Química!Orgânica!I! Aula!2! - lqbo.ufscar.br · tracted to the positively charged nucleus of the...

Química Orgânica I

Aula 2

Prof. Marco Antonio B Ferreira [email protected]

3351-‐8075

1

UNIVERSIDADE FEDERAL DE SÃO CARLOS

www.lqbo.ufscar.br

2

Teoria de Ligação de Valência •  Como átomos formam ligações?

Segundo Lewis: Busca por ter sua camada de valência completa a parFr do comparFlhamento de elétrons.

TLV: Combinação de 2 orbitais atômicos SEMI-‐PREENCHIDOS leva a formação de um novo orbital.

20 C H A P T E R 1 Electronic Structure and Bonding • Acids and Bases

1.6 An Introduction to Molecular Orbital Theory

How do atoms form covalent bonds in order to form molecules? The Lewis model,which describes how atoms attain a complete octet by sharing electrons, tells us onlypart of the story. A drawback of the model is that it treats electrons like particles anddoes not take into account their wavelike properties.

Molecular orbital (MO) theory combines the tendency of atoms to fill their octetsby sharing electrons (the Lewis model) with their wavelike properties—assigningelectrons to a volume of space called an orbital. According to MO theory, covalentbonds result from the combination of atomic orbitals to form molecular orbitals—orbitals that belong to the whole molecule rather than to a single atom. Like an atomicorbital that describes the volume of space around the nucleus of an atom where anelectron is likely to be found, a molecular orbital describes the volume of space arounda molecule where an electron is likely to be found. Like atomic orbitals, molecular or-bitals have specific sizes, shapes, and energies.

Let’s look first at the bonding in a hydrogen molecule As the 1s atomic orbitalof one hydrogen atom approaches the 1s atomic orbital of a second hydrogen atom,they begin to overlap. As the atomic orbitals move closer together, the amount of over-lap increases until the orbitals combine to form a molecular orbital. The covalent bondthat is formed when the two s atomic orbitals overlap is called a sigma bond. A bond is cylindrically symmetrical—the electrons in the bond are symmetrically dis-tributed about an imaginary line connecting the centers of the two atoms joined by thebond. (The term comes from the fact that cylindrically symmetrical molecular or-bitals possess symmetry.)

During bond formation, energy is released as the two orbitals start to overlap, be-cause the electron in each atom not only is attracted to its own nucleus but also is at-tracted to the positively charged nucleus of the other atom (Figure 1.2). Thus, theattraction of the negatively charged electrons for the positively charged nuclei is whatholds the atoms together. The more the orbitals overlap, the more the energy decreases

H H1s atomic

orbital1s atomic

orbital

H H H Hmolecular orbital

=

ss

s1S2(H2).

0

104 kcal/molPote

ntia

l ene

rgy

0.74 ÅInternuclear distance

bond length!104 kcal/mol

bonddissociationenergy

+

!

" hydrogenatoms are closetogether

" hydrogenatoms are farapart

Figure 1.2 NThe change in energy that occurs astwo 1s atomic orbitals approacheach other. The internucleardistance at minimum energy is thelength of the covalent bond.H¬ H

Movie:bond formationH2

BRUI01-001_059r4 20-03-2003 2:58 PM 0

Ex.: Formação da molécula de H2. Ligação covalente sigma (σ)

Section 1.6 An Introduction to Molecular Orbital Theory 21

Maximum stability corresponds to mini-mum energy.

* Joules are the Système International (SI) units for energy, although manychemists use calories. We will use both in this book.

1 kcal = 4.184 kJ.

+

+

!

++

nucleusof the hydrogenatom

node

phase of the orbital


+

!

waves reinforce each other, resulting in bonding

waves cancel each other, andno bond forms

destructive combination

constructive combination > Figure 1.3The wave functions of twohydrogen atoms can interact toreinforce, or enhance, each other(top) or can interact to cancel eachother (bottom). Note that wavesthat interact constructively are in-phase, whereas waves that interactdestructively are out-of-phase.

until the atoms approach each other so closely that their positively charged nuclei startto repel each other. This repulsion causes a large increase in energy. We see that max-imum stability (i.e., minimum energy) is achieved when the nuclei are a certain dis-tance apart. This distance is the bond length of the new covalent bond. The length ofthe bond is 0.74

As Figure 1.2 shows, energy is released when a covalent bond forms. When thebond forms, (or 435 kJ mol)* of energy is released. Breaking the

bond requires precisely the same amount of energy. Thus, the bond strength—alsocalled the bond dissociation energy—is the energy required to break a bond, or theenergy released when a bond is formed. Every covalent bond has a characteristic bondlength and bond strength.

Orbitals are conserved—the number of molecular orbitals formed must equal thenumber of atomic orbitals combined. In describing the formation of an bond,however, we combined two atomic orbitals, but discussed only one molecular orbital.Where is the other molecular orbital? It is there, but it contains no electrons.

Atomic orbitals can combine in two different ways: constructively and destructive-ly. They can combine in a constructive, additive manner, just as two light waves orsound waves may reinforce each other (Figure 1.3). This is called a (sigma) bond-ing molecular orbital. Atomic orbitals can also combine in a destructive way, cancel-ing each other. The cancellation is similar to the darkness that occurs when two lightwaves cancel each other or to the silence that occurs when two sound waves canceleach other (Figure 1.3). This destructive type of interaction is called a antibondingmolecular orbital. An antibonding orbital is indicated by an asterisk 1*2.

S*

S

H¬H

>104 kcal>molH¬H

Å.H¬H

The bonding molecular orbital and antibonding molecular orbital are shownin the molecular orbital diagram in Figure 1.4. In an MO diagram, the energies are rep-resented as horizontal lines; the bottom line is the lowest energy level, the top line thehighest energy level. We see that any electrons in the bonding orbital will most likelybe found between the nuclei. This increased electron density between the nuclei iswhat binds the atoms together. Because there is a node between the nuclei in the anti-bonding molecular orbital, any electrons that are in that orbital are more likely to befound anywhere except between the nuclei, so the nuclei are more exposed to one an-other and will be forced apart by electrostatic repulsion. Thus, electrons that occupythis orbital detract from, rather than aid, the formation of a bond between the atoms.

s*s

BRUI01-001_059r4 20-03-2003 2:58 PM Page 21

3


1.6 An Introduction to Molecular Orbital Theory

How do atoms form covalent bonds in order to form molecules? The Lewis model,which describes how atoms attain a complete octet by sharing electrons, tells us onlypart of the story. A drawback of the model is that it treats electrons like particles anddoes not take into account their wavelike properties.

Molecular orbital (MO) theory combines the tendency of atoms to fill their octetsby sharing electrons (the Lewis model) with their wavelike properties—assigningelectrons to a volume of space called an orbital. According to MO theory, covalentbonds result from the combination of atomic orbitals to form molecular orbitals—orbitals that belong to the whole molecule rather than to a single atom. Like an atomicorbital that describes the volume of space around the nucleus of an atom where anelectron is likely to be found, a molecular orbital describes the volume of space arounda molecule where an electron is likely to be found. Like atomic orbitals, molecular or-bitals have specific sizes, shapes, and energies.

Let’s look first at the bonding in a hydrogen molecule As the 1s atomic orbitalof one hydrogen atom approaches the 1s atomic orbital of a second hydrogen atom,they begin to overlap. As the atomic orbitals move closer together, the amount of over-lap increases until the orbitals combine to form a molecular orbital. The covalent bondthat is formed when the two s atomic orbitals overlap is called a sigma bond. A bond is cylindrically symmetrical—the electrons in the bond are symmetrically dis-tributed about an imaginary line connecting the centers of the two atoms joined by thebond. (The term comes from the fact that cylindrically symmetrical molecular or-bitals possess symmetry.)

During bond formation, energy is released as the two orbitals start to overlap, be-cause the electron in each atom not only is attracted to its own nucleus but also is at-tracted to the positively charged nucleus of the other atom (Figure 1.2). Thus, theattraction of the negatively charged electrons for the positively charged nuclei is whatholds the atoms together. The more the orbitals overlap, the more the energy decreases

H H1s atomic

orbital1s atomic

orbital

H H H Hmolecular orbital

=

ss

s1S2(H2).

0

104 kcal/molPote

ntia

l ene

rgy

0.74 ÅInternuclear distance

bond length!104 kcal/mol

bonddissociationenergy

+

!

" hydrogenatoms are closetogether

" hydrogenatoms are farapart

Figure 1.2 NThe change in energy that occurs astwo 1s atomic orbitals approacheach other. The internucleardistance at minimum energy is thelength of the covalent bond.H¬ H

Movie:bond formationH2

BRUI01-001_059r4 20-03-2003 2:58 PM Page 20

Teoria de Ligação de Valência


Maximum stability corresponds to mini-mum energy.

* Joules are the Système International (SI) units for energy, although manychemists use calories. We will use both in this book.

1 kcal = 4.184 kJ.

+

+

!

++

nucleusof the hydrogenatom

node



+

!

waves reinforce each other, resulting in bonding

waves cancel each other, andno bond forms

destructive combination

constructive combination > Figure 1.3The wave functions of twohydrogen atoms can interact toreinforce, or enhance, each other(top) or can interact to cancel eachother (bottom). Note that wavesthat interact constructively are in-phase, whereas waves that interactdestructively are out-of-phase.

until the atoms approach each other so closely that their positively charged nuclei startto repel each other. This repulsion causes a large increase in energy. We see that max-imum stability (i.e., minimum energy) is achieved when the nuclei are a certain dis-tance apart. This distance is the bond length of the new covalent bond. The length ofthe bond is 0.74

As Figure 1.2 shows, energy is released when a covalent bond forms. When thebond forms, (or 435 kJ mol)* of energy is released. Breaking the

bond requires precisely the same amount of energy. Thus, the bond strength—alsocalled the bond dissociation energy—is the energy required to break a bond, or theenergy released when a bond is formed. Every covalent bond has a characteristic bondlength and bond strength.

Orbitals are conserved—the number of molecular orbitals formed must equal thenumber of atomic orbitals combined. In describing the formation of an bond,however, we combined two atomic orbitals, but discussed only one molecular orbital.Where is the other molecular orbital? It is there, but it contains no electrons.

Atomic orbitals can combine in two different ways: constructively and destructive-ly. They can combine in a constructive, additive manner, just as two light waves orsound waves may reinforce each other (Figure 1.3). This is called a (sigma) bond-ing molecular orbital. Atomic orbitals can also combine in a destructive way, cancel-ing each other. The cancellation is similar to the darkness that occurs when two lightwaves cancel each other or to the silence that occurs when two sound waves canceleach other (Figure 1.3). This destructive type of interaction is called a antibondingmolecular orbital. An antibonding orbital is indicated by an asterisk 1*2.

S*

S

H¬H

>104 kcal>molH¬H

Å.H¬H

The bonding molecular orbital and antibonding molecular orbital are shownin the molecular orbital diagram in Figure 1.4. In an MO diagram, the energies are rep-resented as horizontal lines; the bottom line is the lowest energy level, the top line thehighest energy level. We see that any electrons in the bonding orbital will most likelybe found between the nuclei. This increased electron density between the nuclei iswhat binds the atoms together. Because there is a node between the nuclei in the anti-bonding molecular orbital, any electrons that are in that orbital are more likely to befound anywhere except between the nuclei, so the nuclei are more exposed to one an-other and will be forced apart by electrostatic repulsion. Thus, electrons that occupythis orbital detract from, rather than aid, the formation of a bond between the atoms.

s*s

BRUI01-001_059r4 20-03-2003 2:58 PM Page 21

4

Introdução a Teoria do Orbital Molecular A combinação de N orbitais atômicos deve gerar N “orbitais moleculares”.

(LCAM -‐ Linear combina-on of atomic orbital).

σ∗ = cA.φA − cb.φB

σ = cA.φA + cb.φB orbital σ (ligante)

orbital σ* (anY-‐ligante)

-

HA HB

φA φB σ∗

+

+

HA HB H H

φA φB σ

Notar que este orbital não oferece possibilidade atrata4va pois coloca

os orbitais fora da região compreendida da ligação

5

Introdução a Teoria do Orbital Molecular E as energias???

<1 <2

HB 1s1HA 1s1

V

V*

<V

<V* Orbital molecular antiligante

Orbital molecular ligante

LUMO

HOMO

Cálculos teóricos para a molécula de H2

ΔE1

ΔE2

ΔE2 > ΔE1

6

Introdução a Teoria do Orbital Molecular

E o caso do He2????

7

Introdução a Teoria do Orbital Molecular Section 1.6 An Introduction to Molecular Orbital Theory 23

nodes

nodenode

!" antibonding molecular orbital

! bonding molecular orbital

2p atomicorbital

2p atomicorbital

Ener

gy

> Figure 1.5End-on overlap of two p orbitals toform a bonding molecular orbitaland a antibonding molecularorbital.

s*s

Side-to-side overlap of two p atomic or-bitals forms a bond. All other cova-lent bonds in organic molecules are bonds.

S

P

formed. The antibonding molecular orbital has three nodes. (Notice that after eachnode, the phase of the molecular orbital changes.)

Unlike the bond formed as a result of end-on overlap, side-to-side overlap of twop atomic orbitals forms a pi bond (Figure 1.6). Side-to-side overlap of two in-phase p atomic orbitals forms a bonding molecular orbital, whereas side-to-sideoverlap of two out-of-phase p orbitals forms a antibonding molecular orbital. The

bonding molecular orbital has one node—a nodal plane that passes through both nu-clei. The antibonding molecular orbital has two nodal planes. Notice that bondsare cylindrically symmetrical, but bonds are not.

The extent of overlap is greater when p orbitals overlap end-on than when theyoverlap side-to-side. This means that a bond formed by the end-on overlap of p or-bitals is stronger than a bond formed by the side-to-side overlap of p orbitals. It alsomeans that a bonding molecular orbital is more stable than a bonding molecularorbital because the stronger the bond, the more stable it is. Figure 1.7 shows a molec-ular orbital diagram of two identical atoms using their three degenerate atomic orbitalsto form three bonds—one bond and two bonds.ps

psp

s

psp*

pp*

p1P2s

s*

A bond is stronger than a bond.PS

#" antibonding molecular orbital

# bonding molecular orbital

2p atomicorbital

2p atomicorbital

nodal plane

nodal plane

nodal plane

Ener

gy

> Figure 1.6Side-to-side overlap of two parallelp orbitals to form a bondingmolecular orbital and a antibonding molecular orbital.

p*p

BRUI01-001_059r4 20-03-2003 2:58 PM Page 23

.

. .

. .

.

8



nodes

nodenode



2p atomicorbital

2p atomicorbital

Ener

gy


s*s


S

P





psp

s

psp*

pp*

p1P2s

s*




2p atomicorbital

2p atomicorbital

nodal plane

nodal plane

nodal plane

Ener

gy

> Figure 1.6Side-to-side overlap of two parallelp orbitals to form a bondingmolecular orbital and a antibonding molecular orbital.

p*p

BRUI01-001_059r4 20-03-2003 2:58 PM Page 23

.

. .

.

. .

9


Section 1.6An Introduction to Molecular Orbital Theory23

nodes

node node



2p atomicorbital

2p atomicorbital

Energy

>Figure 1.5End-on overlap of two porbitals toform a bonding molecular orbitaland a antibonding molecularorbital.

s*s

Side-to-side overlap of two patomic or-bitals forms a bond. All other cova-lent bonds in organic molecules are bonds.

S

P

formed. The antibonding molecular orbital has threenodes. (Notice that after eachnode,the phase of the molecular orbital changes.)

Unlike the bond formed as a result of end-on overlap,side-to-side overlap of twopatomic orbitals forms a pibond(Figure1.6). Side-to-side overlap of two in-phase patomic orbitals forms a bonding molecular orbital,whereas side-to-sideoverlap of two out-of-phase porbitals forms a antibonding molecular orbital. The

bonding molecular orbital has one node—a nodal plane that passes through both nu-clei. The antibonding molecular orbital has two nodal planes. Notice that bondsare cylindrically symmetrical,but bonds are not.

The extent of overlap is greater when porbitals overlap end-on than when theyoverlap side-to-side. This means that a bond formed by the end-on overlap of por-bitals is stronger than a bond formed by the side-to-side overlap of porbitals. It alsomeans that a bonding molecular orbital is more stable than a bonding molecularorbital because the stronger the bond,the more stable it is. Figure1.7shows a molec-ular orbital diagram of two identical atoms using their three degenerate atomic orbitalsto form three bonds—one bond and two bonds. p s

p sp

s

ps p*

pp*

p1P2 s

s*

A bond is stronger than a bond. P S



2p atomicorbital

2p atomicorbital

nodal plane

nodal plane

nodal plane

Energy

>Figure 1.6Side-to-side overlap of two parallelporbitals to form a bondingmolecular orbital and a antibonding molecular orbital.

p*p

BRUI01-001_059r4 20-03-2003 2:58 PM Page 23

Section 1.7 Bonding in Methane and Ethane: Single Bonds 25

their atomic orbitals, and because electron pairs repel each other, the bonding electronsand lone-pair electrons around an atom are positioned as far apart as possible.

Because organic chemists generally think of chemical reactions in terms of thechanges that occur in the bonds of the reacting molecules, the VSEPR model oftenprovides the easiest way to visualize chemical change. However, the model is inade-quate for some molecules because it does not allow for antibonding orbitals. We willuse both the MO and the VSEPR models in this book. Our choice will depend onwhich model provides the best description of the molecule under discussion. We willuse the VSEPR model in Sections 1.7–1.13.

PROBLEM 14!

Indicate the kind of molecular orbital that results when the orbitals arecombined as indicated:

1.7 Bonding in Methane and Ethane: Single Bonds

We will begin the discussion of bonding in organic compounds by looking at the bond-ing in methane, a compound with only one carbon atom. Then we will examine thebonding in ethane (a compound with two carbons and a carbon–carbon single bond),in ethene (a compound with two carbons and a carbon–carbon double bond), and inethyne (a compound with two carbons and a carbon–carbon triple bond).

Next, we will look at bonds formed by atoms other than carbon that are commonlyfound in organic compounds—bonds formed by oxygen, nitrogen, and the halogens.Because the orbitals used in bond formation determine the bond angles in a molecule,you will see that if we know the bond angles in a molecule, we can figure out whichorbitals are involved in bond formation.

Bonding in MethaneMethane has four covalent bonds. Because all four bonds have the samelength and all the bond angles are the same (109.5°), we can conclude that the four

bonds in methane are identical.Four different ways to represent a methane molecule are shown here.

In a perspective formula, bonds in the plane of the paper are drawn as solid lines,bonds protruding out of the plane of the paper toward the viewer are drawn as solidwedges, and those protruding back from the plane of the paper away from the viewerare drawn as hatched wedges.

C¬H

C¬H(CH4)

+

a.

+b.

+c.

+d.

1s, s*, p, or p*2

H

C

109.5°perspective formula

of methaneball-and-stick model

of methanespace-filling model

of methaneelectrostatic potential

map for methane

H H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 25




PROBLEM 14!








C¬H

C¬H(CH4)

+

a.

+b.

+c.

+d.

1s, s*, p, or p*2

H

C





map for methane

H H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 25

. .

. .

Orbital molecular σ


The MO diagram shows that the bonding molecular orbital is more stable—is lowerin energy—than the individual atomic orbitals. This is because the more nuclei anelectron “feels,” the more stable it is. The antibonding molecular orbital, with lesselectron density between the nuclei, is less stable—is of higher energy—than theatomic orbitals.

After the MO diagram is constructed, the electrons are assigned to the molecularorbitals. The aufbau principle and the Pauli exclusion principle, which apply to elec-trons in atomic orbitals, also apply to electrons in molecular orbitals: Electrons alwaysoccupy available orbitals with the lowest energy, and no more than two electrons canoccupy a molecular orbital. Thus, the two electrons of the bond occupy thelower energy bonding molecular orbital (Figure 1.4), where they are attracted to bothpositively charged nuclei. It is this electrostatic attraction that gives a covalent bond itsstrength. Therefore, the greater the overlap of the atomic orbitals, the stronger is thecovalent bond. The strongest covalent bonds are formed by electrons that occupy themolecular orbitals with the lowest energy.

The MO diagram in Figure 1.4 allows us to predict that would not be as stableas because has only one electron in the bonding orbital. We can also predictthat does not exist: Because each He atom would bring two electrons, wouldhave four electrons—two filling the lower energy bonding molecular orbital and theremaining two filling the higher energy antibonding molecular orbital. The two elec-trons in the antibonding molecular orbital would cancel the advantage to bondinggained by the two electrons in the bonding molecular orbital.

PROBLEM 13!

Predict whether or not exists.

Two p atomic orbitals can overlap either end-on or side-to-side. Let’s first look atend-on overlap. End-on overlap forms a bond. If the overlapping lobes of the p or-bitals are in-phase (a blue lobe of one p orbital overlaps a blue lobe of the other p or-bital), a bonding molecular orbital is formed (Figure 1.5). The electron density ofthe bonding molecular orbital is concentrated between the nuclei, which causes theback lobes (the nonoverlapping lobes) of the molecular orbital to be quite small. The bonding molecular orbital has two nodes—a nodal plane passing through each of thenuclei.

If the overlapping lobes of the p orbitals are out-of-phase (a blue lobe of one p or-bital overlaps a green lobe of the other p orbital), a antibonding molecular orbital iss*

sss

s

He2

+

He2He2

H2

+H2

H2

+

H—H

When two atomic orbitals overlap, twomolecular orbitals are formed—onelower in energy and one higher in ener-gy than the atomic orbitals.

In-phase overlap forms a bonding MO;out-of-phase overlap forms an anti-bonding MO.



1s atomicorbital

1s atomicorbital

Ener

gynodeFigure 1.4 N

Atomic orbitals of and molecularorbitals of Before covalent bondformation, each electron is in anatomic orbital. After covalent bondformation, both electrons are in thebonding molecular orbital. Theantibonding molecular orbital isempty.

H2.H–

BRUI01-001_059r4 20-03-2003 2:58 PM Page 22




PROBLEM 14!








C¬H

C¬H(CH4)

+

a.

+b.

+c.

+d.

1s, s*, p, or p*2

H

C





map for methane

H H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 25

. Orbital molecular σ*

10

E o caso do carbono????

Segundo a TLV: Combinação de 2 orbitais atômicos SEMI-‐PREENCHIDOS leva a formação de um novo orbital.

C (Z = 6) 1s2 2s2 2px1 2py1

C

Lewis

C+ 2H

HH C H

H

Temos aqui uma falha no modelo

Section 1.6An Introduction to Molecular Orbital Theory23

nodes

node node



2p atomicorbital

2p atomicorbital

Energy

>Figure 1.5End-on overlap of two porbitals toform a bonding molecular orbitaland a antibonding molecularorbital.

s*s

Side-to-side overlap of two patomic or-bitals forms a bond. All other cova-lent bonds in organic molecules are bonds.

S

P

formed. The antibonding molecular orbital has threenodes. (Notice that after eachnode,the phase of the molecular orbital changes.)

Unlike the bond formed as a result of end-on overlap,side-to-side overlap of twopatomic orbitals forms a pibond(Figure1.6). Side-to-side overlap of two in-phase patomic orbitals forms a bonding molecular orbital,whereas side-to-sideoverlap of two out-of-phase porbitals forms a antibonding molecular orbital. The

bonding molecular orbital has one node—a nodal plane that passes through both nu-clei. The antibonding molecular orbital has two nodal planes. Notice that bondsare cylindrically symmetrical,but bonds are not.

The extent of overlap is greater when porbitals overlap end-on than when theyoverlap side-to-side. This means that a bond formed by the end-on overlap of por-bitals is stronger than a bond formed by the side-to-side overlap of porbitals. It alsomeans that a bonding molecular orbital is more stable than a bonding molecularorbital because the stronger the bond,the more stable it is. Figure1.7shows a molec-ular orbital diagram of two identical atoms using their three degenerate atomic orbitalsto form three bonds—one bond and two bonds. p s

p sp

s

ps p*

pp*

p1P2 s

s*

A bond is stronger than a bond. P S



2p atomicorbital

2p atomicorbital

nodal plane

nodal plane

nodal plane

Energy

>Figure 1.6Side-to-side overlap of two parallelporbitals to form a bondingmolecular orbital and a antibonding molecular orbital.

p*p

BRUI01-001_059r4 20-03-2003 2:58 PM Page 23




PROBLEM 14!








C¬H

C¬H(CH4)

+

a.

+b.

+c.

+d.

1s, s*, p, or p*2

H

C





map for methane

H H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 25




PROBLEM 14!








C¬H

C¬H(CH4)

+

a.

+b.

+c.

+d.

1s, s*, p, or p*2

H

C





map for methane

H H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 25

. . . . Como fica o formato desta ligação σ ligante C-‐H?

Geometria angular

11

E o caso do carbono????




PROBLEM 14!








C¬H

C¬H(CH4)

+

a.

+b.

+c.

+d.

1s, s*, p, or p*2

H

C





map for methane

H H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 25




PROBLEM 14!








C¬H

C¬H(CH4)

+

a.

+b.

+c.

+d.

1s, s*, p, or p*2

H

C





map for methane

H H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 25




PROBLEM 14!








C¬H

C¬H(CH4)

+

a.

+b.

+c.

+d.

1s, s*, p, or p*2

H

C





map for methane

H H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 25

Molécula apolar EletronegaFvidade C e H

semelhantes

Todos os comprimentos de ligação iguais dC-‐H = 1,10

Carbono tetraédrico

12

......segundo a TOM????

Diacil de visualizar/pouco didáYco!!!!!

13


The potential map of methane shows that neither carbon nor hydrogen carries muchof a charge: There are neither red areas, representing partially negatively chargedatoms, nor blue areas, representing partially positively charged atoms. (Compare thismap with the potential map for water on p. 14). The absence of partially charged atomscan be explained by the similar electronegativities of carbon and hydrogen, whichcause carbon and hydrogen to share their bonding electrons relatively equally.Methane is a nonpolar molecule.

You may be surprised to learn that carbon forms four covalent bonds since youknow that carbon has only two unpaired electrons in its ground-state electronic con-figuration (Table 1.2). But if carbon were to form only two covalent bonds, it wouldnot complete its octet. Now we need to come up with an explanation that accounts forcarbon’s forming four covalent bonds.

If one of the electrons in the 2s orbital were promoted into the empty 2p atomic or-bital, the new electronic configuration would have four unpaired electrons; thus, fourcovalent bonds could be formed. Let’s now see whether this is feasible energetically.

Because a p orbital is higher in energy than an s orbital, promotion of an electronfrom an s orbital to a p orbital requires energy. The amount of energy required is

The formation of four bonds releases of energy be-cause the bond dissociation energy of a single bond is If theelectron were not promoted, carbon could form only two covalent bonds, which wouldrelease only So, by spending (or 402 kJ mol) to promotean electron, an extra (or 879 kJ mol) is released. In other words, promo-tion is energetically advantageous (Figure 1.9).

>210 kcal>mol>96 kcal>mol210 kcal>mol.

105 kcal>mol.C¬H420 kcal>molC¬H96 kcal>mol.

s

p p p

s

p p ppromotion

before promotion after promotion

Linus Carl Pauling (1901–1994)was born in Portland, Oregon. Afriend’s home chemistry laboratorysparked Pauling’s early interest inscience. He received a Ph.D. from theCalifornia Institute of Technologyand remained there for most of hisacademic career. He received theNobel Prize in chemistry in 1954 forhis work on molecular structure. LikeEinstein, Pauling was a pacifist, win-ning the 1964 Nobel Peace Prize forhis work on behalf of nucleardisarmament.

Pote

ntia

l ene

rgy

promotion96 kcal/mol

420 kcal/mol

4 covalentbonds

Figure 1.9 NAs a result of electron promotion,carbon forms four covalent bondsand releases 420 kcal mol ofenergy. Without promotion, carbonwould form two covalent bondsand release 210 kcal mol of energy.Because it requires 96 kcal mol topromote an electron, the overallenergy advantage of promotion is114 kcal>mol.

>>

>

We have managed to account for the observation that carbon forms four covalentbonds, but what accounts for the fact that the four bonds in methane are identi-cal? Each has a bond length of 1.10 and breaking any one of the bonds requires thesame amount of energy (105 kcal mol, or 439 kJ mol). If carbon used an s orbital andthree p orbitals to form these four bonds, the bond formed with the s orbital would bedifferent from the three bonds formed with p orbitals. How can carbon form four iden-tical bonds, using one s and three p orbitals? The answer is that carbon uses hybridorbitals.

Hybrid orbitals are mixed orbitals—they result from combining orbitals. The con-cept of combining orbitals, called orbital hybridization, was first proposed by LinusPauling in 1931. If the one s and three p orbitals of the second shell are combined andthen apportioned into four equal orbitals, each of the four resulting orbitals will be onepart s and three parts p. This type of mixed orbital is called an (stated “s-p-three”not “s-p-cubed”) orbital. (The superscript 3 means that three p orbitals were mixed

sp3

>>Å,

C¬H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 26

......usando hibridização para explicar a valência e geometria do carbono!!!!

April, 1931 THE NATURE OF THE CHEMICAL BOND 1367

[CONTRIBUTION FROM GATES CHEMICAL LABORATORY, CALIFORNIA INSTITUTE OF TECHNOLOGY, No. 2801

THE NATURE OF THE CHEMICAL BOND. APPLICATION OF RESULTS OBTAINED FROM THE

QUANTUM MECHANICS AND FROM A THEORY OF PARAMAGNETIC SUSCEPTIBILITY TO THE STRUCTURE

OF MOLECULES BY LINUS PAULING

RECEIVED FEBRUARY 17,1931 PUBLISHED APRIL 6, 1931

During the last four years the problem of the nature of the chemical bond has been attacked by theoretical physicists, especially Heitler and London, by the application of the quantum mechanics. This work has led to an approximate theoretical calculation of the energy of formation and of other properties of very simple molecules, such as Hz, and has also pro- vided a formal justification of the rules set up in 1916 by G. N. Lewis for his electron-pair bond. In the following paper it will be shown that many more results of chemical significance can be obtained from the quantum mechanical equations, permitting the formulation of an extensive and powerful set of rules for the electron-pair bond supplementing those of Lewis. These rules provide information regarding the relative strengths of bonds formed by different atoms, the angles between bonds, free rotation or lack of free rotation about bond axes, the relation between the quantum numbers of bonding electrons and the number and spatial arrangement of the bonds, etc. A complete theory of the magnetic moments of molecules and complex ions is also developed, and it is shown that for many compounds involving elements of the transition groups this theory together with the rules for electron-pair bonds leads to a unique assignment of electron structures as well as a definite determination of the type of bonds involved.'

I. The Electron-Pair Bond The Interaction of Simple Atoms.-The discussion of the wave equation

for the hydrogen molecule by Heitler and London,2 S ~ g i u r a , ~ and Wang4 showed that two normal hydrogen atoms can interact in either of two ways, one of which gives rise to repulsion with no molecule formation, the other

A preliminary announcement of some of these results was made three years ago [Linus Pauling, Proc. Nat. Acad. Sci., 14, 359 (1928)l. Two of the results (90' bond angles for p eigenfunctions, and the existence, but not the stability, of tetrahedral eigenfunctions) have been independently discovered by Professor J. C. Slater and an- nounced at meetings of the National Academy of Sciences (Washington, April, 1930) and the American Physical Society (Cleveland, December, 1930).

W. Heitler and F. London, 2. Physik, 44, 455 (1927). Y. Sugiura, ibid. , 45, 484 (1927). S. C. Wang, Phys. Rev., 31, 579 (1928).

J. Am. Chem. Soc., 1931, 53 (4), pp 1367–1400

14

When assigning electrons to MOs, the same rules apply as for writing electron con-figurations of atoms. Electrons fill the MOs in order of increasing orbital energy, and themaximum number of electrons in any orbital is 2. The 2 electrons of H2 occupy thebonding orbital, have opposite spins, and both are held more strongly than they wouldbe in separated hydrogen atoms. There are no electrons in the antibonding orbital.

For a molecule as simple as H2, it is hard to see much difference between thevalence bond and molecular orbital methods. The most important differences appear inmolecules with more than two atoms—a very common situation indeed. In those cases,the valence bond method continues to view a molecule as a collection of bonds betweenconnected atoms. The molecular orbital method, however, leads to a picture in which thesame electron can be associated with many, or even all, of the atoms in a molecule.

In the remaining sections of this chapter we will use a modification of valencebond theory to describe CH and CC bonds in some fundamental types of organic com-pounds.

1.15 BONDING IN METHANE AND ORBITAL HYBRIDIZATION

A vexing puzzle in the early days of valence bond theory concerned the bonding inmethane (CH4). Since covalent bonding requires the overlap of half-filled orbitals of theconnected atoms, carbon with an electron configuration of 1s22s22px

12py1 has only two

half-filled orbitals (Figure 1.20a), so how can it have bonds to four hydrogens?

1.15 Bonding in Methane and Orbital Hybridization 35

Incr

easi

ng e

nerg

y

Antibonding

Bonding

1s 1s

Hydrogen 1satomic orbital

Molecular orbitalsof H2


FIGURE 1.19 Two molecu-lar orbitals are generated bycombining two hydrogen 1sorbitals. One molecularorbital is a bonding molecu-lar orbital and is lower inenergy than either of theatomic orbitals that combineto produce it. The othermolecular orbital is anti-bonding and is of higherenergy than either atomicorbital. Each arrow indicatesone electron; the electronspins are opposite in sign.The bonding orbital containsboth electrons of H2.

2p

2s

2p

2s

2sp3

Ene

rgy

Ground electronicstate of carbon

Higher energy electronicstate of carbon

sp3 hybridstate of carbon

(c)(b)(a)

FIGURE 1.20 (a) Electronconfiguration of carbon inits most stable state. (b) Anelectron is “promoted” fromthe 2s orbital to the vacant2p orbital. (c) The 2s orbitaland the three 2p orbitals arecombined to give a set offour equal-energy sp3-hybridized orbitals, each ofwhich contains one electron.

Hibridização ......usando hibridização para explicar a valência e geometria do metano!!!!

promoção

Neste caso, os orbitais não apontam para os cantos de um tetraedro.






12py1 has only two



Incr

easi

ng e

nerg

y

Antibonding

Bonding

1s 1s





2p

2s

2p

2s

2sp3

Ene

rgy




(c)(b)(a)


In the 1930s Linus Pauling offered an ingenious solution to the puzzle. He beganwith a simple idea: “promoting” one of the 2s electrons to the empty 2pz orbital givesfour half-filled orbitals and allows for four C±H bonds (Figure 1.20b). The electronconfiguration that results (1s22s12px

12py12pz

1), however, is inconsistent with the fact thatall of these bonds are equivalent and directed toward the corners of a tetrahedron. Thesecond part of Pauling’s idea was novel: mix together (hybridize) the four valenceorbitals of carbon (2s, 2px, 2py, and 2pz) to give four half-filled orbitals of equal energy(Figure 1.20c). The four new orbitals in Pauling’s scheme are called sp3 hybrid orbitalsbecause they come from one s orbital and three p orbitals.

Figure 1.21 depicts some of the spatial aspects of orbital hybridization. Each sp3

hybrid orbital has two lobes of unequal size, making the electron density greater on oneside of the nucleus than the other. In a bond to hydrogen, it is the larger lobe of a car-bon sp3 orbital that overlaps with a hydrogen 1s orbital. The orbital overlaps corre-sponding to the four C±H bonds of methane are portrayed in Figure 1.22. Orbital over-lap along the internuclear axis generates a bond with rotational symmetry—in this casea C(2sp3)±H(1s) ! bond. A tetrahedral arrangement of four ! bonds is characteristicof sp3-hybridized carbon.

The peculiar shape of sp3 hybrid orbitals turn out to have an important consequence.Since most of the electron density in an sp3 hybrid orbital lies to one side of a carbonatom, overlap with a half-filled 1s orbital of hydrogen, for example, on that side producesa stronger bond than would result otherwise. If the electron probabilities were equal onboth sides of the nucleus, as it would be in a p orbital, half of the time the electron wouldbe remote from the region between the bonded atoms, and the bond would be weaker.Thus, not only does Pauling’s orbital hybridization proposal account for carbon formingfour bonds rather than two, these bonds are also stronger than they would be otherwise.

36 CHAPTER ONE Chemical Bonding

2s

sp3

Combine one 2s and three 2p orbitals to give four equivalent sp3 hybrid orbitals:

x x x x

y

z z z z

y y y

!

!

!""

2px 2py 2pz

The two lobes of each sp3

hybrid orbital are ofdifferent size. More ofthe electron density isconcentrated on one side of the nucleus than on the other.

"

"

"!4

FIGURE 1.21 Representa-tion of orbital mixing in sp3

hybridization. Mixing of ones orbital with three porbitals generates four sp3

hybrid orbitals. Each sp3

hybrid orbital has 25% scharacter and 75% p charac-ter. The four sp3 hybridorbitals have their majorlobes directed toward thecorners of a tetrahedron,which has the carbon atomat its center.






12py1 has only two



Incr

easi

ng e

nerg

y

Antibonding

Bonding

1s 1s





2p

2s

2p

2s

2sp3

Ene

rgy




(c)(b)(a)


15

......orbitais híbridos sp3: 4 orbitais atômicos geram 4 orbitais híbridos.


s orbital

the s orbital adds tothe lobe of the p orbital

p orbital the s orbital subtracts fromthe lobe of the p orbital

> Figure 1.10The s orbital adds to one lobe ofthe p orbital and subtracts from theother lobe of the p orbital.

p

s

p psp3 sp3 sp3 sp3

hybridization

H

H

b.a.

H

H

CC

> Figure 1.12(a) The four orbitals aredirected toward the corners of atetrahedron, causing each bondangle to be 109.5°.(b) An orbital picture of methane,showing the overlap of each orbital of carbon with the sorbital of a hydrogen. (For clarity,the smaller lobes of the orbitals are not shown.)

sp3

sp3

sp3

with one s orbital to form the hybrid orbitals.) Each orbital has 25% s characterand 75% p character. The four orbitals are degenerate—they have the same energy.

Like a p orbital, an orbital has two lobes. The lobes differ in size, however, be-cause the s orbital adds to one lobe of the p orbital and subtracts from the other lobe ofthe p orbital (Figure 1.10). The stability of an orbital reflects its composition; it ismore stable than a p orbital, but not as stable as an s orbital (Figure 1.11). The largerlobe of the orbital is used in covalent bond formation.sp3

sp3

sp3

p p p

ssp3 sp3 sp3sp3hybridization

4 orbitals are hybridizedhybrid orbitals

sp3sp3

> Figure 1.11An s orbital and three p orbitalshybridize to form four orbitals.An orbital is morestable than a p orbital, but not asstable as an s orbital.

sp3sp3

The four orbitals arrange themselves in space in a way that allows them to getas far away from each other as possible (Figure 1.12a). This occurs because electronsrepel each other and getting as far from each other as possible minimizes the repulsion(Section 1.6). When four orbitals spread themselves into space as far from each otheras possible, they point toward the corners of a regular tetrahedron (a pyramid with four

sp3

BRUI01-001_059r4 20-03-2003 2:58 PM Page 27


s orbital




p

s

p psp3 sp3 sp3 sp3

hybridization

H

H

b.a.

H

H

CC


sp3

sp3

sp3



sp3

sp3

p p p



sp3sp3


sp3sp3


sp3

BRUI01-001_059r4 20-03-2003 2:58 PM Page 27

16


s orbital




p

s

p psp3 sp3 sp3 sp3

hybridization

H

H

b.a.

H

H

CC


sp3

sp3

sp3



sp3

sp3

p p p



sp3sp3


sp3sp3


sp3

BRUI01-001_059r4 20-03-2003 2:58 PM Page 27


s orbital




p

s

p psp3 sp3 sp3 sp3

hybridization

H

H

b.a.

H

H

CC


sp3

sp3

sp3



sp3

sp3

p p p



sp3sp3


sp3sp3


sp3

BRUI01-001_059r4 20-03-2003 2:58 PM Page 27

Distribuição espacial dos orbitais PROBLEM 1.20 Construct an orbital diagram like that of Figure 1.20 for nitro-gen in ammonia, assuming sp3 hybridization. In what kind of orbital is theunshared pair? What orbital overlaps are involved in the N±H bonds?

1.16 sp3 HYBRIDIZATION AND BONDING IN ETHANE

The orbital hybridization model of covalent bonding is readily extended to carbon–carbon bonds. As Figure 1.23 illustrates, ethane is described in terms of a carbon–carbon ! bond joining two CH3 (methyl) groups. Each methyl group consists of an sp3-hybridized carbon attached to three hydrogens by sp3–1s ! bonds. Overlap of theremaining half-filled orbital of one carbon with that of the other generates a ! bondbetween them. Here is a third kind of ! bond, one that has as its basis the overlap oftwo sp3-hybridized orbitals. In general, you can expect that carbon will be sp3-hybridizedwhen it is directly bonded to four atoms.

PROBLEM 1.21 Describe the bonding in methylsilane (H3CSiH3), assuming thatit is analogous to that of ethane. What is the principal quantum number of theorbitals of silicon that are hybridized?

The orbital hybridization model of bonding is not limited to compounds in whichall the bonds are single, but can be adapted to compounds with double and triple bonds,as described in the following two sections.

1.16 sp3 Hybridization and Bonding in Ethane 37

H

HH

H

CComing toward you

Going awayfrom you

In the planeof the paper

In the planeof the paper

109.5!

H(1s)±C(2sp3)" bond

FIGURE 1.22 The sp3 hy-brid orbitals are arranged in a tetrahedral fashionaround carbon. Each orbitalcontains one electron andcan form a bond with ahydrogen atom to give atetrahedral methane mole-cule. (Note: Only the majorlobe of each sp3 orbital isshown. As indicated in Fig-ure 1.21, each orbital con-tains a smaller back lobe,which has been omitted forthe sake of clarity.)

FIGURE 1.23 Orbital over-lap description of thesp3–sp3 ! bond between thetwo carbon atoms ofethane.

The C±H and C±C bonddistances in ethane are 111and 153 pm, respectively,and the bond angles areclose to tetrahedral.

4 ligações σ iguais.

Você consegue desenhar esta estrutura?




PROBLEM 14!








C¬H

C¬H(CH4)

+

a.

+b.

+c.

+d.

1s, s*, p, or p*2

H

C





map for methane

H H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 25

17




PROBLEM 14!








C¬H

C¬H(CH4)

+

a.

+b.

+c.

+d.

1s, s*, p, or p*2

H

C





map for methane

H H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 25




PROBLEM 14!








C¬H

C¬H(CH4)

+

a.

+b.

+c.

+d.

1s, s*, p, or p*2

H

C





map for methane

H H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 25

. .

. .


The MO diagram shows that the bonding molecular orbital is more stable—is lowerin energy—than the individual atomic orbitals. This is because the more nuclei anelectron “feels,” the more stable it is. The antibonding molecular orbital, with lesselectron density between the nuclei, is less stable—is of higher energy—than theatomic orbitals.

After the MO diagram is constructed, the electrons are assigned to the molecularorbitals. The aufbau principle and the Pauli exclusion principle, which apply to elec-trons in atomic orbitals, also apply to electrons in molecular orbitals: Electrons alwaysoccupy available orbitals with the lowest energy, and no more than two electrons canoccupy a molecular orbital. Thus, the two electrons of the bond occupy thelower energy bonding molecular orbital (Figure 1.4), where they are attracted to bothpositively charged nuclei. It is this electrostatic attraction that gives a covalent bond itsstrength. Therefore, the greater the overlap of the atomic orbitals, the stronger is thecovalent bond. The strongest covalent bonds are formed by electrons that occupy themolecular orbitals with the lowest energy.

The MO diagram in Figure 1.4 allows us to predict that would not be as stableas because has only one electron in the bonding orbital. We can also predictthat does not exist: Because each He atom would bring two electrons, wouldhave four electrons—two filling the lower energy bonding molecular orbital and theremaining two filling the higher energy antibonding molecular orbital. The two elec-trons in the antibonding molecular orbital would cancel the advantage to bondinggained by the two electrons in the bonding molecular orbital.

PROBLEM 13!

Predict whether or not exists.

Two p atomic orbitals can overlap either end-on or side-to-side. Let’s first look atend-on overlap. End-on overlap forms a bond. If the overlapping lobes of the p or-bitals are in-phase (a blue lobe of one p orbital overlaps a blue lobe of the other p or-bital), a bonding molecular orbital is formed (Figure 1.5). The electron density ofthe bonding molecular orbital is concentrated between the nuclei, which causes theback lobes (the nonoverlapping lobes) of the molecular orbital to be quite small. The bonding molecular orbital has two nodes—a nodal plane passing through each of thenuclei.

If the overlapping lobes of the p orbitals are out-of-phase (a blue lobe of one p or-bital overlaps a green lobe of the other p orbital), a antibonding molecular orbital iss*

sss

s

He2

+

He2He2

H2

+H2

H2

+

H—H

When two atomic orbitals overlap, twomolecular orbitals are formed—onelower in energy and one higher in ener-gy than the atomic orbitals.

In-phase overlap forms a bonding MO;out-of-phase overlap forms an anti-bonding MO.



1s atomicorbital

1s atomicorbital

Ener

gy

nodeFigure 1.4 NAtomic orbitals of and molecularorbitals of Before covalent bondformation, each electron is in anatomic orbital. After covalent bondformation, both electrons are in thebonding molecular orbital. Theantibonding molecular orbital isempty.

H2.H–

BRUI01-001_059r4 20-03-2003 2:58 PM Page 22




PROBLEM 14!








C¬H

C¬H(CH4)

+

a.

+b.

+c.

+d.

1s, s*, p, or p*2

H

C





map for methane

H H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 25

. Orbital σ*


s orbital




p

s

p psp3 sp3 sp3 sp3

hybridization

H

H

b.a.

H

H

CC


sp3

sp3

sp3



sp3

sp3

p p p



sp3sp3


sp3sp3


sp3

BRUI01-001_059r4 20-03-2003 2:58 PM Page 27

Os orbitais σ e σ* encontram-‐se centrados neste eixo. Para cada ligação C-‐H, temos um par de orbitais σ e σ*.

Orbital σ

18


faces, each an equilateral triangle). Each of the four bonds in methane isformed from overlap of an orbital of carbon with the s orbital of a hydrogen(Figure 1.12b). This explains why the four bonds are identical.

The angle formed between any two bonds of methane is 109.5°. This bond angle iscalled the tetrahedral bond angle. A carbon, such as the one in methane, that formscovalent bonds using four equivalent orbitals is called a tetrahedral carbon.

The postulation of hybrid orbitals may appear to be a theory contrived just to makethings fit—and that is exactly what it is. Nevertheless, it is a theory that gives us a verygood picture of the bonding in organic compounds.

Note to the student

It is important to understand what molecules look like in three dimensions. As you studyeach chapter, make sure to visit the Web site www.prenhall.com/bruice and look at thethree-dimensional representations of molecules that can be found in the molecule gallerythat accompanies the chapter.

Bonding in EthaneThe two carbon atoms in ethane are tetrahedral. Each carbon uses four orbitals toform four covalent bonds:

One orbital of one carbon overlaps an orbital of the other carbon to formthe bond. Each of the remaining three orbitals of each carbon overlapsthe s orbital of a hydrogen to form a bond. Thus, the bond isformed by overlap, and each bond is formed by overlap(Figure 1.13). Each of the bond angles in ethane is nearly the tetrahedral bond angleof 109.5°, and the length of the bond is 1.54 Ethane, like methane, is anonpolar molecule.

Å.C¬C

sp3–sC¬Hsp3–sp3C¬CC¬H

sp3C¬Csp3sp3

CH

H

H

C

H

H

H

ethane

sp3

sp3

C¬Hsp3

C¬H

H H

HH

H H H H

H H

H H

C C C C

Electron pairs spread themselves intospace as far from each other as possible.

! Figure 1.13An orbital picture of ethane. The bond is formed by overlap, and each bond is formed by overlap. (The smaller lobes of the orbitals are not shown.)sp3sp3–s

C¬ Hsp3–sp3C¬ C

perspective formulaof ethane

ball-and-stick model of ethane

°1.54 A

°1.10 A 109.6°

space-filling modelof ethane

electrostatic potentialmap for ethane

CC

H

H

H

H

HH

3-D Molecule:Methane

BRUI01-001_059r4 20-03-2003 2:58 PM Page 28





Note to the student




Å.C¬C


sp3C¬Csp3sp3

CH

H

H

C

H

H

H

ethane

sp3

sp3

C¬Hsp3

C¬H

H H

HH

H H H H

H H

H H

C C C C



C¬ Hsp3–sp3C¬ C



°1.54 A

°1.10 A 109.6°



CC

H

H

H

H

HH


BRUI01-001_059r4 20-03-2003 2:58 PM Page 28





Note to the student




Å.C¬C


sp3C¬Csp3sp3

CH

H

H

C

H

H

H

ethane

sp3

sp3

C¬Hsp3

C¬H

H H

HH

H H H H

H H

H H

C C C C



C¬ Hsp3–sp3C¬ C



°1.54 A

°1.10 A 109.6°



CC

H

H

H

H

HH


BRUI01-001_059r4 20-03-2003 2:58 PM Page 28





Note to the student




Å.C¬C


sp3C¬Csp3sp3

CH

H

H

C

H

H

H

ethane

sp3

sp3

C¬Hsp3

C¬H

H H

HH

H H H H

H H

H H

C C C C



C¬ Hsp3–sp3C¬ C



°1.54 A

°1.10 A 109.6°



CC

H

H

H

H

HH


BRUI01-001_059r4 20-03-2003 2:58 PM Page 28





Note to the student




Å.C¬C


sp3C¬Csp3sp3

CH

H

H

C

H

H

H

ethane

sp3

sp3

C¬Hsp3

C¬H

H H

HH

H H H H

H H

H H

C C C C



C¬ Hsp3–sp3C¬ C



°1.54 A

°1.10 A 109.6°



CC

H

H

H

H

HH


BRUI01-001_059r4 20-03-2003 2:58 PM Page 28





Note to the student




Å.C¬C


sp3C¬Csp3sp3

CH

H

H

C

H

H

H

ethane

sp3

sp3

C¬Hsp3

C¬H

H H

HH

H H H H

H H

H H

C C C C



C¬ Hsp3–sp3C¬ C



°1.54 A

°1.10 A 109.6°



CC

H

H

H

H

HH


BRUI01-001_059r4 20-03-2003 2:58 PM Page 28





Note to the student




Å.C¬C


sp3C¬Csp3sp3

CH

H

H

C

H

H

H

ethane

sp3

sp3

C¬Hsp3

C¬H

H H

HH

H H H H

H H

H H

C C C C



C¬ Hsp3–sp3C¬ C



°1.54 A

°1.10 A 109.6°



CC

H

H

H

H

HH


BRUI01-001_059r4 20-03-2003 2:58 PM Page 28

Você consegue desenhar esta estrutura?

σ (C(sp3)-‐H(1s))

σ (C(sp3)-‐C(sp3))

Por que ligação C-‐C é maior que C-‐H?

19

All the bonds in methane and ethane are sigma bonds because they are allformed by the end-on overlap of atomic orbitals. All single bonds found in organiccompounds are sigma bonds.

PROBLEM 15!

What orbitals are used to form the 10 covalent bonds in propane

The MO diagram illustrating the overlap of an orbital of one carbon with an orbital of another carbon (Figure 1.14) is similar to the MO diagram for the end-onoverlap of two p orbitals, which should not be surprising since orbitals have 75%p character.

sp3

sp3sp3

(CH3CH2CH3)?

(s)

Section 1.8 Bonding in Ethene: A Double Bond 29



sp3 atomicorbital

sp3 atomicorbital

Ener

gy

> Figure 1.14End-on overlap of two orbitalsto form a bonding molecularorbital and a antibondingmolecular orbital.

s*s

sp3

All single bonds found in organic com-pounds are sigma bonds.

3-D Molecule:Ethane

1.8 Bonding in Ethene: A Double Bond

Each of the carbon atoms in ethene (also called ethylene) forms four bonds, but each isbonded to only three atoms:

To bond to three atoms, each carbon hybridizes three atomic orbitals. Because threeorbitals (an s orbital and two of the p orbitals) are hybridized, three hybrid orbitals areobtained. These are called orbitals. After hybridization, each carbon atom hasthree degenerate orbitals and one p orbital:

To minimize electron repulsion, the three orbitals need to get as far from eachother as possible. Therefore, the axes of the three orbitals lie in a plane, directedtoward the corners of an equilateral triangle with the carbon nucleus at the center. Thismeans that the bond angles are all close to 120°. Because the hybridized carbonsp2

sp2

p p p

ssp2 sp2

psp2

hybridization

three orbitals are hybridizedhybrid orbitals

sp2sp2

C C

Hethene

(ethylene)

H

H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 29





Note to the student




Å.C¬C


sp3C¬Csp3sp3

CH

H

H

C

H

H

H

ethane

sp3

sp3

C¬Hsp3

C¬H

H H

HH

H H H H

H H

H H

C C C C



C¬ Hsp3–sp3C¬ C



°1.54 A

°1.10 A 109.6°



CC

H

H

H

H

HH


BRUI01-001_059r4 20-03-2003 2:58 PM Page 28

Os orbitais σ e σ* encontram-‐se centrados neste eixo.

.

. .

. .

.

20

A Dupla ligação ......usando hibridização para explicar a geometria do eteno.

All the bonds in methane and ethane are sigma bonds because they are allformed by the end-on overlap of atomic orbitals. All single bonds found in organiccompounds are sigma bonds.

PROBLEM 15!

What orbitals are used to form the 10 covalent bonds in propane

The MO diagram illustrating the overlap of an orbital of one carbon with an orbital of another carbon (Figure 1.14) is similar to the MO diagram for the end-onoverlap of two p orbitals, which should not be surprising since orbitals have 75%p character.

sp3

sp3sp3

(CH3CH2CH3)?

(s)

Section 1.8 Bonding in Ethene: A Double Bond 29



sp3 atomicorbital

sp3 atomicorbital

Ener

gy

> Figure 1.14End-on overlap of two orbitalsto form a bonding molecularorbital and a antibondingmolecular orbital.

s*s

sp3

All single bonds found in organic com-pounds are sigma bonds.

3-D Molecule:Ethane

1.8 Bonding in Ethene: A Double Bond

Each of the carbon atoms in ethene (also called ethylene) forms four bonds, but each isbonded to only three atoms:

To bond to three atoms, each carbon hybridizes three atomic orbitals. Because threeorbitals (an s orbital and two of the p orbitals) are hybridized, three hybrid orbitals areobtained. These are called orbitals. After hybridization, each carbon atom hasthree degenerate orbitals and one p orbital:

To minimize electron repulsion, the three orbitals need to get as far from eachother as possible. Therefore, the axes of the three orbitals lie in a plane, directedtoward the corners of an equilateral triangle with the carbon nucleus at the center. Thismeans that the bond angles are all close to 120°. Because the hybridized carbonsp2

sp2

p p p

ssp2 sp2

psp2

hybridization

three orbitals are hybridizedhybrid orbitals

sp2sp2

C C

Hethene

(ethylene)

H

H

H

BRUI01-001_059r4 20-03-2003 2:58 PM Page 29

1.17 sp2 HYBRIDIZATION AND BONDING IN ETHYLENE

Ethylene is a planar molecule, as the structural representations of Figure 1.24 indi-cate. Because sp3 hybridization is associated with a tetrahedral geometry at carbon,it is not appropriate for ethylene, which has a trigonal planar geometry at both of itscarbons. The hybridization scheme is determined by the number of atoms to whichthe carbon is directly attached. In ethane, four atoms are attached to carbon by !bonds, and so four equivalent sp3 hybrid orbitals are required. In ethylene, three atomsare attached to each carbon, so three equivalent hybrid orbitals are required. As shownin Figure 1.25, these three orbitals are generated by mixing the carbon 2s orbital withtwo of the 2p orbitals and are called sp2 hybrid orbitals. One of the 2p orbitals isleft unhybridized.


2s 2s

2sp2

Ene

rgy




(c)(b)(a)

2p 2p 2p

FIGURE 1.25 (a) Electron configuration of carbon in its most stable state. (b) An electron is“promoted” from the 2s orbital to the vacant 2p orbital. (c) The 2s orbital and two of the three2p orbitals are combined to give a set of three equal-energy sp2-hybridized orbitals. One ofthe 2p orbitals remains unchanged.

Another name for ethyleneis ethene.

134 pm

121.4!

110 pmCœC

H

H

H

H

(a)

117.2!

(b)

FIGURE 1.24 (a) Allthe atoms of ethylene lie inthe same plane. All the bondangles are close to 120°, andthe carbon–carbon bond dis-tance is significantly shorterthan that of ethane. (b) Aspace-filling model of ethyl-ene.

Aqui temos a hibridização dos orbitais: 2s + 2px + 2py

21

Figure 1.26 illustrates the mixing of orbitals in sp2 hybridization. The three sp2

orbitals are of equal energy; each has one-third s character and two-thirds p character.Their axes are coplanar, and each has a shape much like that of an sp3 orbital.

Each carbon of ethylene uses two of its sp2 hybrid orbitals to form ! bonds to twohydrogen atoms, as illustrated in the first part of Figure 1.27. The remaining sp2 orbitals,one on each carbon, overlap along the internuclear axis to give a ! bond connecting thetwo carbons.

As Figure 1.27 shows, each carbon atom still has, at this point, an unhybridized2p orbital available for bonding. These two half-filled 2p orbitals have their axes per-pendicular to the framework of ! bonds of the molecule and overlap in a side-by-sidemanner to give what is called a pi (!) bond. According to this analysis, the carbon–car-bon double bond of ethylene is viewed as a combination of a ! bond plus a " bond.The additional increment of bonding makes a carbon–carbon double bond both strongerand shorter than a carbon–carbon single bond.

Electrons in a " bond are called ! electrons. The probability of finding a " elec-tron is highest in the region above and below the plane of the molecule. The plane ofthe molecule corresponds to a nodal plane, where the probability of finding a " electronis zero.

In general, you can expect that carbon will be sp2-hybridized when it is directlybonded to three atoms.

1.17 sp2 Hybridization and Bonding in Ethylene 39

x

z

y

x

z

y

x

z

y

Combine one 2s and two 2p orbitals

x

z

y

Leave thisorbital alone

3!

Three sp2 hybrid orbitals

x

z

y

2pz

FIGURE 1.26 Representation of orbital mixing in sp2 hybridization. Mixing of one s orbitalwith two p orbitals generates three sp2 hybrid orbitals. Each sp2 hybrid orbital has one-third scharacter and two-thirds p character. The axes of the three sp2 hybrid orbitals are coplanar.One 2p orbital remains unhybridized, and its axis is perpendicular to the plane defined by theaxes of the sp2 orbitals.

One measure of the strengthof a bond is its bond dissoci-ation energy. This topic willbe introduced in Section 4.17and applied to ethylene inSection 5.2.

22

1.18 sp HYBRIDIZATION AND BONDING IN ACETYLENE

One more hybridization scheme is important in organic chemistry. It is called sphybridization and applies when carbon is directly bonded to two atoms, as it is in acety-lene. The structure of acetylene is shown in Figure 1.28 along with its bond distancesand bond angles.

Since each carbon in acetylene is bonded to two other atoms, the orbital hybridiza-tion model requires each carbon to have two equivalent orbitals available for the for-mation of ! bonds as outlined in Figures 1.29 and 1.30. According to this model the car-bon 2s orbital and one of the 2p orbitals combine to generate a pair of two equivalentsp hybrid orbitals. Each sp hybrid orbital has 50% s character and 50% p character. Thesetwo sp orbitals share a common axis, but their major lobes are oriented at an angle of180° to each other. Two of the original 2p orbitals remain unhybridized. Their axes areperpendicular to each other and to the common axis of the pair of sp hybrid orbitals.


Another name for acetyleneis ethyne.

Begin with two sp2 hybridized carbon atoms and four hydrogen atoms:

sp2

sp2

H

sp2sp2

sp2

sp2

sp2 hybrid orbitals of carbonoverlap to form ! bonds tohydrogens and to each otherC(2sp2) –H(1s)

! bond

C(2sp2) –C(2sp2) ! bond

C(2p) –C(2p) " bond

p orbitals that remain on carbonsoverlap to form " bond

Half-filled 2porbital

In plane ofpaper

HH

H

FIGURE 1.27 The carbon–carbon double bond in ethyl-ene has a ! component and a" component. The ! compo-nent arises from overlap ofsp2-hybridized orbitals alongthe internuclear axis. The "component results from aside-by-side overlap of 2porbitals.

1.18 sp HYBRIDIZATION AND BONDING IN ACETYLENE

One more hybridization scheme is important in organic chemistry. It is called sphybridization and applies when carbon is directly bonded to two atoms, as it is in acety-lene. The structure of acetylene is shown in Figure 1.28 along with its bond distancesand bond angles.

Since each carbon in acetylene is bonded to two other atoms, the orbital hybridiza-tion model requires each carbon to have two equivalent orbitals available for the for-mation of ! bonds as outlined in Figures 1.29 and 1.30. According to this model the car-bon 2s orbital and one of the 2p orbitals combine to generate a pair of two equivalentsp hybrid orbitals. Each sp hybrid orbital has 50% s character and 50% p character. Thesetwo sp orbitals share a common axis, but their major lobes are oriented at an angle of180° to each other. Two of the original 2p orbitals remain unhybridized. Their axes areperpendicular to each other and to the common axis of the pair of sp hybrid orbitals.


Another name for acetyleneis ethyne.

Begin with two sp2 hybridized carbon atoms and four hydrogen atoms:

sp2

sp2

H

sp2sp2

sp2

sp2

sp2 hybrid orbitals of carbonoverlap to form ! bonds tohydrogens and to each otherC(2sp2) –H(1s)

! bond

C(2sp2) –C(2sp2) ! bond

C(2p) –C(2p) " bond

p orbitals that remain on carbonsoverlap to form " bond

Half-filled 2porbital

In plane ofpaper

HH

H

FIGURE 1.27 The carbon–carbon double bond in ethyl-ene has a ! component and a" component. The ! compo-nent arises from overlap ofsp2-hybridized orbitals alongthe internuclear axis. The "component results from aside-by-side overlap of 2porbitals.

23


HH

H HH

! bond ! bond

" bond

" bond

a. b. c.

H

H

H

HC C CC

" bond formedby sp2– s overlap

" bond formed bysp2– sp2 overlap

C C

H H

H

atom is bonded to three atoms that define a plane, it is called a trigonal planar carbon.The unhybridized p orbital is perpendicular to the plane defined by the axes of the orbitals (Figure 1.15).

The carbons in ethene form two bonds with each other. This is called a doublebond. The two carbon–carbon bonds in the double bond are not identical. One of thebonds results from the overlap of an orbital of one carbon with an orbital of theother carbon; this is a sigma bond because it is formed by end-on overlap(Figure 1.16a). Each carbon uses its other two orbitals to overlap the s orbital of ahydrogen to form the bonds. The second carbon–carbon bond results fromside-to-side overlap of the two unhybridized p orbitals. Side-to-side overlap of p or-bitals forms a pi bond (Figure 1.16b). Thus, one of the bonds in a double bond is a

bond and the other is a bond. All the bonds are bonds.sC¬Hps(p)

C¬Hsp2

(s)sp2sp2

sp2

side view top view

p

120°sp2

sp2sp2

Figure 1.15 NAn hybridized carbon. The threedegenerate orbitals lie in aplane. The unhybridized p orbital isperpendicular to the plane. (Thesmaller lobes of the orbitals arenot shown.)

sp2

sp2sp2

! Figure 1.16(a) One bond in ethene is a bond formed by overlap, and the bondsare formed by overlap. (b) The second bond is a bond formed by side-to-sideoverlap of a p orbital of one carbon with a p orbital of the other carbon. (c) There is anaccumulation of electron density above and below the plane containing the two carbonsand four hydrogens.

pC¬ Csp2–sC¬ Hsp2–sp2sC¬ C

3-D Molecule:Ethene

The two p orbitals that overlap to form the bond must be parallel to each otherfor maximum overlap to occur. This forces the triangle formed by one carbon andtwo hydrogens to lie in the same plane as the triangle formed by the other carbonand two hydrogens. This means that all six atoms of ethene lie in the same plane,and the electrons in the p orbitals occupy a volume of space above and below theplane (Figure 1.16c). The electrostatic potential map for ethene shows that it is anonpolar molecule with an accumulation of negative charge (the orange area) abovethe two carbons. (If you could turn the potential map over, a similar accumulation ofnegative charge would be found on the other side.)

p

1.33 A°

1.08 A°

121.7°

116.6°C C

H

H

H

Ha double bond consists of

one " bond and one ! bondball-and-stick model

of ethenespace-filling model

of etheneelectrostatic potential map

for ethene

BRUI01-001_059r4 20-03-2003 2:58 PM Page 30


HH

H HH

! bond ! bond

" bond

" bond

a. b. c.

H

H

H

HC C CC



C C

H H

H




C¬Hsp2

(s)sp2sp2

sp2

side view top view

p

120°sp2

sp2sp2


sp2

sp2sp2



3-D Molecule:Ethene


p

1.33 A°

1.08 A°

121.7°

116.6°C C

H

H

H





for ethene

BRUI01-001_059r4 20-03-2003 2:58 PM Page 30


HH

H HH

! bond ! bond

" bond

" bond

a. b. c.

H

H

H

HC C CC



C C

H H

H




C¬Hsp2

(s)sp2sp2

sp2

side view top view

p

120°sp2

sp2sp2


sp2

sp2sp2



3-D Molecule:Ethene


p

1.33 A°

1.08 A°

121.7°

116.6°C C

H

H

H





for ethene

BRUI01-001_059r4 20-03-2003 2:58 PM Page 30

24

Você consegue desenhar estas estruturas?

Desenhe o eteno em ao menos 3 perspec-vas

diferentes!!!!

Agora escolha uma que vc consiga desenhar todos os orbitais envolvidos nas ligações C-‐H e C=C.

1)

2)

25


nodes

nodenode



2p atomicorbital

2p atomicorbital

Ener

gy


s*s


S

P





psp

s

psp*

pp*

p1P2s

s*




2p atomicorbital

2p atomicorbital

nodal plane

nodal plane

nodal plane

Ener

gy> Figure 1.6Side-to-side overlap of two parallelp orbitals to form a bondingmolecular orbital and a antibonding molecular orbital.

p*p

BRUI01-001_059r4 20-03-2003 2:58 PM Page 23

26


HH

H HH

! bond ! bond

" bond

" bond

a. b. c.

H

H

H

HC C CC



C C

H H

H




C¬Hsp2

(s)sp2sp2

sp2

side view top view

p

120°sp2

sp2sp2


sp2

sp2sp2



3-D Molecule:Ethene


p

1.33 A°

1.08 A°

121.7°

116.6°C C

H

H

H





for ethene

BRUI01-001_059r4 20-03-2003 2:58 PM Page 30





Note to the student




Å.C¬C


sp3C¬Csp3sp3

CH

H

H

C

H

H

H

ethane

sp3

sp3

C¬Hsp3

C¬H

H H

HH

H H H H

H H

H H

C C C C



C¬ Hsp3–sp3C¬ C



°1.54 A

°1.10 A 109.6°



CC

H

H

H

H

HH


BRUI01-001_059r4 20-03-2003 2:58 PM Page 28

Compare os comprimentos de ligação (C-‐H) e (CC), explicando as diferenças observadas entre o e4leno e o etano.


HH

H HH

! bond ! bond

" bond

" bond

a. b. c.

H

H

H

HC C CC



C C

H H

H




C¬Hsp2

(s)sp2sp2

sp2

side view top view

p

120°sp2

sp2sp2


sp2

sp2sp2



3-D Molecule:Ethene


p

1.33 A°

1.08 A°

121.7°

116.6°C C

H

H

H





for ethene

BRUI01-001_059r4 20-03-2003 2:58 PM Page 30

Densidade nega4va na ligação C=C denota maior

concentração eletrônica em decorrência da ligação π.

Força de ligação: C=C: 152 kcal/mol C-‐C: 88 kcal/mol

27

Algumas estruturas interessantes envolvendo apenas os átomos de

carbono.

Dê o nome de cada um dos compostos. Dê também a hibridização dos átomos de carbono em cada caso.

Química!Orgânica!I! Aula!2! - lqbo.ufscar.br · tracted to the positively charged nucleus of the...

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