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Notes​ ​by​ ​Aaron​ ​Hui
Lecture​ ​3
...
17​ ​-​ ​Quantum​ ​Theory
● Maxwell​ ​(1873)​ ​proposed​ ​that​ ​visible​ ​light​ ​consists​ ​of​ ​electromagnetic​ ​waves
○ All​ ​EM​ ​radiation​ ​c=λv
● Small​ ​λ,​ ​high​ ​v,​ ​and​ ​most​ ​energy
...
0​ ​x​ ​10​4​​ ​Hz​ ​(cycles/s​ ​=​ ​s-1​
​ )
...
00 * 10 8 m/s = λ (6
...
0 * 10 3 m *

10 9 nm
m

= 5
...
63​ ​x​ ​10​-34​​ ​J*s
● Einstein​ ​in​ ​1905
○ Photoelectric​ ​effect:​ ​eject​ ​e-​ ​from​ ​the​ ​surface​ ​of​ ​certain​ ​metals​ ​exposed​ ​to
light​ ​of​ ​a​ ​minimum​ ​frequency​ ​(threshold​ ​frequency)
○ Photon​ ​-​ ​“particle”​ ​of​ ​light
■

E p = hv

■ E​ ​=​ ​energy​ ​of​ ​1​ ​photon
○ Light​ ​has​ ​both​ ​wave​ ​nature​ ​and​ ​particle​ ​nature
...

Calculate​ ​the​ ​energy​ ​(in​ ​joules)​ ​associated​ ​with​ ​each​ ​photon​ ​if​ ​the​ ​wavelength​ ​of
the​ ​X​ ​rays​ ​is​ ​0
...
​ ​ ​What​ ​is​ ​the​ ​energy​ ​for​ ​a​ ​mol​ ​of​ ​photons?

Ep =

hc
λ

λ = 0
...
63*10

= 1
...
00*10 m/s)
1
...
29*10 −15 J
photon

23

*

1
...
022*10 photons
mol

= 7
...

● Line​ ​Emission​ ​Spectrum​ ​-​ ​light​ ​emission​ ​only​ ​at​ ​specific​ ​wavelengths
(“fingerprint”)
● Bohr’s​ ​Model​ ​of​ ​the​ ​Atom​ ​(1913)
○ Lines​ ​in​ ​spectra​ ​correspond​ ​to​ ​energy​ ​of​ ​e-​​ ​ ​(quantized)
...

● Quantum​ ​Mechanics
○ Describes​ ​how​ ​e-​ ​act​ ​as​ ​both​ ​particles​ ​and​ ​waves
○ Electrons​ ​do​ ​not​ ​travel​ ​in​ ​circular​ ​orbits
○ Orbital​ ​-​ ​a​ ​probability​ ​map​ ​(region​ ​of​ ​space)​ ​of​ ​where​ ​an​ ​e-​ ​is​ ​most​ ​likely​ ​to
be​ ​found
...

● Principal​ ​Energy​ ​Shells

○ Quantum​ ​numbers​ ​-​ ​describe​ ​distribution​ ​of​ ​electrons​ ​in​ ​an​ ​atom
■ Principle​ ​Shell​ ​-​ ​orbitals​ ​of​ ​approximately​ ​the​ ​same​ ​size​ ​and​ ​energy
■ 𝛹​ ​=​ ​function​ ​(​n​,​ ​l​,​ ​m​l​,​ ​m​s​)
● Represented​ ​by​ ​the​ ​principal​ ​quantum​ ​number,​ ​n​ ​(n=1,2,3…)
● Higher​ ​“n”,​ ​higher​ ​energy,​ ​farther​ ​from​ ​nucleus
● #​ ​orbitals​ ​=​ ​n2​
● n​ ​tells​ ​us​ ​the​ ​distance​ ​of​ ​e-​ ​from​ ​the​ ​nucleus
○ Subshells​ ​-​ ​part​ ​of​ ​the​ ​principal​ ​shell,​ ​orbitals​ ​have​ ​the​ ​same​ ​energy​ ​and
similar​ ​shape
■ 𝛹​ ​=​ ​fn(n,​ ​l​,​ ​m​l​,​ ​m​s​)
Angular​ ​momentum​ ​quantum​ ​number​ ​l
■ Same​ ​values​ ​as​ ​n​ ​and​ ​l
■ #​ ​of​ ​subshells​ ​=​ ​n
■ l​ ​=​ ​0,1,2,3,
...
,​ ​0,
...
​ ​ ​So​ ​3​ ​orbitals
○ How​ ​many​ ​electrons​ ​can​ ​be​ ​placed​ ​in​ ​the​ ​3rd​ ​subshell?
■ n​ ​=​ ​3,​ ​l​ ​=​ ​2,​ ​m​l​​ ​=​ ​-2,​ ​-1,​ ​0,​ ​1,​ ​2​ ​so​ ​5​ ​orbitals,​ ​and​ ​each​ ​orbital
contains​ ​2​ ​e-​ ​each,​ ​so​ ​=​ ​10​ ​electrons​ ​in​ ​3rd​ ​subshell
...

■ n​ ​=​ ​3,​ ​l​ ​=​ ​1,​ ​m​l​​ ​=​ ​-1,​ ​0,​ ​1​ ​,​ ​m​s​​ ​=​ ​+½,​ ​-½
○ What​ ​is​ ​the​ ​total​ ​number​ ​of​ ​orbitals​ ​associated​ ​with​ ​the​ ​principle​ ​quantum
number​ ​n​ ​=​ ​4?
■ N​2​​ ​=​ ​42​​ ​ ​=​ ​16​ ​orbitals
● Electron​ ​configuration​ ​-​ ​distribution​ ​of​ ​electrons​ ​among​ ​the​ ​atomic​ ​orbitals
...

● Only​ ​2e-​ ​with​ ​opposite​ ​spins​ ​can​ ​occupy​ ​with​ ​the​ ​same
atomic​ ​orbital
● Each​ ​subshell​ ​has​ ​a​ ​max​ ​#​ ​of​ ​e○ s​ ​=​ ​2
○ p​ ​=​ ​6

○ d​ ​=​ ​10
○ f​ ​=14
○ Aufbau​ ​Principle​​ ​-​ ​electrons​ ​occupy​ ​the​ ​next​ ​available​ ​orbital​ ​with​ ​the
lowest​ ​energy
...

○ Using​ ​the​ ​periodic​ ​table​ ​to​ ​determine​ ​electron​ ​configuration
■ Row​ ​=​ ​energy​ ​shell
■ Length​ ​of​ ​row​ ​depends​ ​on​ ​types​ ​of​ ​subshells
● Paramagnetic​ ​-​ ​unpaired​ ​electrons​ ​so​ ​attracted​ ​to​ ​a​ ​magnet
● Diamagnetic​ ​-​ ​all​ ​electrons​ ​paired​ ​-​ ​slightly​ ​repelled​ ​by​ ​magnet
● Valence​ ​electrons​ ​-​ ​electrons​ ​in​ ​outermost​ ​principal​ ​shell​ ​(largest​ ​n)
○ Si:​ ​1s2​ ​2s2​ ​2p6​ ​3s2​ ​3​ ​p2
■ Highest​ ​shell​ ​3:​ ​2​ ​e-​ ​from​ ​s​ ​and​ ​2​ ​e-​ ​from​ ​p,​ ​so​ ​2+2​ ​=​ ​4​ ​V
...

○ Valence​ ​shell​ ​configuration​ ​-​ ​electron​ ​config​ ​of​ ​valence​ ​electrons
● Tellurium​ ​(Te):​ ​[Kr]​ ​5s2​ ​4d10​ ​5p4
○ Valence​ ​shell​ ​config:​ ​5s2​ ​5p4
○ Total​ ​V
...
E
...

● O:​ ​1s2​ ​2s2​ ​2p4
O​2-​ ​:​ ​1s2​ ​2s2​ ​2p6
Al:​ ​[Ne]​ ​3s2​ ​2p1
Al​3+​​ ​:​ ​[Ne]
Al3+​ ​and​ ​N3-​ ​are​ ​isoelectric​ ​with​ ​Ne
...
30
...

● Periodic​ ​law​ ​-​ ​if​ ​arrange​ ​elements​ ​by​ ​Z,​ ​their​ ​chemical​ ​and​ ​physical​ ​properties
vary​ ​periodically
...

○ Effective​ ​nuclear​ ​charge​ ​(Z​eff​)​ ​-​ ​“positive​ ​charge”​ ​felt​ ​by​ ​an​ ​electron
...

● Atomic​ ​radius​ ​(atomic​ ​size)​ ​-​ ​one-half​ ​the​ ​distance​ ​between​ ​the​ ​2​ ​nuclei​ ​in​ ​2
adjacent​ ​atoms
...

○ Larger​ ​Z​eff​​ ​→​ ​stronger​ ​hold​ ​of​ ​nucleus​ ​on​ ​e-​ ​→​ ​smaller​ ​the​ ​atomic​ ​radius
○ ↓​ ​across​ ​a​ ​period
■ ↑​ ​#​ ​of​ ​protons​ ​but​ ​small​ ​shell​ ​of​ ​e● Put​ ​in​ ​order​ ​of​ ​decreasing​ ​atomic​ ​radius
○ P,​ ​Si,​ ​N
■ Si​ ​>​ ​P​ ​>​ ​N
○ C,​ ​Li,​ ​Be
■ Li​ ​>​ ​Be​ ​>​ ​C
● Ionic​ ​Radius
○ Size​ ​of​ ​an​ ​ion
...

■ Often​ ​loses​ ​its​ ​outer​ ​shell
■ Nuclear​ ​charge​ ​remains​ ​the​ ​same​ ​but​ ​fewer​ ​electrons​ ​(atom
shrinks)
...

■ Nuclear​ ​charge​ ​remains​ ​the​ ​same​ ​but​ ​more​ ​e-​ ​(weaker​ ​so​ ​expands)
○ Summary:​ ​cation​ ​<​ ​atom​ ​<​ ​anion
● Comparing​ ​Ions​ ​to​ ​each​ ​other
○ Ions​ ​in​ ​the​ ​same​ ​group​ ​(family)
■ Increases​ ​down​ ​a​ ​group​ ​(if​ ​same​ ​charge)
● Size​ ​of​ ​shell​ ​increases​ ​as​ ​you​ ​go​ ​down​ ​a​ ​group
...

I 1 + X (g) → X

+

I 2 + X (g) → X

2+

I 3 + X (g) → X

3+

+e−

I 1 − 1 st IE

(g)

+e−

I 2 − 2 nd IE

(g)

+e−

I 3 − 3 nd IE

(g)

● IE​ ​is​ ​switched​ ​for​ ​columns​ ​2​ ​&​ ​3,​ ​also​ ​for​ ​5​ ​&​ ​6​ ​due​ ​to​ ​extra​ ​stability​ ​of​ ​half
filled​ ​p​ ​subshell​ ​and​ ​full​ ​s​ ​subshell
...

○ Look​ ​at​ ​energy​ ​configs

● Ionization​ ​energy​ ​and​ ​metallic​ ​character
○ Metallic​ ​character​ ​based​ ​is​ ​IE
○ More​ ​metallic​ ​character,​ ​more​ ​the​ ​element​ ​exhibits​ ​physical​ ​and​ ​chemical
properties​ ​of​ ​metals
...

○ Low​ ​EA​ ​=​ ​hard​ ​to​ ​gain​ ​e○ High​ ​EA​ ​=​ ​easy​ ​to​ ​gain​ ​e● Electronegativity​ ​(EN)​ ​-​ ​the​ ​ability​ ​to​ ​attract​ ​electrons​ ​in​ ​a​ ​covalent​ ​bond
○ Closer​ ​to​ ​Fluorine,​ ​more​ ​EN
...
Absorbance​ ​(A)​ ​(how​ ​much​ ​light​ ​the​ ​sample​ ​absorbs)​ ​(higher​ ​the​ ​#​ ​the​ ​more​ ​light
it​ ​absorbs)
2
...
T =

P
PO

* 100

Direct​ ​linear​ ​relationship​ ​between​ ​A​ ​and​ ​known​ ​concentration:​ ​use​ ​A​ ​to​ ​find
concentration
A = − logT
If​ ​T​ ​=​ ​
...
​ ​ ​Its​ ​a​ ​logarithmic​ ​curve,​ ​so​ ​use​ ​linear​ ​A
...

What​ ​is​ ​the​ ​absorbance​ ​of​ ​a​ ​solution​ ​that​ ​has​ ​40%​ ​transmittance​ ​at​ ​500nm?
A = − log(
...
40

As​ ​we​ ​change​ ​λ:
● Intensity​ ​of​ ​light​ ​changes
● Sensitivity​ ​of​ ​phototube​ ​changes
● Amount​ ​of​ ​light​ ​absorbed​ ​by​ ​cell​ ​and​ ​H​2​O​ ​also​ ​varies
We​ ​will​ ​“blank”​ ​at​ ​each​ ​λ​:
● Set​ ​absorbance​ ​of​ ​H​2​O​ ​+​ ​cell​ ​=​ ​0
● Blank​ ​for​ ​each​ ​λ
...

Red​ ​dye​ ​and​ ​yellow​ ​dye
Goal:​ ​Find​ ​λ​ ​where​ ​1​ ​dye​ ​absorbs​ ​well​ ​but​ ​no​ ​absorbance​ ​from​ ​other​ ​dye
...
Graph​ ​plotting​ ​both​ ​red​ ​and​ ​yellow​ ​dye​ ​on​ ​same​ ​graph
2
...

3
...
0​ ​mM​ ​and​ ​h2o
...
0​ ​mM
v2​ ​=​ ​10​ ​mL
Need​ ​to​ ​make​ ​10​ ​mL​ ​of​ ​each​ ​solution
Exp
...
4
...

● Lewis​ ​Dot​ ​Symbols
○ Element​ ​symbol​ ​+​ ​valence​ ​electrons​ ​as​ ​dots
■ Keep​ ​dots​ ​as​ ​far​ ​apart​ ​as​ ​possible​ ​(e
...
​ ​singly​ ​at​ ​first)
● The​ ​Ionic​ ​Bond
○ Electrostatic​ ​force​ ​between​ ​ions​ ​in​ ​an​ ​ionic​ ​compound
■ Involves​ ​transfer​ ​of​ ​electrons
● (QUIZ)​ ​Electrostatic​ ​(Lattice)​ ​Energy
○ Lattice​ ​energy​ ​(E)​ ​-​ ​energy​ ​required​ ​to​ ​separate​ ​1​ ​mol​ ​of​ ​solid​ ​ionic
compound​ ​into​ ​gaseous​ ​ions
■ Measure​ ​of​ ​stability​ ​of​ ​an​ ​ionic​ ​solid
○ Coulomb’s​ ​Law
■

E=k

Q+Q−
r

■ Q​+​​ ​-​ ​charge​ ​on​ ​cation
■ Q​-​​ ​-​ ​charge​ ​on​ ​anion
■ r​ ​-​ ​distance​ ​between​ ​ions
■ Lattice​ ​energy​ ​(E)​ ​increases​ ​as​ ​Q​ ​increases​ ​(absolute​ ​value)​ ​and/or
as​ ​r​ ​decreases
...
​ ​(Based​ ​on​ ​Hess’s​ ​Law)
○

Example:
Calculate​ ​the​ ​lattice​ ​energy​ ​for​ ​LiF​ ​if​ ​the​ ​heat​ ​of​ ​sublimation​ ​of​ ​Li​ ​is​ ​155
...
1​ ​kJ/mol
...
2​ ​(heat​ ​of​ ​sublimation)

½​ ​F​2​(g)​ ​→​ ​2F(g)

ΔH°​ ​(kJ)​ ​=​ ​½(155)​ ​(bond​ ​energy/enthalpy)

Li(g)​ ​→​ ​Li​+​​ ​(g)​ ​+​ ​e-

ΔH°​ ​(kJ)​ ​=​ ​ ​520​ ​(IE)

e-​ ​+​ ​F(g)​ ​→​ ​F-​​ ​ ​(g)

ΔH°​ ​(kJ)​ ​=​ ​-328​ ​(EA)

Li​+​​ ​(g)​ ​+​ ​F-​​ ​ ​(g)​ ​→​ ​LiF(s)

ΔH°​ ​(kJ)​ ​=​ ​ ​X​ ​(opposite​ ​of​ ​lattice​ ​energy)

Li(s)​ ​+​ ​½​ ​F​2​(g)​ ​→​ ​LiF(s)

ΔH​f​(LiF​ ​(s))​ ​=​ ​-594
...
2​ ​+​ ​½​ ​(155)​ ​+​ ​520​ ​-​ ​320​ ​+​ ​X)​ ​kJ​ ​=​ ​-594
...
5​ ​kJ/mol,​ ​the​ ​second
ionization​ ​energy​ ​of​ ​Ca​ ​is​ ​1145​ ​kJ/mol,​ ​the​ ​second​ ​ionization​ ​energy​ ​of​ ​Ca
is​ ​1145​ ​kJ/mol,​ ​the​ ​electron​ ​affinity​ ​of​ ​Cl​ ​is​ ​-349​ ​kJ,​ ​and​ ​the​ ​heat​ ​of
formation​ ​of​ ​CaCl​2​​ ​is​ ​-795​ ​kJ/mol
...
5

Ca​+​​ ​(g)​ ​→​ ​Ca​2+​​ ​(g)​ ​+​ ​e-

ΔH°​ ​(kJ)​ ​=​ ​1145

2e-​ ​+​ ​2Cl(g)​ ​→​ ​2Cl​-​​ ​(g)

ΔH°​ ​(kJ)​ ​=​ ​2(-349)

Ca​2+​(g)​ ​+​ ​2Cl​-​(g)​ ​→​ ​CaCl​2​(s)

ΔH°​ ​(kJ)​ ​=​ ​-2195

Ca(s)​ ​+​ ​Cl​2​(g)​ ​→​ ​CaCl​2​(s)

ΔH​f​(CaCl​2​​ ​(s))​ ​=​ ​-795

(x+242+589
...
0

Nonpolar​ ​covalent

C​ ​and​ ​N​ ​(𝛿+​ ​ ​𝛿-)

0
...
9

Polar​​ ​covalent​ ​(less)

Li​ ​and​ ​F​ ​(+​ ​-)

3
...
6
...

●

Writing​ ​Lewis​ ​Structures​ ​-​ ​Covalent​ ​Compounds

1
...


Count​ ​total​ ​number​ ​of​ ​valence​ ​ea
...
Subtract​ ​1​ ​for​ ​each​ ​+​ ​charge

3
...


Complete​ ​an​ ​octet​ ​for​ ​all​ ​atoms​​ ​except​ ​H
a
...


If​ ​central​ ​atom​ ​doesn’t​ ​have​ ​an​ ​octet​ ​form,​ ​double​ ​and​ ​triple​ ​bonds
● Two​ ​possible​ ​skeletal​ ​structures​ ​for​ ​formaldehyde​ ​(CH​2​O)

○

H-:C=:O-H

○

H↘
​ ​ ​ ​ ​ ​C=:O:
H↗

● The​ ​correct​ ​Lewis​ ​structure​ ​will​ ​minimize​ ​formal​ ​charge​ ​(FC)
● FC​ ​=​ ​Valence​ ​e-​ ​(VE)​ ​-​ ​#​ ​of​ ​nonbonding​ ​e-​ ​(NBE)​ ​-​ ​½​ ​#​ ​of​ ​electrons
shared​ ​in​ ​covalent​ ​bonds​ ​(BE)
● Second​ ​structure​ ​is​ ​the​ ​better​ ​structure​ ​because​ ​it​ ​has​ ​no​ ​FC
● Carbon​ ​hates​​ ​lone​ ​pairs
...

● Resonance​ ​structures​ ​-​ ​two​ ​or​ ​more​ ​Lewis​ ​structures​ ​that​ ​are​ ​“averaged”​ ​to​ ​more
accurately​ ​represent​ ​the​ ​real​ ​molecule
...
​ ​ ​It​ ​does​ ​NOT​ ​shuttle​ ​back​ ​and​ ​forth​ ​between​ ​forms
...
11
...


Use​ ​bond​ ​energies

b
...
7 − 0 − 52
...
0 kJ/mol ​ ​(EXACT)
● Valence​ ​shell​ ​electron​ ​pair​ ​repulsion​ ​(VESPR)​ ​model:

○

Model​ ​to​ ​predict​ ​the​ ​geometry​ ​of​ ​molecules

○

Electrons​ ​repel​ ​one​ ​another​ ​so​ ​place​ ​as​ ​far​ ​apart​ ​as​ ​possible

○

2​ ​Rules:
■ Double​ ​and​ ​triple​ ​bonds​ ​are​ ​treated​ ​like​ ​single​ ​bonds
■ If​ ​molecule​ ​has​ ​resonance​ ​structure,​ ​apply​ ​theory​ ​to​ ​any​ ​one
● Predicting​ ​Molecular​ ​Geometry

1
...


Determine​ ​#​ ​of​ ​electron​ ​groups​ ​(both​ ​bonded​ ​atoms​ ​and​ ​lone​ ​pairs)​ ​surrounding

the​ ​central​​ ​atom
3
...


For​ ​bond​ ​angles,​ ​remember​ ​that​ ​lone​ ​pairs​ ​are​ ​more​ ​repulsive​ ​than​ ​bonding​ ​pairs
● Electron​ ​geometry​ ​-​ ​arrangement​ ​of​ ​electron​ ​groups
● Molecular​ ​geometry​ ​-​ ​shape​ ​(just​ ​atoms​ ​not​ ​lone​ ​pairs)
● When​ ​drawing​ ​shape,​ ​must​ ​show​ ​lone​ ​pairs​ ​on​ ​central​ ​atom​ ​+​ ​multiple​ ​bonds
Electron
Groups

#​ ​Lone
Pairs

Electron
Geometry

Molecular
Geometry

Drawing
Shaper

Bond
Angles

2

0

Linear

Linear

Cl​ ​-​ ​Be​ ​-​ ​Cl

180

3

0

Trigonal

Trigonal

O

120

planar

Planar

||
C
H↗​ ​ ​ ​ ​ ​ ​ ​ ​↖H

Trigonal
planar

Bent


...
5

4

1

Tetrahedral

Trigonal
Pyramidal


...
5

4

2

Tetrahedral

bent


...
5

<109
...


<90

6

2

Octahedral

Square
planar

​ ​ ​ ​ ​ ​ ​ ​ ​
...


90

● Adding​ ​lone​ ​pairs​ ​affect​ ​the​ ​molecular​ ​geometry​​ ​(shape​ ​arrangement​ ​of​ ​atoms)
but​ ​not​ ​the​ ​electron​ ​geometry​ ​(arrangement​ ​of​ ​electrons)
○

Lone​ ​pairs​ ​are​ ​more​ ​repulsive​ ​than​ ​atoms
● Number​ ​of​ ​molecular​ ​geometries​ ​equals​ ​Electron​ ​groups​ ​-​ ​1
...
​ ​Cont’d

Part​ ​2​ ​&​ ​3
● Pick​ ​λ
○

Maximize​ ​absorbance

○

No​ ​abs​ ​from​ ​other​ ​dye
● 22​ ​data​ ​points

Lecture​ ​4
...
17
● Geometry​ ​of​ ​Molecules​ ​with​ ​Multiple​ ​Central​ ​Atoms
○ Difficult​ ​to​ ​determine​ ​an​ ​overall​ ​geometry​ ​of​ ​molecule
○ Describe​ ​geometry​ ​around​ ​each​ ​central​ ​atom
○ CH​3​CH​4​COOH

■
■ Tetrahedral​ ​(first​ ​group​ ​around​ ​C),​ ​Tetrahedral​ ​(second​ ​group
around​ ​C),​ ​Trigonal​ ​planar​ ​(third​ ​group​ ​around​ ​C),​ ​Bent​ ​(O-H)
○ CH​3​CH​2​OH

■
○ CH​3​NHCH​3

■
● Predicting​ ​Molecular​ ​Geometry

○ What​ ​are​ ​the​ ​molecular​ ​geometries​ ​and​ ​bond​ ​angles​ ​of​ ​Al​I​3​​ ​and​ ​SCl​4

■

(trigonal​ ​planar)

(seesaw)

● Electronegativity​ ​and​ ​Polarity
○ Electronegativity​ ​-​ ​ability​ ​of​ ​an​ ​element​ ​to​ ​attract​ ​electrons​ ​within​ ​a
covalent​ ​bond
○ Polar​ ​Covalent​ ​Bond​ ​-​ ​covalent​ ​bond​ ​in​ ​which​ ​bonding​ ​electrons​ ​spend
more​ ​time​ ​near​ ​one​ ​atom​ ​than​ ​the​ ​other
...
0

Nonpolar​ ​covalent

(𝛿+)​ ​C​ ​&​ ​N​ ​(𝛿-)

0
...
0

Polar​ ​Covalent​ ​(more)

(+)​ ​Li​ ​&​ ​F​ ​(-)

3
...
Valence​ ​bond​ ​(VB)​ ​theory
2
...

○ Valence​ ​bond​ ​theory​ ​states:​ ​a​ ​stable​ ​molecule​ ​forms​ ​when​ ​the​ ​potential
energy​ ​of​ ​the​ ​system​ ​decreases​ ​to​ ​a​ ​minimum
...
Mix​ ​at​ ​least​ ​2​ ​nonequivalent​ ​atomic​ ​orbits​ ​(e
...
​ ​s​ ​and​ ​p)
2
...
#​ ​hybrid​ ​orbitals​ ​=​ ​#​ ​pure​ ​atomic​ ​orbitals
4
...

■ VESPR​ ​Geometry​ ​(electron​ ​geo):​ ​trigonal​ ​planar
● Formation​ ​of​ ​sp​ ​Hybrid​ ​Orbitals
○ BeF​2
■ 2​ ​electron​ ​groups​ ​around​ ​Be,​ ​so​ ​2​ ​hybrid​ ​orbitals
■ So​ ​Be​ ​forms​ ​sp​ ​orbitals,​ ​or​ ​two​ ​2sp​ ​orbitals​ ​in​ ​this​ ​case
...
Draw​ ​the​ ​Lewis​ ​structure
2
...
​ ​ ​What​ ​is​ ​the​ ​hybridization​ ​of​ ​the​ ​central
(underlined)​ ​atom?​ ​Describe​ ​which​ ​orbitals​ ​overlap​ ​during​ ​the
hybridization​ ​process:
■ P​F​3
● 4​ ​electron​ ​groups​ ​so​ ​four​ ​3sp3​ ​orbits​ ​around​ ​P
● But​ ​only​ ​three​ ​3sp3​ ​orbitals​ ​of​ ​P​ ​overlap​ ​with​ ​the​ ​2p​ ​of​ ​F​ ​(1
orbital​ ​is​ ​a​ ​lone​ ​pair)
■ P​Br​5
● 5​ ​electron​ ​groups​ ​so​ ​five​ ​3sp​3​d
● All​ ​five​ ​3sp​3​d​ ​orbitals​ ​of​ ​P​ ​overlap​ ​with​ ​the​ ​4p​ ​of​ ​Br
■ Be​H​2
● 2​ ​electron​ ​groups​ ​so​ ​two​ ​2sp​ ​hybridization
● All​ ​two​ ​2sp​ ​orbitals​ ​of​ ​Be​ ​overlap​ ​with​ ​1s​ ​orbital​ ​of​ ​H
Lecture​ ​4
...
17
● Hybridization​ ​in​ ​Molecules​ ​with​ ​Double​ ​and​ ​Triple​ ​Bonds
○ sp​2​​ ​hybridization​ ​of​ ​a​ ​Carbon​ ​Atom​ ​(double​ ​bond)
■ Ground​ ​state:​ ​2s​ ​↑↓​​ ​2p​ ​↑​ ​↑​ ​_
■ Promotion​ ​of​ ​electron:​ ​2s​ ​↑​ ​2p​ ​↑​ ​↑​ ​↑

■ sp​2​ ​hybridized​ ​orbitals:​ ​sp​2​​ ​↑​ ​↑​ ​↑​ ​2p​z​​ ​↑
● Sigma​ ​bond​ ​(σ)​ ​-​ ​orbitals​ ​overlap​ ​end-to-end,​ ​electron​ ​density​ ​between​ ​the​ ​2
atoms
● Pi​ ​bond​ ​(𝜋)​ ​-​ ​orbitals​ ​overlap​ ​sideways,​ ​electron​ ​density​ ​above​ ​and​ ​below​ ​plan​ ​of
nuclei
...

● Remember​ ​only​ ​central​ ​atom​ ​hybridizes
...

○ Antibonding​ ​MO​ ​-​ ​higher​ ​energy​ ​and​ ​lower​ ​stability​ ​than​ ​the​ ​atomic
orbitals​ ​it​ ​came​ ​from
...
#​ ​of​ ​MOs​ ​=​ ​#​ ​of​ ​atomic​ ​orbitals​ ​combined
2
...

3
...
Each​ ​MO​ ​holds​ ​2​ ​electrons​ ​with​ ​opposite​ ​spins
5
...

6
...
Abhi​ ​im​ ​dying​ ​pls​ ​halp​ ​|​ ​dude​ ​do​ ​you​ ​wanna​ ​find​ ​some​ ​time​ ​this​ ​week
to​ ​study​ ​together?
● Bond​ ​order​ ​=​ ​½​ ​(#​ ​of​ ​e-​ ​in​ ​bonding​ ​MOs​ ​-​ ​#​ ​of​ ​e-​ ​in​ ​antibonding​ ​MOs)
Bond​ ​order

Type​ ​of​ ​Bond

1

Single

2

Double

3

Triple

● All​ ​e-​ ​paired​ ​=​ ​diamag
...
​ ​Orb
...
​ ​ ​Higher​ ​bond​ ​dissociation
energy​ ​(more​ ​energy​ ​to​ ​break)
○ Shortest​ ​bond​ ​length:​ ​H​2
○ Least​ ​stable:​ ​He​2​ ​(in​ ​fact​ ​it​ ​does​ ​not​ ​exist)
○ Highest​ ​bond​ ​dissoc
...
​ ​Orb
...
​ ​energy
● σ​2p​ ​and​ ​𝜋​2p​​ ​switch​ ​places​ ​between​ ​N​ ​and​ ​O

● Quiz
○ Ch
...
​ ​4-​ ​know​ ​how​ ​to​ ​do​ ​beer’s​ ​law​ ​calculations,​ ​best​ ​wavelength​ ​to​ ​do
calculation,​ ​figure​ ​out​ ​concentrations​ ​from​ ​beer’s​ ​law​ ​plot
...
375​ ​(for​ ​example),​ ​solve​ ​for​ ​K
...
​ ​ ​%Trans=​ ​-log​ ​A
○ SA​ ​4​ ​-​ ​formal​ ​charge,​ ​lewis​ ​dot​ ​structure
○ SA​ ​5​ ​-​ ​Models​ ​360:​ ​bond​ ​lengths​ ​in​ ​Angstroms​ ​(10​-10​​ ​m),​ ​calculate​ ​bond
order
...




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