UNIVERSIDADE+FEDERAL+DE+SÃO+CARLOS+ Fundamentos+de… · A drawback of the model is that it treats...

Fundamentos de QO Aula 1

Prof. Marco Antonio B Ferreira [email protected]

3351-‐8075

1

UNIVERSIDADE FEDERAL DE SÃO CARLOS

www.lqbo.ufscar.br

Avaliação

17/09 – 1ª Avaliação 29/10 – 2ª Avaliação 03/12 – 3ª Avaliação

Aprovado: (P1 + P2 + P3)/3 ≥ 6

Exercícios: ao fim de cada aula será aplicado exercícios.

2

Biografia •  BRUICE, P.Y., Química Orgânica,4ª ed., vol. 1, São

Paulo, Pearson & Prin\ce Hall, 2006.

•  SOLOMONS, T.W.G. & FRYHLE, G., Química Orgânica, LTC, 8ª ed., Rio de Janeiro, vol 1 e 2, 2004.

•  CLAYDEN, J., GREEVES, N., WARREN, S. & WOTHERS, P., Organic Chemistry, Oxford, Oxford University, 2001.

3

O que é a Química Orgânica?

Jöns Jakob Berzelius (1779–1848)

1807: Cunhou o termo Química Orgânica

“Estudo de compostos derivados de fontes naturais (COMPOSTOS ORGÂNICOS seriam

derivados destas fontes).”

VITALISMO: “Força Vital” presente em organismos vivos era necessário para a síntese de compostos orgânicos. Compostos inorgânicos poderiam ser sinte\zados em laboratório e a par\r de compostos inorgânicos.

4

O que é a Química Orgânica? Friedrich Wöhler (1800–1882)

Aluno de Berzelius

1828: Síntese da Ureia.

INORGÂNICO ORGÂNICO

Wöhler e Berzelius estavam mais interessados no conceito de isomerismo desta reação. ORIGEM DA TEORIA ESTRUTURAL (Arranjo dos átomos define uma substância)

5

A Química Orgânica Moderna Compostos Orgânicos: Compostos que contém átomos de carbono.

VIVEMOS NA ERA DA QUÍMICA ORGÂNICA

•  >90 % de todos os compostos conhecidos são compostos por carbono.

•  Química Orgânica é crucial para nosso modo de vida: Agricultura, Roupas, Polímeros/Materiais, Medicina, CombusFvel.

6

A Química Orgânica Moderna ÁREAS DA QUÍMICA ORGÂNICA

Mecanismo

Síntese

Estrutura

Como e porque os átomos são

dispostos para formar moléculas.

Como e por que as reações químicas

acontecem.

Preparação de moléculas complexas a parNr de moléculas mais simples usando reações químicas. 7

Por que o carbono é tão especial? Eletroposijvos

Eletronegajvos

•  Carbono pode fazer até 4 ligações covalentes (estável), com outros átomos de carbono ou com outros elementos da tabela (GRANDE DIVERSIDADE ESTURTURAL).

> 16 milhões de compostos conhecidos!!!! 8

Estrutura Atômica

Quem não viu/lembra, fazer uma revisão usando o livro do Atkins.

9

Estrutura Eletrônica dos átomos O que aprendemos da Mecânica Quânjca?

•  Dualidade onda-‐parqcula da matéria:

Prof. Marco: 115 kg andando a 5 m/s. λ = 1,15 x 10-‐36 m

ℎ = 6,62606957 × 10-‐34 m2 kg / s

Elétron: 9,10938291(40)×10−31 kg andando a 1/1000 da velocidade da luz.

λ = 2,43 x 10-‐9 m

10

Estrutura Eletrônica dos átomos

Difração de elétrons •  Orbitais Atômicos

11


Ĥ Ψ = E Ψ Equação de Schrödinger

Operador Hamiltoniano

12

Ψ (função de onda) = função matemá\ca que descreve a região de maior probabilidade em encontrar os elétrons de um átomo.


energia QUANTIZADA

13

Cada elétron em um átomo é descrito por 4 números quânjcos....


Camada Sub-‐camada orbitais

14


15


16


17


Energia dos orbitais

1s

2s

3s

2p

3p 3d

Penetração radial

quanto mais próximo do núcleo, menor a energia

do orbital.

Orbitais degenerados: (2px, 2py, 2pz), (3px, 3py, 3pz), (3dxy, 3dxz, 3dyz, 3dx2-‐y2, 3dz2) 18

Princípio de auxau: Elétron sempre ocupa o orbital de menor energia em um átomo no seu estado fundamental.

Princípio de exclusão de Pauling: Dois elétrons não podem ocupar o mesmo estado quân\co simultaneamente.

C Z = 6

1s2 2s2 2p2

1s

2s

2px 2py 2pz

spins anA-‐paralelos

spins paralelos


Regra de Hund: Elétrons numa mesma subcamada tendem a permanecer desemparelhados (em orbitais separados), com spins paralelos.

Quantos elétrons na camada de valência?

19


20

Ligações Iônicas e Covalentes Átomos perdem, recebem ou comparjlham elétrons a fim de terem sua camada de valência completa. Isso confere estabilidade ao sistema. (G. N. Lewis)

A. Ligação iônica: Ocorre entre átomos com grande diferença de eletronega\vidade (EN) ΔEN>2.

Ex.: NaCl (EN: Na = 0,93 ; Cl = 3,16)

Na Na+ + e-‐

1s2 2s2 2p6 3s1 1s2 2s2 2p6 = [Ne]

Cl + e-‐ Cl-‐

1s2 2s2 2p6 3s2 3p5 1s2 2s2 2p6 3s2 3p6 = [Ar]

Usando estruturas de Lewis:

ΔE = -‐153 kcal/mol

21

Ligações Iônicas e Covalentes

22

Ligações Iônicas e Covalentes B. Ligação covalente: Ocorre entre átomos com eletronega\vidade similar.

23

Ligações Iônicas e Covalentes Dois átomos podem comparjlhar mais do que um par de elétrons.

24

Ligações Covalentes Polares

25


Ligação covalente apolar (átomos de mesma eletronega\vidade)

Ligação covalente polar (átomos com eletronega\vidade diferente)

Dipolo na molécula

26


Mapa de Potencial eletrostájco (mostra como a densidade eletrônica da molécula é distribuída entre os átomos)

EN: H = 2,1; Li = 1,0; F = 4,0 27


Polarizabilidade (habilidade de distorção da densidade eletrônica)

Família: Aumenta descendo a tabela. Período: Aumenta da esquerda pra direita.

28

Carga Formal (CF) CF = EV – (NL+EL/2) onde:

EV = número de elétrons de valência, NLP = número de elétrons não ligantes, NL = número de elétron de ligações

Carga formal não denota que a carga

esteja toda concentrada

naquele átomo.

SEMPRE INDIQUE A CARGA FORMAL NOS EXERCÍCIOS E PROVAS 29

Estrutura Condensada Podemos omi\r uma ou mais ligações covalentes:

30

Estrutura Condensada

31

1) Dê as estruturas de Lewis de cada um dos seguintes compostos:

4) Examinando o mapa de potencial eletrostá\co do LiH, HF e H2, responda as seguintes questões: a) Qual destes compostos são polares?

b) Por que o LiH tem o hidrogênio com maior densidade eletrônica?

c) Quais destes compostos apresenta o hidrogênio mais carregado posi\vamente?

Questões:

32

3) Dê configuração eletrônica dos seguintes elementos: K (Z = 19), Cl (Z = 17) e Br (Z = 35).

2) Dê as estruturas de Lewis de cada um dos seguintes compostos:

Section 1.4 Representation of Structure 17

Kekulé structure Condensed structures

Two or more identical groups considered bonded to the “first” atom on the left can be shown (in parentheses) to the left of thatatom, or hanging from the atom.

An oxygen doubly bonded to a carbon can be shown hanging off the carbon or to the right of the carbon.

O

or O)CH3CH3CH2CCH3 CH3CH2COCH3 or CH3CH2C(

O

or OCH3CH2CH2CH CH3CH2CH2CHO or CH3CH2CH2CH

O

orCH3CH2COH CH3CH2CO2H or CH3CH2COOH

O

orCH3CH2COCH3 CH3CH2CO2CH3 or CH3CH2COOCH3

H HC

H

(CH3)2NCH2CH2CH3 CH3NCH2CH2CH3or

CH3

C

HH

H

C

H

HH H H

H

C

C

H

N

H HC (CH3)2CHCH2CH2CH3 CH3CHCH2CH2CH3or

CH3

H

HH H H

H

C

C

HH

H

C

H

H

C

H

C

PROBLEM 10 SOLVED

Draw the Lewis structure for each of the following:

a. d. g. j. NaOHb. e. h. k.c. f. i. l.

SOLUTION TO 10a The only way we can arrange one N and three O’s and avoidsingle bonds is to place the three O’s around the N. The total number of valence

electrons is 23 (5 for N, and 6 for each of the three O’s). Because the species has one neg-ative charge, we must add 1 to the number of valence electrons, for a total of 24. We thenuse the 24 electrons to form bonds and fill octets with lone-pair electrons.

When all 24 electrons have been assigned, we see that N does not have a complete octet. Wecomplete N’s octet by using one of oxygen’s lone pairs to from a double bond. (It doesn’tmake any difference which oxygen atom we choose.) When we check each atom to seewhether it has a formal charge, we find that two of the O’s are negatively charged and the Nis positively charged, for an overall charge of

O

N+O! O !

-1.

O

NO O

O¬O

Na2CO3-CH3N2NO2

-NH4Cl+C2H5HCO3

-NO2

+CH3NH3

+CO2NO3

-

Table 1.5 (continued)

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33

Introdução a Teoria do Orbital Molecular •  Como átomos formam ligações?

Segundo Lewis: Busca por ter sua camada de valência completa a par\r do compar\lhamento de elétrons.

TOM: Combinação de orbitais atômicos levam a formação de orbitais moleculares.

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 Page 20

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

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

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34

Introdução a Teoria do Orbital Molecular

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 Page 20

35

Introdução a Teoria do Orbital Molecular


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

+

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 23Section 1.6 An Introduction to Molecular Orbital Theory 23

nodes

nodenode



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


S

P





psp

s

psp*

pp*

p1P2s

s*




2p atomicorbital

2p atomicorbital

nodal plane

nodal plane

nodal plane

Ener

gy


p*p

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


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


p*p

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

+


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


p*p

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


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


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

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UNIVERSIDADE+FEDERAL+DE+SÃO+CARLOS+ Fundamentos+de… · A drawback of the model is that it treats...

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