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Transcript
Chapter 1
Structure and Bonding
© 2006 Thomson Higher Education
Scientific Revolution
Scientific revolution leads to:
• Safer and more effective medicines
• Cures for genetic diseases
• Improved quality of life
有机化学(Organic Chemistry ) –研究碳化合物的化学
Carbon
Organic chemistry
• All organic compounds contain the element of
carbon
•
•
•
•
•
Shares four electrons
Forms four strong covalent bonds
Bonds to other carbons to create chains and
rings
Not all carbon compounds are derived from
living organisms
Over 99% of 26 million known compounds
contain carbon
Chapter 1 Structure and bonding
1.1 Atomic Structure
• Atom: (直径2x10-10 m)
• Nucleus
• Protons (+)
10-14 ~ 10-15 m
• neutrons
• Electrons (-) 质量忽略不计,在大约 10-10 m 距离的范围内绕
核运动
1. 2 Atomic Structure: Orbitals
Quantum Mechanical Model (量子理学模型)
Behavior of a specific electron in an atom
described by mathematical expression called a
wave equation (原子中一个电子的行为可以用波动
方程的数学表达式来描述)
Solution of wave equation is called a wave
function(波动方程的解称为波函数)
Wave function is an orbital(波函数是一个轨道)
Orbital denoted by Greek letter psi, 
A plot of  2 describes volume of space around
nucleus the electron is most likely to occupy
Electron cloud has no specific boundary.
 2 表示原子核周围电子最容易出现的空间区域。
Wave Properties
of Electrons
• Standing wave vibrates in fixed location.
• Wave function, , mathematical description of
size, shape, orientation
• Amplitude(振幅) may be positive or negative
• Node: amplitude is zero
_
+
-
+
=>
Atomic Structure: Orbitals
Four different kinds of orbitals for electrons
•
Denoted s, p, d, and f
•
s and p orbitals most important in organic chemistry
• s orbitals
• Spherical, nucleus at center
• p orbitals
• Dumbbell-shaped, nucleus at middle
• d orbitals
• Four cloverleaf-shaped and one dumbbell-doughnut
Atomic Structure
• Atomic number (Z)
• Mass number (A)
• Atomic weight
• Orbital (): s, p, d, & f.
px, py, and pz
Atomic Structure: Orbitals
p Orbital
• In each shell, beginning with the second, there are three
perpendicular p orbitals, px, py, and pz, of equal energy
• Lobes of a p orbital are separated by a region of zero electron
density, a node
• Each lobe has a different algebraic(代数) sign, + and -,
represented by different colors
• Algebraic signs are not charges
1.3 Atomic Structure:
Electron Configuration(电子构型)
Orbitals are grouped in electron shells of increasing size and
energy
Electron Shell
•
•
•
•
Each orbital can be occupied by two electrons
First shell contains one s orbital, denoted 1s, which holds only two
electrons
Second shell contains four orbitals, one s orbital (2s) and three p
orbitals (2p), which hold a total of eight electrons
Third shell contains nine orbitals, one s orbital (3s), three p orbitals
(3p), and five d orbitals (3d), which hold a total of 18 electrons
Atomic Structure: Electron Configuration
P 15
Ground-state electron configuration
•
The most stable, lowest-energy electron
configuration of a molecule or atom
Three rules:German
1. Aufbau principle(构造原理) Aufbau meaning "building up
•
Lowest-energy orbitals fill first: 1s  2s  2p  3s 
3p  4s  3d
2. Pauli Exclusion Principle ( 泡利不相容原理)
•
Electron spin can have only two orientations, up 
and down 
•
Only two electrons can occupy an orbital, and they
must be of opposite spin to have unique wave
equations
Atomic Structure: Electron
Configuration
Hund‘s rule (洪特规则)
3.
•
If two or more empty orbitals of equal energy are
available, electrons occupy each orbital with parallel
spins until all orbitals have one electron
复习
Three rules:German
1. Aufbau principle Lowest-energy orbitals
fill first: 1s  2s  2p  3s  3p 
4s  3d
2. Pauli Exclusion Principle
Electron spin can have only two orientations, up  and
down  Only two electrons can occupy an orbital, and
they must be of opposite spin to have unique wave
equations
3. Hund's rule
•
If two or more empty orbitals of equal energy are
available, electrons occupy each orbital with
parallel spins until all orbitals have one electron
Atomic orbitals
Carbon has six electrons and is in row 2 of the periodic
table. This means that there are two shells of atomic
orbitals available for these electrons. The first shell closest
to the nucleus has a single s orbital – the 1s orbital. The
second shell has a single s orbital (the 2s orbital) and three
p orbitals (3 × 2p).
C 1s2 2s2 2px1 2py1
Carbon is in the second row of the periodic table and has six
electrons
● aufbau principle
● Pauli exclusion principle
● Hund’s rule.
1s2 2s2 2px1 2py1
复习
Atomic Structure
S
S
s,p
S
P
1.4 The Nature of Chemical Bonds:
Molecular Orbital Theory
(化学键的本质:分子轨道理论)
σ bonds
π bonds
Cl2: p-p overlap
Constructive overlap along the same
axis forms a sigma bond.
Cl–Cl
=>
σ bonds
σ
H–H
H2: s-s overlap
antibonding MO.
bonding MO.
H–Cl
CH4
Sigma bonds:σ bonds have a circular cross-section and
are formed by the head-on overlap of two atomic orbitals.
This is a strong interaction and so sigma bonds are strong
bonds.
A covalent bond binds two atoms together in a molecular
structure and is formed when atomic orbitals overlap to
produce a molecular orbital – so called because the
orbital belongs to the molecule as a whole rather than to
one specific atom.
antibonding MO.
bonding MO.
p bonds
•
•
•
Additive combination of two 2p orbitals
• Leads to formation of a low energy p bonding MO
• The p bonding MO is formed by combining p orbital lobes with the
same algebraic sign
• No node between nuclei
Subtractive combination of two 2p orbitals
• Leads to formation of a high energy p* antibonding MO
• The p* antibonding MO is formed by combining p orbital lobes
with different algebraic signs
• Node between nuclei
Only the bonding p MO is occupied
1.5 The Nature of Chemical Bonds:
Valence Bond Theory
Valence bond theory
• Bonding theory that describes a
covalent bond as resulting from
the overlap of two atomic
orbitals
• Electrons are paired in the
overlapping orbitals and are
attracted to nuclei of both atoms,
thus bonding the two atoms
together
•
•
•
H–H bond results from the
overlap of two singly occupied
hydrogen 1s orbitals
H-H bond is cylindrically
symmetrical
Bonds formed by head-on
overlap of two atomic orbitals
along a line drawn between the
nuclei are sigma (s) bonds
1.5 Valence Bond Theory
Bond strength
• H2 molecule has 436 kJ/mol less energy than the
starting 2 H atoms, the product is more stable than
the reactant and the H-H bond has a strength of 436
kJ/mol
• Conversely, the bond dissociation energy of H2 is 436
kJ/mol because it requires 436 kJ/mol of energy to
break the H2 bond
1.5 The Nature of Chemical Bonds:
Valence Bond Theory
There is an optimum distance between nuclei that leads to
maximum stability called the bond length
Bond length
• The distance between
nuclei at the minimum
energy point
• Because a covalent bond
is dynamic, like a spring,
the characteristic bond
length is the equilibrium
distance between the
nuclei of two atoms that
are bonded to each other
两个原子核之间有一个使键稳
定性最大的距离(74ppm)。
复
习
σ bonds
π bonds
1.6 Development of Chemical Theory
•
In 1858 August Kekulé (凯库勒,德国化学家) and Archibald Couper(库帕,
苏格兰) independently proposed that carbon is tetravalent (always forms four
bonds)Emil Erlenmeyer (德国化学家,提出用直短线代表化学键,两划即双键,
三划即三键的表示方法) proposed a carbon-carbon triple bond for acetylene,
Alexander Crum Brown (布朗) proposed a carbon-carbon double bond In
1865, Kekulé suggested that carbon chains can double back to form rings of
atoms In 1874, Jacobus van’t Hoff荷兰化学家范霍夫 and Joseph Le Bel (勒
贝尔) proposed four atoms to which carbon is bonded sit at the corner of a
regular tetrahedron
盲人摸象
1.6 Development of Chemical Theory
Atoms bond because the compound that results is more
stable and lower in energy than the separate atoms
•
•
Energy is released from the chemical system when a
bond forms
Energy is consumed by the system when a bond
breaks
Valence shell
• Outer most electron shell of an atom
• Eight electrons in valence shell (an electron octet)
impart special stability to noble-gas elements in 8A
• Main group elements are governed by their tendency
to take on electron configuration of the nearest noble
gas
1.6 Development of Chemical Theory
Ionic compounds
• Some elements achieve an octet configuration by gaining or
losing electrons
• When an electron is gained or lost from a neutral atom an ion
is formed
• Ions are charged because they have different numbers of
protons and electrons
• Ions are held together by an electrostatic attraction, like in
Na+ Cl-, forming an ionic bond
Carbon achieves an octet configuration by sharing
electrons
Covalent Bond
• A bond formed by sharing electrons between atoms
Molecule
• A neutral collection of atoms held together by covalent bonds
1.6 Development of Chemical Theory
Lewis structures (electron-dot structures)
• Representations of covalent bonds in molecules
Kekulé structures (line-bond structures)
1.6 Development of Chemical Theory
Number of covalent bonds depends on how many additional valence
electrons needed to reach noble-gas configuration
• H (1s) needs one more electron to attain (1s2)
• N (2s22p3) needs three more electrons to attain (2s22p6)
Lone-pair electrons
• Valence-shell electron pairs not used for bonding
• Lone-pair electrons can act as nucleophiles and react with
electrophiles
1.7 Hybridization
a carbon atom has two unpaired electrons and so we would
expect carbon to form two bonds.However, carbon forms four
bonds!
when a carbon atom forms bonds,it can ‘mix’ the s and p
orbitals of its second shell (the valence shell). This is known
as hybridization and it allows carbon to form the four bonds
which we observe in reality.There are three ways in which
this mixing process can take place.
● the 2s orbital is mixed with all three 2p orbitals. This is known
as sp3 hybridization;
● the 2s orbital is mixed with two of the 2p orbitals. This is known
as sp2 hybridization;
● the 2s orbital is mixed with one of the 2p orbitals. This is known
as sp hybridization.
sp3 hybridization
Energy levels
M
Methane
molecule forms 4 C−H sigma bonds and
e
t
is three
dimensional tetrahedral shape,bond
h
angleais 109.5°,cabon atom is sp3 hybridized in
n
methane
molecule
.
e
sp3 Hybrid Orbitals and the
Structure of Ethane
Orbital hybridization accounts for the bonding together of carbon
atoms into chains and rings
Ethane C2H6
• Tetrahedral
• Bond angles are near 109.5º
• Carbon-carbon single bond
• Formed by s overlap of
sp3 hybrids from each
carbon
• The remaining sp3
hybrids of each carbon
overlap with 1s orbitals of
three hydrogen atoms to
form six carbon-hydrogen
bonds
Ethane:
CH3-CH3
Question:
σ bonds C-H C-C?
3 Sigma bonds and one lone pair
:NH3
Geometry? Pyramidal,why?
2 sigma bonds and 2 lone pair
O
H
H
BENT
sp2 Hybrid Orbitals
有那些键?
sp2 Hybrid Orbitals and the
Structure of Ethylene
Ethylene C2H4
•
•
•
Carbon-carbon double bond
• Four shared electrons
Planar (flat)
Bond angles 120º
sp2 hybrid orbitals
• A hybrid orbital derived by
combination of an s atomic
orbital with 2p atomic orbitals
• One p orbital remains nonhybridized
πBonding
• Pi bonds form after sigma bonds.
• Sideways overlap of parallel p orbitals.
=>
Each ethene molecule forms 1 C–C σ bond ,
4 C–H σ bond and 1 π bond. Ethene is a flat,
rigid molecule where each carbon is trigonal
planar. Carbon atom is sp2 hybridized
and bond angle is 120°in ethene molecule.
• Ethylene
•
•
•
C2H4
s bond
π bond
Multiple Bonds
• A double bond (2 pairs of shared electrons)
consists of a sigma bond and a pi bond.
• A triple bond (3 pairs of shared electrons)
consists of a sigma bond and two pi bonds.
sp2 Hybrid Orbitals
sp2 Hybrid Orbitals
•
combination of an s orbital
with 2 pi orbitals - three sp2
hybrid orbitals
trigonal planar: 120o bond angles.
O
H 3C
H
Carbon has 3 sigma and one pi bond
O
Delocalization occurs in 1,3-butadiene where there are
alternating single and double bonds. All four carbons
in 1,3-butadiene are sp2 hybridized and so each of these
carbons has a half-filled p orbital which can interact to give
two π bonds. However, a certain amount of overlap is also
possible between the p orbitals of the middle two carbon
atoms and so the bond connecting the two alkenes has
some double bond character borne out by the observation
that this bond is shorter in length than a typical single
bond. This delocalization also results in increased
stability.
SP HYBRIDIZATION
• Acetylene
•
•
•
C2H2
One s-bond
& two p-bonds
=>
Each ethyne molecule forms 1 C–C σ bond ,
2 C–H σ bond. The 2py and 2pz orbitals of
each carbon atom can overlap side-on to
form two π bonds. Each carbon is sp
hybridized and bond angle is 180°in ethyne.
Ethyne is linear.
sp Hybrid Orbitals and the
Structure of Acetylene
Acetylene
•
•
Linear
Carbon-carbon triple bond
• Six shared electrons
• Bond angles are 180º
sp hybridized orbital
•
A hybrid orbital derived
from the combination of
one s and one p atomic
orbital
• The two sp hybrids are
oriented at an angle of
180º to each other
• Two 2p orbitals remain
non-hybridized
HCN
sp Hybrid Orbitals
Combine one 2 s with one 2 p sigma orbital
Linear 180 o
H-C
N
Note: both C and N are sp hybridized
H-C
N
sp Hybrid Orbitals and the Structure
of Acetylene
复习 :说出下列分子以及杂化类型
22_502
H
H
H
H
C
H
C
H
22_504
2p
2p
2p
H1s
H1s
2p
H
C
H
C
sp
2p
2p
2p
sp
2p
sp
22_501
sp2
sp2
sp2
H1s
sp
2
sp2
C
C
sp2
H1s
2p
A double bond is a s bond and a p bond.
Double bond B.D.E.
= 146 kcal/mol
s bond B.D.E.
= 83 kcal/mol
Therefore p B.D.E. must
= 63 kcal/mol.
A p bond is weaker than a s bond.
Bonds are more reactive than s bonds,
p bonds are considered
to be a functional group.
Double and Triple Bonds
H
1.33A
H
H
H
H
HH
1.08A
1.09A
HH
H
1.54A
sp2 are 1/3 s, whereas sp3 are 1/4 s in character.
(s orbitals are closer to the nucleus and lower in
energy).The carbon-carbon bond in ethene is
shorter and stronger than in ethane partly because
of the sp2-sp2 overlap being stronger than sp3-sp3,
but especially because of the extra p bond in ethene.
σ and π bonds Identifying σ and π bonds in a molecule
is quite easy as long as you remember
the following rules:
● all bonds in organic structures are either sigma (σ) or
pi (π) bonds;
● all single bonds are σ bonds;
● all double bonds are made up of one σ bond and one
π bond;
● all triple bonds are made up of one σ bond and two π
bonds
Hybridized centers: All the atoms in an organic
structure (except hydrogen) are either sp, sp2 or
sp3 hybridized
含双键和三键的分子比含有单键的分子活泼得多。
1.8 Polar Covalent Bonds:
Electronegativity
• Covalent bonds can have ionic character
• These are polar covalent bonds
• Bonding electrons attracted more strongly by one
atom than by the other
• Electron distribution between atoms is not
symmetrical
72
Bond Polarity and Electronegativity
• Electronegativity (EN): intrinsic ability of an atom to
•
•
•
•
•
•
attract the shared electrons in a covalent bond
Differences in EN produce bond polarity
Arbitrary scale. As shown in Figure 2.2,
electronegativities are based on an arbitrary scale
F is most electronegative (EN = 4.0), Cs is least (EN
= 0.7)
Metals on left side of periodic table attract electrons
weakly, lower EN
Halogens and other reactive nonmetals on right side
of periodic table attract electrons strongly, higher
electronegativities
EN of C = 2.5
73
The Periodic Table and Electronegativity
74
Bond Polarity and Inductive Effect
(键的极化和诱导效应)
•
•
•
•
•
诱导效应(Inductive Effect) :一个原子使键极化的能力。
Nonpolar Covalent Bonds: atoms with similar EN
Polar Covalent Bonds: Difference in EN of atoms <
2
Ionic Bonds: Difference in EN > 2
• C–H bonds, relatively nonpolar C-O, C-X bonds
(more electronegative elements) are polar
Bonding electrons toward electronegative atom
• C acquires partial positive charge, +
• Electronegative atom acquires partial negative
charge, Inductive effect: shifting of electrons in a bond in
response to EN of nearby atoms
75
Electrostatic Potential Maps
分子的静电势图
• Electrostatic potential
maps show calculated
charge distributions
• Colors indicate electronrich (red) and electronpoor (blue) regions
• Arrows indicate direction
of bond polarity
体系中静电势大的部位容易受亲电试剂
的攻击发生反应。
76
Bond Polarity & Electronegativity
• Covalent bond to ionic bond
EN = 0
0< EN <2.0
EN>2.0
1.8 Bond Polarity & Electronegativity
• Electrostatic potential maps
• Electron rich (red)
• Electron poor (blue)
1.9 Resonance
•
•
•
•
Some molecules are have structures that cannot be shown with a single
representation
In these cases we draw structures that contribute to the final structure but which
differ in the position of the p bond(s) or lone pair(s)
Such a structure is delocalized and is represented by resonance forms
The resonance forms are connected by a double-headed arrow
1.9 Resonance Hybrids
•
•
•
A structure with resonance forms does not alternate between
the forms
Instead, it is a hybrid of the two resonance forms, so the
structure is called a resonance hybrid
For example, benzene (C6H6) has two resonance forms with
alternating double and single bonds
•
In the resonance hybrid, the actual structure, all its C-C
bonds are equivalent, midway between double and single
80
1.9 Rules for Resonance Forms
• Individual resonance forms are imaginary - the real
•
•
•
•
structure is a hybrid (only by knowing the contributors
can you visualize the actual structure)
Resonance forms differ only in the placement of their
p or nonbonding electrons
Different resonance forms of a substance don’t have
to be equivalent
Resonance forms must be valid Lewis structures: the
octet rule applies
The resonance hybrid is more stable than any
individual resonance form would be
81
1.9 Curved Arrows and Resonance
Forms
• We can imagine that electrons move in pairs to
convert from one resonance form to another
• A curved arrow shows that a pair of electrons moves
from the atom or bond at the tail of the arrow to the
atom or bond at the head of the arrow
82
1.9 Drawing Resonance Forms
• Any three-atom grouping with a multiple bond
has two resonance forms
83
1.9 Different Atoms in Resonance
Forms
•
•
•
•
Sometimes resonance forms involve different atom types as well
as locations
The resulting resonance hybrid has properties associated with
both types of contributors
The types may contribute unequally
The “enolate” derived from acetone is a good illustration, with
delocalization between carbon and oxygen
84
2,4-Pentanedione
• The anion derived from 2,4-pentanedione
• Lone pair of electrons and a formal negative
charge on the central carbon atom, next to a
C=O bond on the left and on the right
• Three resonance structures result
85
86
87
Summary
•
•
•
•
Organic molecules often have polar covalent bonds as a result
of unsymmetrical electron sharing caused by differences in the
electronegativity of atoms
The polarity of a molecule is measured by its dipole moment, .
(+) and () indicate formal charges on atoms in molecules to
keep track of valence electrons around an atom
Some substances must be shown as a resonance hybrid of
two or more resonance forms that differ by the location of
electrons.