Chapp in a hat

7/29/2019 Chapp in a hat

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1-1William H. Brown Beloit College

William H. BrownChristopher S. FooteBrent L. Iverson

Eric Anslynhttp://academic.cengage.com/chemistry/brown

Chapter 1Covalent Bonding and Shapes of Molecules


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Organic Chemistry

The study of the compounds of carbon.

Over 10 million compounds have been identified. About 1000 new ones are identified each day!

C is a small atom. It forms single, double and triple bonds.

It is intermediate in electronegativity (2.5).

It forms strong bonds with C, H, O, N, and many metals.


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Schematic View of an Atom

A small dense nucleus,diameter 10-14 - 10-15 m,

which contains positivelycharged protons and mostof the mass of the atom.

An extranuclear space,diameter 10-10 m, whichcontains negativelycharged electrons.


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Some useful terms

Shell: Define the probability of finding the e- invarious regions of the space. The energy of the e-

in the shells is quantized

Quantization: only specific values of energy arepossible. The shells can ONLY occur at quantized

energy in which three effects balance each other:electrostatic attraction, electrostatic repulsion andthe wavelike nature of e- that prefers to be

delocalized Delocalization: spreading of electron density over

a large volume of space (away from the nuclei)


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Electron Configuration of Atoms

Electrons are confined to regions of space calledprincipal energy levels (shells).

Each shell can hold 2n2 electrons (n= 1,2,3,4......).

Shell

Number ofElectrons S hell

Can Hold

Relative Energiesof Electrons

in These Shells

32

1882

4

321

higher

lower


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Shells are divided into subshells called orbitals,which are designated by the letters s, p, d, f........

s(one per shell)

p(set of three per shell 2 and higher)

d(set of five per shell 3 and higher) ..

The distribution of Orbitals in Shells

Shell O rbitals Contained in That Shell

3

2

1 1s

2s , 2px, 2py, 2pz

3s , 3px, 3py, 3pz, p lus five 3d orbitals


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Aufbau (Build-Up) Principle: Orbitals fill in order of increasing energy from lowest

energy to highest energy.

Pauli Exclusion Principle:

No more than two electrons may be present in anorbital. If two electrons are present, their spins must bepaired.

Hunds Rule:

When orbitals of equal energy are available but thereare not enough electrons to fill all of them, one electronis added to each orbital before a second electron isadded to any one of them; the spins of the electrons indegenerateorbitals (same energy) should be aligned.


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The pairing of electron spins


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Electron Configuration

Energy-level diagram: A pictorial designation ofwhere electrons are placed in an electronconfiguration. For example, the energy-level diagram

for the ground-state electron configuration of carbonis 1s2 2s2 2p2. For chlorine: 1s2 2s2 2p6 3s2 3p5.

1s

2s

2p

Energy-leveldiagram for chlorine(atomic number 17)

Energy

3s

3p

1s

2s

2p

Energy-level

diagram for carbonatomic number 6

Energ

y


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1s

2s

2p

3s

3p

4s

3d

4p

5s

4d

5pE

Approximate order of filling orbitals with electrons


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The Concept of Energy

In the discussion of energy-level diagrams, linesare drawn on the diagram to depict relativeenergies.

Energy: The ability to do work. The higher inenergy an entity, the more work it can perform.

Potential energy: Stored energy. Unstable structures have energy that is waiting to

be released if given the opportunity. When the

energy is released, work is done, such as, burninggasoline to drive the pistons in an internalcombustion engine that propels the automobile.


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In the ground state of carbon, electrons are placedin accordance with the quantum chemistry

principles (aufbau, Hunds rule, Pauli exclusionprinciple, etc.) that dictate the lowest energy form of

carbon.

If we place the electrons in a different manner (as

for example with one electron in the 2s and three

electrons in the 2p) we would have a higher energylevel referred to as an excited state. When theelectrons are rearranged back to the ground state,

energy is released.


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Electrons in the lowest energy orbital, 1s, are heldtightest to the nucleus and are the hardest toremove from the atom.

First ionization energy: The energy needed to

remove the most loosely held electron from anatom or molecule.


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Lewis Dot Structures (Simple bonding theories)

Lewis electron-dot diagrams are very simplifiedbut very useful modelsfor analyzing bonding in

moleculesValence electrons are those in the outer shell of an atomand they are the electrons involved in chemical reactions

and bonding

The Lewis symbol is the elements symbolplus one dot per valence electron

S......

[Ne]3s23p4


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Lewis Dot Structures

Table 1.4. Lewis Dot Structures for Elements 1-18

N OB

H

Li Be

N a

He

Cl

F

S

N e

A r

C

S iA l P

1A 2A 3A 4A 5A 6A 7A 8A

M g

.

.

. .

..

. .

. .

..

.

:

:

:

::::::::

::::::.:::

:. .. ::

:::..

.

.

..


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Lewis Model of Bonding

Atoms interact in such a way that each participating atomacquires an electron configuration that is the same as that ofthe noble gas nearest it in atomic number.

An atom that gains electrons becomes an anion.

An atom that loses electrons becomes a cation.

The attraction of anions and cations leads to the formationof ionic solids. This ionic interaction is often referred to asan ionic bond.

An atom may share electrons with one or more atoms tocomplete its valence shell; a chemical bond formed bysharing electrons is called a covalent bond. Bonds may bepartially ionic or partially covalent; these bonds are called

polar covalent bonds


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Electronegativity

Electronegativity: A measure of an atoms attraction for the electrons it

shares with another atom in a chemical bond.

Pauling scale

Generally increases left to right in a row.

Generally increases bottom to top in a column.


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Some important EN values

H

2,2

*

Li1,0

Be1,6

B2,0

C2,5

N3,0

O3,5

F4,0

*

Na

0,9

Mg

1,3

Al

1,6

Si

1,9

P

2,2

S

2,6

Cl

3,2

*

K

0,8

Ca

1,0

Sc

1,4

Ge

2,0

As

2,2

Se

2,5

Br

3,0

Kr

3,3

Rb

0,8

Sr

0,9

Y

1,2

Sn

1,9

Sb

2,0

Te

2,1

I

2,7

Xe

3,0

Cs

0,8

Ba

0,9


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Formation of Ions

A rough guideline: Ions will form if the difference in electronegativity between

interacting atoms is 1.9 or greater.

Example: sodium (EN 0.9) and fluorine (EN 4.0)

We use a single-headed (barbed) curved arrow to showthe transfer of one electron from Na to F.

In forming Na+F-, the single 3s electron from Na istransferred to the partially filled valence shell of F.

Na + F Na+

F-

+ F-(1s22s22p6)Na +(1s22s22p6)F(1s22s22p5)+Na(1s22s22p63s1)


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Covalent Bonds

The simplest covalent bond is that in H2 The single electrons from each atom combine to form

an electron pair.

The shared pair functions in two ways simultaneously;it is shared by the two atoms and fills the valence shell

of each atom.

The number of shared pairs

One shared pair forms a single bond

Two shared pairs form a double bond

Three shared pairs form a triple bond

H H H-H+ H0 = -435 kJ (-104 kcal)/mol


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Polar and Nonpolar Covalent Bonds

Although all covalent bonds involve sharing ofelectrons, they differ widely in the degree of sharing.

We divide covalent bonds into

nonpolar covalent bonds and

polar covalent bonds.

Difference in

Electronegativity

Between Bonded Atoms Typ e of Bond

Less than 0.5

0.5 to 1.9

Greater than 1.9

N onpolar covalent

Polar covalent

Ions f orm


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Polar and Nonpolar Covalent Bonds

An example of a polar covalent bond is that of H-Cl.

The difference in electronegativity between Cl and H is

3.0 - 2.1 = 0.9.

We show polarity by using the symbols + and -, or byusing an arrow with the arrowhead pointing toward the

negative end and a plus sign on the tail of the arrow atthe positive end.

H Cl

+ -

H Cl


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Polar Covalent Bonds

Bond dipole moment ( ): A measure of the polarity of a covalent bond.

The product of the charge on either atom of a polar

bond times the distance between the two nuclei.

The table shows average bond dipole moments ofselected covalent bonds.

H-OH-NH-C

H-S C-I

C-FC-ClC-Br C-N

-

C-O

C=N

C=O

-

BondDipole

(D) Bond

1.4

1.5

1.5

1.4

1.2

Bond

BondDipole

(D)

1.3

0.3 0.7

0.2

0.7

2.3

3.5

BondDipole

(D)Bond


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Octet Rule

Atoms tend to gain, lose or share electrons

until they are surrounded by eight valence electrons(i.e., until they resemble a noble gas)

Molecules share pairs of electrons in bondsand may also have lone pairs

: :O

H H

C OO:

::

:


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Drawing Lewis Structures

1. Sum the valence electrons from all atoms For anions add one electron for each negative charge

For cations subtract one electron for each positive charge

2. Connect atoms with single bonds (except for simple moleculesit has to be determined experimentally)

3. Complete octets for atoms bonded to central atom (except H - 2e)

4. Place leftover electrons on central atom, even if this goes beyond octet

5. If not enough electrons to provide octet to central atom, try multiplebonds

(alternative methods can be found in different books )


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Lewis Structures - Table 1.8

In neutral molecules

hydrogen has one bond.

carbon has 4 bonds and no lone pairs.

nitrogen has 3 bonds and 1 lone pair.

oxygen has 2 bonds and 2 lone pairs. halogens have 1 bond and 3 lone pairs.

H2O (8) NH3 (8)CH4 (8) HCl (8)

C2H4 (12) C2H2 (10) CH2O (12) H2CO3 (24)

H-O-H H-N-H

H

H-C-H

H

H H-Cl

H-C C-H

H

H

C O

H

H

C C

H

HO O

CH H

O

Ethylene

Hydrogen chlorideMethaneAmmoniaWater

Carbonic acidFormaldehydeAcetylene


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Formal Charge

Formal charge: The charge on an atom in amolecule or a polyatomic ion.

- Apparent electronic charge of each atom in a Lewisstructure (or the charge an atom would have if electronpairs were shared equally)- Each atom owns one e- of a bonding pair- A purely covalent model


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Formal Charge

To derive formal charge1. Write a correct Lewis structure for the molecule or ion.

2. Assign each atom all its unshared (nonbonding) electrons

and one-half its shared (bonding) electrons.

3. Compare this number with the number of valenceelectrons in the neutral, unbonded atom.

4. The sum of all formal charges is equal to the total charge

on the molecule or ion.

Number ofvalence electrons

in the neutral,

unbonded atom

Allunshared

electrons

One half ofall shared

electrons

+Formalcharge

=


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Formal Charge

C N-

C: (4 valence electrons) - (2 non bonding + 3 bonding) = -1N: (5 valence electrons) - (2 non bonding + 3 bonding) = 0

-1 0

Number ofvalence electrons

in the neutral,unbonded atom

Allunshared

electrons

One half ofall shared

electrons+

Formalcharge

=


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Formal Charge

C O

OC N

N

CN O

FormalCharge C N O

+ 1

- 1

0

C


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Formal Charge

Example: Draw Lewis structures, and show which atom ineach bears the formal charge.

NH2-

HCO3-

CO32 -

CH3 NH3+

HCOO-

CH3COO-

(b) (c)

(d) (e) (f)

(a)

A E i h O R l


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Apparent Exceptions to the Octet Rule

Molecules that contain atoms of Group 3Aelements, particularly boron and aluminum.

::

:

F B

F

F

Cl Al

Cl

Cl

6 electrons in thevalence shells of boron

and aluminum

Boron trifluoride Aluminum chloride

: :

: :

: :

:

:

:

:

:

:

:

::

O


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Expanded shells (hypervalence)

When it is impossible to write a structure consistent with the octet rule

increase the number of electrons around the central atom

Cl P

Cl

Cl

Cl

Cl

10e around P

Only for elements from 3rd row (P, S) and heavier, which can make use of empty dorbitals

(Maximum 18e: 2s, 6p, 10d)

A E i h O R l


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The P in trimethylphosphine obeys the octet rule by havingthree bonds and one unshared pair.

A common depiction of phosphoric acid, however, has fivebonds to P, which is explained by invoking the use of 3dorbitals to accommodate the additional bonds.

:

:

Phosphorus

pentachloride

Phosphoric

acid

P

Cl

Cl Cl

Cl ClCH3 -P-CH3

CH3

Trimethyl-

phosphine

H-O-P-O-H

O

O-H

:

:

:

:

:

:

:

:

:

:

:

::

::

:

:

:

:

:

:

F ti l G


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Functional Groups

Functional group: An atom or group of atomswithin a molecule that shows a characteristic set ofphysical and chemical properties.

Functional groups are important for three reasons;they are:

1. the units by which we divide organic compounds intoclasses.

2. the sites of characteristic chemical reactions.

3. the basis for naming organic compounds.

Al h l


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Alcohols

Contain an -OH (hydroxyl) group bonded to atetrahedral carbon atom.

Ethanol may also be written as a condensed

structural formula.

H-C-C-O-H

H

H

H

H:

:-C-O-H

Ethanol

(an alcohol)

Functional

group

CH3 -CH2 -OH CH3 CH2 OHor

Al h l


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Alcohols

Alcohols are classified as primary (1), secondary (2), ortertiary (3) depending on the number of carbon atomsbonded to the carbon bearing the -OH group.

CH3 -C-OHCH

3-C-OH

CH3

H

CH3-C-OH

CH3

CH3

A 1 alcohol A 2 alcohol A 3 alcohol

H

H

Al h l


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Alcohols

There are two alcohols with molecular formula C3H8O.

CH3

CH2

CH2

OH

CH3CHCH3

OH

H-C-C-C-O-H

H

H

H

H

H

H

C-C-C-HH

HO H

H H H

H

or

or

a 2 alcohol

a 1 alcohol

A i


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Amines

Contain an amino group; an sp3-hybridized nitrogenbonded to one, two, or three carbon atoms.

An amine may be 1, 2, or 3.

CH3 N H

H

CH3 N H

CH3

CH3 N CH3

CH3

Methylamine(a 1 amine)

Dimethylamine(a 2 amine)

Trimethylamine(a 3 amine)

: : :

Ald h d d K t


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Aldehydes and Ketones

Contain a carbonyl (C=O) group.

C H

O

Functional

group

Acetone

(a ketone)

CH3 -C- H CH3 -C-CH3

O O

C

O

Functional

group

Acetaldehyde(an aldehyde)

::::

C rb li A id


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Carboxylic Acids

Contain a carboxyl (-COOH) group.

C O

O

H CH3 -C-O-H

O

CH3COOH CH3 CO2 H

: :

::

or or

Acetic acid

(a carboxylic acid )

Functionalgroup

Carboxylic Esters


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Carboxylic Esters

Ester: A derivative of a carboxylic acid in which thecarboxyl hydrogen is replaced by a carbon group.

C O

O

Functional

group

CH3

-C-O-CH2

-CH3

Ethyl acetate

(an ester)

O

Carboxylic Amide


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Carboxylic Amide

Carboxylic amide, commonly referred to as anamide: A derivative of a carboxylic acid in which the -OH of the -COOH group is replaced by an amine.

The six atoms of the amide functional group lie in a planewith bond angles of approximately 120.

CH3 -C-N-H

H

Acetamide(a 1 amid e)

O

C N

O

Functionalgroup

V l Sh ll El t P i R l i Th (VSEPR)


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Valence Shell Electron Pair Repulsion Theory (VSEPR)

Electrons in molecules appear in bonding pairs or lone pairs

Each pair of electrons repels all other pairs

Molecules adopt geometries with electron pairs as far from each other aspossible

Electron pairs define regions of space where they are likely to be:Between nuclei for bonding pairsClose to one nucleus for lone pairs

those regions are called electron domains

(a very approximate but very useful way of predicting molecular shapes)



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Geometries of electron

domains



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Molecular Geometries



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Lone pairs are largerthan bonding pairs

VSEPR Model


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VSEPR Model

Example: predict all bond angles for these molecules andions.

( a ) NH4+ ( b ) CH3 NH2

( f ) H2 CO3 (g) HCO3-( e) CH3 CH= CH2

( i) CH3COOH( h) CH3 CHO

( d) CH3 OH

( j ) BF4-

Polar and Nonpolar Molecules


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To determine if a molecule is polar, we need todetermine

if the molecule has polar bonds and

the arrangement of these bonds in space.

Molecular dipole moment (): The vector sum of theindividual bond dipole moments in a molecule.

reported in Debyes (D)

Electrostatic Potential (elpot) Maps


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Relative electron density distribution in moleculesis important because it allows us to identify sites

of chemical reactivity.

Many reactions involve an area of relatively high

electron density on one molecule reacting with an areaof relatively low electron density on another molecule.

It is convenient to keep track of overall molecular

electron density distributions using computer graphics.



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In electrostatic potential maps (elpots) (chargedensities are easily computed by programs such as spartan)

Areas of relatively high calculated electron density are

shown in red. Areas of relatively low calculated electron density are

shown in blue.

Intermediate electron densities are represented byintermediate colors.



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These molecules have polar bonds, but eachmolecule has a zero dipole moment.

O C O

Carbon dioxide

= 0 D

B

F

F

F

Boron trifluoride

= 0 D

C

Cl

ClClCl

Carbon tetrachloride

= 0 D



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These molecules have polar bonds and are polarmolecules.

N

HH

H

O

H H

Water

= 1.85D

Ammonia

= 1.47D

directionof dip ole

moment

directionof dip ole

moment



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Formaldehyde has polar bonds and is a polar molecule.

Formaldehyde= 2.33 D

directionof dipole

moment H HC

O

Quantum or Wave Mechanics


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Albert Einstein: E = h (energy is quantized) light has particle properties.

Louis deBroglie: wave/particle duality

Erwin Schrdinger: wave equation

wave function, : A solution to a set of equations thatdepicts the energy of an electron in an atom.

each wave function is associated with a unique set ofquantum numbers.

each wave function represents a region of three-dimensional space and is called an orbital.

2

is the probability of finding an electron at a given pointin space.

hm


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For most aspects of organic chemistry, it is best to considerthe wavelike properties of electrons. In this course, weconcentrate on wave functions and shapes associated with s

and p orbitals because they are the orbitals most ofteninvolved in covalent bonding in organic compounds.

When we describe orbital interactions, we are referring tointeractions of waves. Waves interact constructively ordestructively (adding or subtracting). When two wavesoverlap, positive phasing adds constructively with positivephasing. Positive and negative phasing add destructively,meaning they cancel.

Shapes of Atomic s and p Orbitals


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Shapes of Atomic sand pOrbitals

All s orbitals have theshape of a sphere withthe center of the sphereat the nucleus.

Figure 1.8 (a) Calculatedand (b) cartoonrepresentations showingan arbitrary boundary

surface containing about95% of the electrondensity.



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Figure 1.9 (a) Three-dimensional representations of the2px, 2py, and 2pz atomic orbitals computed using theSchrdinger equation. Nodal planes are shaded.



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Figure 1.9(b) Cartoon representations of the 2px, 2py, and 2pzatomic orbitals.

Molecular Orbital Theory


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MO theory begins with the hypothesis thatelectrons in atoms exist in atomic orbitals and

electrons in molecules exist in molecular orbitals.



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Rules: Combination of n atomic orbitals (mathematically adding

and subtracting wave functions) gives n MOs (new wave

functions). MOs are arranged in order of increasing energy.

MO filling is governed by the same rules as for atomic

orbitals: Aufbau principle: fill beginning with LUMO (lowest

unoccupied molecular orbital)

Pauli exclusion principle: no more than 2e- in a MO Hunds rule: when two or more MOs of equivalent

energy are available, add 1e- to each before filling any

one of them with 2e-.



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Figure 1.10 MOs derived from combination by (a)addition and (b) subtraction of two 1satomic orbitals.

Bonding-Combined VB&MO


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Bonding Combined VB&MO

Bonding molecular orbital: A MO in whichelectrons have a lower energy than they wouldhave in isolated atomic orbitals.

Sigma () bonding molecular orbital: A MO inwhich electron density is concentrated betweentwo nuclei along the axis joining them and iscylindrically symmetrical.

Covalent Bonding-Combined VB & MO


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Covalent Bonding Combined VB & MO

Antibonding MO: A MO in which electrons have ahigher energy than they would in isolated atomicorbitals.

Sigma star () antibonding molecular orbital: A MOin which population with electrons actually causesrepulsion of the nuclei involved.

Covalent Bonding-Combined VB & MO


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Covalent Bonding Combined VB & MO

Figure 1.11 A MO energy diagram for H2. (a) Groundstate and (b) lowest excited state.


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VB: Hybridization of Atomic Orbitals


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What we find instead are bond angles of approximately109.5 in molecules with only single bonds, 120 inmolecules with double bonds, and 180 in moleculeswith triple bonds.

To account for the observed bond angles, LinusPauling proposed that atomic orbitals for each atomshould be thought of as first combining to form newatomic orbitals, called hybrid orbitals, which theninteract to form bonds by overlapping with orbitals fromother atoms.

Hybrid orbitals are formed by combinations of atomicorbitals by a process called hybridization.Mathematically, this is accomplished by combining

wave functions for the 2sand 2porbital.


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y

The mathematical combination of one 2s atomicorbital wave function and two 2p atomic orbitalwave functions forms three equivalent sp2 hybrid

orbitals.

sp 2 Hybridization, with electron

population for carbon to form double

bonds

2s

2p

sp2Energy 2p


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Consider BH3. VSEPRtells us that BH3 istrigonal planar, with

120 H-B-H bondangles. In BH3 theunhybridized 2p orbital

is empty.

y


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Figure 1.16 spHybrid orbitals and two 2porbitals onan sphybridized atom.

y


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Combining VB & MO Theories


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g

To create orbitals that are localized betweenadjacent atoms, we add and subtract the atomicorbitals on the adjacent atoms, which are aligned

to overlap with each other. Consider methane, CH4. The sp

3 hybrid orbitals ofcarbon each point to a 1sorbital of hydrogen and,therefore, we add and subtract these atomicorbitals to create molecular orbitals.

As with H2, one resulting MO is lower in energythan the two separated atomic orbitals, and iscalled a bonding orbital. The other is higher in

energy and is antibonding.



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g

Figure 1.17 Molecular orbital mixing diagram forcreation of any C-C bond.


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A carbon-carbon triplebond consists of one bond formed by overlap

of sp hybrid orbitals andtwo bonds formed bythe overlap of parallel

2patomic orbitals.

Covalent Bonding of Carbon


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H-C C-H

C C

H-C-C-HH

H

H

H

H

HH

H

OrbitalHybrid-

ization

Types ofBonds

to Each Carbon Example

sp3 four bonds

sp 2 three bondsand one bond

sp two bondsand two bonds

Ethane

Ethene

Ethyne

Name

PredictedBond

Angles

109.5

120

180

GroupsBonded

to Carbon

4

3

2

Resonance


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For many molecules and ions, no single Lewisstructure provides a truly accurate representation.

Ethanoate ion

(acetate ion)

C

O

O

H3CC

O

O

H3C

-

-

and

:

: :

: :

::

:

:

:

Resonance


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Linus Pauling - 1930s Many molecules and ions are best described by writing

two or more Lewis structures.

Individual Lewis structures are called contributingstructures.

Connect individual contributing structures by double-

headed (resonance) arrows. The molecule or ion is a hybrid of the various

contributing structures.


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Resonance


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Curved arrow: A symbol used to show theredistribution of valence electrons.

In using curved arrows, there are only two allowedtypes of electron redistribution:

from a bond to an adjacent atom. from a lone pair on an atom to an adjacent bond.

Electron pushing is a survival skill in organicchemistry.

learn it well!


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Resonance


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All contributing structures must1. have the same number of valence electrons.

2. obey the rules of covalent bonding:

no more than 2 electrons in the valence shell of H.

no more than 8 electrons in the valence shell of a 2ndperiod element.

3. differ only in distribution of valence electrons; the position ofall nuclei must be the same.

4. have the same number of paired and unpaired electrons.


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93/99

Resonance


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1-94

Preference 2: maximum number of covalentbonds

Structures with a greater number of covalent bondscontribute more than those with fewer covalent bonds.

Greater contribution(8 covalent bonds)

Lesser contribution(7 covalent bonds)

+ +CCH3 OCH3 O

H

HC

H

H

Resonance


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1-95

Preference 3: least separation of unlike charge

Structures with separation of unlike charges contribute less

than those with no charge separation.

CH3 -C-CH3 CH3 -C-CH3

Greater contribution(no separation of

unlike charges)

Lesser contribution(separation of unlike

charges)

O -O::

:::


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Bond Lengths and Bond Strengths


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1-98

H-C-C-H

H

H

H

H

H

C C

H

H H

H-C C-H

C-C

C-C

C-C

C-H

C-H

C-H

Formula BondOrbital

Overlap

Bond Length(pm)

sp3

-sp3

sp2-1s

sp -sp , two 2p-2p

sp -1s

sp3-1s

sp2-sp2, 2p-2p

153.2

111.4

133.9

110.0

121.2

109.0

Ethane

Ethene

Ethyne

Name


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Covalent

Bonds &

Shapes of

Molecules

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