111 ⎤ Eel,,1 total ∝++⎢ ⎥ ⎣ddd12 13 14 ⎦

Square planar orientation= 1.914 opp Why tetrahedron and not this orientation? Compare to tetrahedron sinθ = ⎡ 111 ⎤ hyp Eel,,1 total ∝++⎢ ⎥ ⎣ddd12 13 14 ⎦ 111 1,1- ⎡ ⎤ ()14 Eel,,1 total ∝++⎢ ⎥ ⎣ ⎦ o ddd12 13 14 = d 90 o 2 14 d12 sin()60 1 =1 1 22 ()14 dAA=+=+=()21112() 26014()= o () A 12 sin 4,1- d13 2,1- 2 A11= 1732. () 14 ()14 60o ⎡ 1 1 1 ⎤ Eel,,1 total ∝++⎢ 1914.⎥ = 2 A 3,1- ⎣ 2 2 2 ⎦ 4 A41= 1

d14 ⎡ 1 1 1 ⎤ A ∝++= Calculation suggests that the electrostatic charge repulsion Eel,,1 total ⎢ 1732. ⎥ 4 ⎣1732...1732 1732 ⎦ 2 Energy is proportional to 1.914 for a “square planar” orientation 3 Of four identically charged atoms This means: atom 1 experiences less charge Repulsion from 2, 3, and 4 when tetrahedral

13 Geometries of AX2-AX6 molecules Geometries of AX2-AX6 molecules

Triangular pyramidbipyramid

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

Geometries of AX2-AX6 molecules

octahedron SF6

How can these six guys best position themselves away from each other?

14 Some of the molecules we constructed using Lewis Dot structures had •• •• •• •• •• •• • • • OS−= O OSO=−• AX E •• •• •• •• 2 UNSHARED PAIRS of electrons on the CENTRAL ATOM This geometry, with respect to electron pairs and bonds, What effect will this have on the geometry?. is triangular planar (three guys trying to get out of each others way)

But one of the “terminal atoms” is missing so the •• •• •• •• •• •• AX2E • • molecular geometry differs from triangular planar • OS−= O O=− SO• <120 •• •• •• •• Unshared electron pairs orient •• •• themselves pretty much the same • • • OHH• ++→•• HOH• • as single bonds. Actual degrees observed •• •• is slightly less than 120o because •• •• The observed molecular unshared electron pair expands • • geometry (invisible • NHHH••••+++→ HN• H• • •• electrons) is very Molecular Geometry is “bent” H different

•• • • HN• •H •• AX3E Effect of lone pairs on substituents H (non-central atoms)

This geometry, with respect to electron pairs and bonds, H is tetrahedral (four guys trying to get out of each others way) •• •• •• •• • • • • • • • F • N •F • But one of the “terminal atoms” is missing so the HC−−H HNH •• •••• •• • ••• molecular geometry differs from tetrahedral •F • H H ••

Triangular 107.2o o Pyramid 102.3

F F Explains why amines like ammonia can “steal” a proton F From water – high charge density from the lone pair

15 What happens when we have Two Effect of F is NOT by geometry of it’s lone pairs BUT Unbonded electron pairs on the By it’s electronegativity which pulls electrons along the bond, lowers Density of electrons in the bondings area Central atom, A? Allows N lone pair to expand Compressing the angle •• •• •• •• • • • • • • F N F HNH • •• • ••• •• • • ••• • • vs •F • AX E AX E H •• 2 2 2

107.2o

102.3o

F F F

•• Both are bent, • • HOH• • AX2E2 but the angle is different. •• Depends upon the number This geometry, with respect to electron pairs and bonds, of valence shell electron pairs is tetrahedral (four guys trying to get out of each others way)

•• •• •• But two of the “terminal atoms” are missing so the AX E • 2 • OS−= O molecular geometry differs from tetrahedral •• ••

•• • • HOH• • AX2E2 ••

This shape is “bent”

16 5 ELECTRON PAIRS

Molecular Orientation Triangular bipyramid

Triangular pyramidbipyramid

•• •• •• • • o • FXeF−−• Triangular 180 •• •• •• Linear AX2E3 bipyramid XeF2

X+E=3456 AX E Triangular o AX3 AX4 AX5 AX6 3 2 90 T-shape bipyramid 180o ClF3 AX2EAX2E2 AX2E3 AX5E Why put the E at equator? AX3E2 AX4E2

Comparing where the non-bonded electron pair will go Variations: Axial versus Equatorial orientation If we put the E at the axial orientation they axial 90o S E minimize E-E repulsion; increase E-S repulsion 120o S S equatorial E S If we put E at the equatorial orientation E S E-E repulsion exists, but we decrease the E-S repulsion S E 1 2

E E E = non-bonding electron pair S = bonded electron pair Minimizes impact Electrostatic repulsion is sum of all near neighbor repulsions Of E on S

EEEESESSSTotal,1 =−()equatorial 242() +equatorial −+ ( − )axial+ ( − ) axial

ESSESTotal,2 =−36() equatorial ( +− )axial

17 5 ELECTRON PAIRS Geometries of molecules with expanded octets (Fig. 7.8)

Molecular 5 ELECTRON PAIRS Orientation

Triangular pyramidbipyramid

•• •• •• • • o • FXeF−−• Triangular 180 •• •• •• Linear AX2E3 bipyramid XeF2

AX2E3

AX E Triangular o 3 2 90 T-shape bipyramid 180o ClF3

Molecular Orientation

octahedron SF6

octahedron

octahedron •• • F• AX E • • 4 2 o 90 •• •• • • • FXe− − F • 180o •• ••

• F • • • •• Square X+E=3456 XeF2 planar

AX3 AX4 AX5 AX6 AX2EAX2E2 AX2E3 AX5E AX3E2 AX4E2

18 Why electron configuration is important: controls shape of molecule Geometries of molecules with expanded octets (Fig. 7.8) dictates 3D interaction of molecules Anticancer Drug

Square planar lets it slide into the DNA grove

•• •• Compare Molecular Geometries • • for BeF and CO FBeF−− MULTIPLE Bonds 2 2 • • •• •• 1. We already did BeF2 Has no effect on geometry Multiple bond acts as a single bond

Compare BF3 and SO3 Triangular Planar •• •• •• • • • O• • • • • • F• • • • O• • O• •• •• •• •• •• •• •• •• • OS−=O • • • • • • OS=−O• • OS−−O• • FB−−F• •• •• •• •• •• •• •• ••

19 Compare Molecular Geometries Compare Molecular Geometries

for BeF2 and CO2 for BeF2 and CO2

1. We already did BeF2

Skeleton Carbon is first in formula= central atom Valence shell electrons for CO2? +C 4 •• •• •• •• +2(O) 4 • • OCO− − • FBeF−−• == Negative charge 0 •• •• OCO Total electrons 16 •• •• -2single bonds -4 OCO== remaining 12 Both are linear e required for octets -16 •• •• deficit = multiple bonds -4 OCO== +C 4 •• •• +2(O) 4 Negative charge 0 Total electrons 16 Octet for C = 0 -4bonds -8 Octet for each O = 4 Remaining 8

Figure out geometry with NO CENTRAL ATOM Figure out geometry with NO CENTRAL ATOM

H H •• •• • • • • • • • • • HC• • • • CH• HC• • •C •H

Consider each carbon separately Consider each carbon separately

AX3 AX2

Geometry around Geometry around the carbon =Linear the carbon = Triangular Planar

20 1. Bonding = sharing –electrons between repulsive + nuclei “A” students work 2. Lewis Dot structures help us visualize sharing of electrons Octets (without solutions manual) Double and triple bonds ~ 10 problems/night. Resonance structures and No Clean Socks 3. Formal Charge to help distinguish between alternatives 4. Violations of the Octet Rule 2 electrons Alanah Fitch >8 electrons Flanner Hall 402 5. Using electrons to predict the SHAPE of the molecules 508-3119 [email protected] VESPR Effect of unpaired electrons on the central atom on molecular shape Office Hours W – F 2-3 pm 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, sp2, sp3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

Decreasing size Electronegativity is a measure of Positive Nuclear charge density experienced by Bond Polarity Electron on another atom , scaled to a maxium of 4 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

H–H ∆E.N. = 2.2-2.2=0 nonpolar ∆E.N. = 2.5-2.2= 0.3 slightly polar H–C ∆E.N. = 4-2.2=1.8 strongly polar

H–F

21 Molecular Polarity Arrow indicates the direction in which electrons CH4 Depends on are biased - the negative pole 1. bond polarity F ← Be → F O H 2. molecular shape H H Bonds are polar C a. diatomic molecules are linear Bonds are polar H H ∆E.N. = 4-1.6=2.4 molecule polar if atoms differ ∆E.N. = 3.5-2.2=1.3 H –H–Cl Cl–Cl – polar nonpolar Vectors cancel each other Bonds are weak polar Non-polar Net charge direction ∆E.N. = 2.5-2.2=0.3 molecule Polar molecules line Polar molecule up in an electric field No net charge direction b. Polyatomic molecules can have polar Non polar bonds and still be non-polar molecule

Very polar bonds (F=4; Be=1.6) Non-polar polar “A” students work Changing (without solutions manual) Central atom so 10 problems/night That a lone ~ . Electron pair Bends the molecule Means Vectors Alanah Fitch Don’t cancel Flanner Hall 402 508-3119 [email protected] Changing One atom Office Hours W – F 2-3 pm Means Vectors Don’t cancel

polar

Non polar Very polar bonds (Cl = 3.2; C=2.5)

22 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 The “fly in the ointment” 4. Violations of the Octet Rule 2 electrons >8 electrons 5. Using electrons to predict the SHAPE of the molecules •• •• Repulsion of valence VESPR • FBeF−−• VSEPR shell electrons pushes F Effect of unpaired electrons on the central atom on molecular shape • • •• •• AX2 apart to a 180o orientation Effect of Multiple bonds Linear How to deal with “no central atom” 6. Bond polarity Effect of electronegativity difference between atoms in bond VSEPR model suggests that once Be bonds to F the orbitals Effect of molecular shape are “equivalent” and therefore are equidistant from each other. How to symbolize bond polarity 7. Discrepancies between Electron Orbital Diagrams and VESPR But the electron orbital diagram suggests otherwise that Be has 2 3 Introducing orbital hybridization, s, sp, sp , sp paired electrons and would not make bonds at all. Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

•• •• 12s s 2 p • • F ()↓↑ ()↓↑ ()()()↓↑ ↓↑ ↓ • FBeF−−• •• •• ↓↑ ↓↑ Be () () () () () Formation of Hybrid Atomic Orbitals 1. Be has no unpaired ↓↑ ↓↑ ↓↑ ↓↑ ↓ electrons available for F () () ()()() bonding Orbital diagram of isolated atoms F and Be S + p orbitals = 2 “sp” hybridized orbitals

1s 22sp F ()↓↑ ()↓↑ ()()()↓↑ ↓↑ ↑ 2. Be promotes 1 “2s” electron to a “2p” = Be ()↓↑ ()↓ ↓ () () () orbital 3. = “sp” hybridization F ()↓↑ () ↓↑ ()()() ↓↑ ↓↑ ↑

Orbital diagram of isolated F and hybridized Be atoms The number of hybridized orbitals formed = 1s 22sp 4. The new “sp” electrons F ()↓↑ ()↓↑ ()()()↓↑ ↓↑ ↑ number of atomic orbitals mixed engage in bonding with sp Be ()↓↑ ()↓ ↓ () () () the unpaired electrons Shared Energies of hybridized orbitals intermediate to on F electrons F ()↓↑ () ↓↑ ()()() ↓↑ ↓↑ ↑ the atomic orbitals mixed

Orbital diagram of F and Be in BeF2

23 What does the bond made from the atomic “sp” and •• 12s s 2 p • • F ()↓↑ ()↓↑ ()()()↓↑ ↓↑ ↓ “p” atomic orbitals look like? • F• ↓↑ ↓↑ ↑ 22sp •• •• B () () ()() () 1s • • •• •• ()↓↑ ()()()↓↑ ↓↑ ↑ • FB−−F• ↓↑ ↓↑ ↓↑ ↓↑ ↓ F ()↓↑ •• •• • • sp F () () ()()() FBeF−− ↓↑ ↓↑ ↓↑ ↓↑ ↓ • • Be ()↓↑ ()↓ ↓ () () () •• •• VSEPR AX3 F () () ()()() Triangular Planar F ()↓↑ () ↓↑ ()()() ↓↑ ↓↑ ↑ Orbital diagram of isolated atoms F and B 1. 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 12s s 2 p 3. 3 sp2 atomic orbitals with 1 e each F ()↓↑ ()↓↑ ()()()↓↑ ↓↑ ↓ formed sp2 ↓↑ ↑ ↑↑ B () () () ()() Sigma bond 4. Electrons in these orbitals are Single bond shared with unpaired electrons on ↓↑ ↓↑ ↓↑ ↓↑ ↓ F to create single bonds (blue) F () () ()()() σ ↓↑ ↓↑ ↓↑ ↓↑ ↓ F () () ()()()

H 1s 2s 2 p VSEPR AX H ()↓ What does an sp2 orbital look like? 4 ↓↑ ↓↑ ↑↑ Tetrahedral C () () () ()() −− HCH ↓ H () ↓

H H ()↓

H () 1. C needs four energetically equivalent bonds Orbital diagram of isolated atoms H and C 2. C mixes 1 “2s” and 3 “2p” atomic orbitals

3. 4 sp3 atomic orbitals formed 1s 2s 2 p H ()↓ sp3 3 ↓↑ ↑ ↑↑↑ 4. sp orbitals used to create bondsC () () () () () ↓ H () ↓

H ()↓

H ()

24 H •• • • HC−−H HN• •H •• H H

sp3 orbital on C + s orbital on H Form a sigma bond

So far we have considered •• •• VSEPR sp • • VSEPR • FBeF−−• AX5 •• •• 2 orbitals AX2 Triangular bipyramid Linear

•• • • • F• VSEPR 2 sp •• •• AX3 • FB−−F• 3 orbitals • • Triangular Planar •• •• these guys 3p Have d orbitals H That allow them 4s 3d To have more 3 VSEPR Than 8 electrons sp HC−−H AX4 4 orbitals Tetrahedral H What about the guys with expanded octets?

25 VSEPR

AX5 triangular VSEPR bipyramid AX2 AX3 AX4 AX 3p 3d 5 3s AX6

Cl[]10 Ne ()↓↑ ()()()↓↑ ↓↑ ↓ () () () sp3d () () []↑ ↑↑↑ ↑ () () () PNe10 () ()() ()() () Nothing new here – same as we got with VSEPR ↓↑ ↓↑ ↓↑ ↓

Cl[]10 Ne () ()()() () () ()() () ↓↑ ↓↑ ↓↑ ↓ [] () () ()() () Cl Ne () ()()() 10 ↓↑ ↓↑ ↓↑ ↓

↓↑ ↓↑ ↓↑ ↓ Cl[]10 Ne () ()()() () () ()() () [] () () ()() ()

Cl10 Ne () ()()()

Double and Triple Bonds What does ethylene, C H look Like? Whenever we have a “single” bond we can assume that 2 4 it has the sigma shape, resulting from hybridization between atomic orbitals H H •• •• • • • • Sigma bond HC• • •C •H Single bond σ

For double and triple bonds, we do not need to create more equivalent bonds which can be moved as far apart AX3 as predicted by Valence Shell Electron Pair Repulsion.

We need to simply create additional bonds within the Geometry around shape predicted by VSEPR the carbon = Triangular Planar Pi bond Double bond around single bond π

26 What does acetylene, C2H2, look like?

• • • • • HC• • • • CH•

AX2

Geometry around the carbon =Linear

Benzene, C6H6, is a very common compound H H Usually don’t show H H or

HH RESONANCE Benzene resonance

P orbitals

Sigma bonds Charge Pi (double) delocalization bonds

27 BIG KEY POINT!!!!!! 1. Bonding = sharing –electrons between repulsive + nuclei 2. Lewis Dot structures help us visualize sharing of electrons Delocalization/Resonance Structures Octets Double and triple bonds Nitrate: no clean socks 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 Nitrate is a is not charge dense Poor Nitrate 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 1 charge Low 7. Discrepancies between Electron Orbital Diagrams and VESPR kQ12 Q Charge Introducing orbital hybridization, s, sp, sp2, sp3 Eel = + Large radius density Using orbital hybridization to visualize resonance or smearing out rr12 of electron density in the molecule.

“A” students work (without solutions manual) ~ 10 problems/night.

Alanah Fitch Flanner Hall 402 508-3119 [email protected]

Office Hours W – F 2-3 pm

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