7
Chemical Bonding
化學鍵結
Chapter Goals
1. Lewis Dot Formulas of Atoms 路易士電子點結構式
Ionic Bonding 離子鍵結
2. Formation of Ionic Compounds 形成離子化合物
Covalent Bonding 共價鍵結
3. Formation of Covalent Bonds 形成共價鍵
4. Bond Lengths and Bond Energies 鍵長及鍵能
5. Lewis Formulas for Molecules and Polyatomic Ions 分子極多元
子離子之路易士結構式
6. Writing Lewis Formulas: The Octet Rule 八隅體法則
7. Formal Charges 形式電荷
8. Writing Lewis Formulas: Limitations of the Octet Rule 八隅體法
則之限制
9. Resonance 共振
10. Polar and Nonpolar Covalent Bonds 極性及非極性共價鍵結
11. Dipole Moments 偶極距
12. The Continuous Range of Bonding Types 鍵結形式之連續範圍
2
Introduction
• Attractive forces that hold atoms together in
compounds are called chemical bonds (化合物中將原
子拉進的吸引力稱之為化學鍵結)
• The electrons involved in bonding are usually those
in the outermost (valence) shell (與鍵結有關的電子通
常是指最外層的電子)
3
Introduction
• Chemical bonds are classified into two types:
離子鍵結 results from electrostatic
attractions among ions, which are formed by the transfer
of one or more electrons from one atom to another.
 Ionic bonding

離子鍵為陽離子與陰離子間之庫倫靜電力，陽離子由金屬失去電子形成，陰離
子由非金屬得到電子形成。兩正負電荷之離子吸引力，稱為離子鍵。
 Ionic compound 離子化合物 NaCl
 Covalent bonding 共價鍵結 results from sharing one or
more electron pairs between two atoms.

價鍵為兩原子共用電子對之化學鍵，其結合原子皆為非金屬原子
 Covalent compound H2, Cl2
4
Ionic Compound 離子化合物
Covalent Compound 共價化合物
Solids
Gases, liquid, or solids
High melting points (>400oC)
Low melting points (<300oC)
Soluble in polar solvents, such as
water
insoluble in polar solvents
insoluble in nonpolar solvents, such
as C6H14 and CCl4
Most are soluble in nonpolar
solvents, such as C6H14 ,CCl4
Molten compounds conduct
electricity well because they contain
ions
Liquid and molten compounds do not
conduct electricity
Aqueous solutions conduct electricity
well because they contain ions
Aqueous solutions are usually poor
conductors
Two elements with different
electronegativities, usually a metal
and a nonmetal
Two elements with similar
electronegativities, usually nonmetals
5
Lewis Dot Formulas of Atoms 原子的路易
士電子點結構式
• Lewis dot formulas or Lewis dot representations are a
convenient bookkeeping method for tracking valence
electrons 價電子.
– The electrons in the outermost occupied shells (最外層的
電子數)
• s and p orbitals
– Valence electrons are those electrons that are
transferred or involved in chemical bonding (價電子是指
與化學鍵結有關或轉移的電子)
• They are chemically important
6
Lewis Dot Formulas 路易士結構
of Atoms
Elements that are in the same periodic group have the same
Lewis dot structures (同族元素具相同路易士結構)
1 electron in valence shell
5 electrons in valence shell
Not as useful for the transition and inner transition elements
7
(不適用於過渡元素)
Ionic Bonding 離子鍵結
Formation of Ionic Compounds 形成離子化
合物
•
An ion is an atom or a group of atoms possessing a
net electrical charge.
• Ions come in two basic types:
– positive (+) ions or cations
•These atoms have lost 1 or more electrons.
– negative (-) ions or anions
•These atoms have gained 1 or more electrons.
8
Formation of Ionic Compounds
• Monatomic ions consist of one atom.
– Examples:
• Na+ sodium ion, Ca2+, Al3+ -- cations
• Cl- chloride ion, O2-, N3- -- anions
• Polyatomic ions contain more than one atom.
– Examples:
• NH4+ ammonium ion -- cation
• NO2-, CO32-, SO42- sulfate ion – anions
– The atoms of a polyatomic ion are held together by
covalent bonds (多原子離子之原子間以共價鍵結)
9
Ionic Bonding is the attraction of oppositely charged
ions (cations and anions) in large numbers to form a
solid. Such a solid compound is called an ionic solid.
• Elements that have low electronegativities and low
ionization enengy — metals
(oxidozed; lose electrons to form cations)
• Elements that have high electronegativities and very
negative electron affinities — nonmetals
(reduced; gain electrons to form anions)
React
Ionic compound
10
Formation of
Ionic Compounds
• Reaction of Group IA Metals with Group VIIA
Nonmetals
1A metal 7A nonmetal
2Li(s) +
F2(g)
 2LiF(s)
Silver
yellow
white solid
Solid
gas
with an 842oC
melting point
11
Formation of
Ionic Compounds
• The underlying reason for the formation of LiF lies in
the electron configurations of Li and F.
1s 2s
2p
Li  
F   
These atoms form ions with these configurations.
Li+ 
same configuration as [He]
Li+ isoelectronic with He
loss one electron
F-    same configuration as [Ne]
F+ isoelectronic with Ne
gained one electron
等電子離子
12
Formation of
Ionic Compounds
• We can also use Lewis dot formulas to represent the neutral
atoms and the ions they form.
13
Formation of
Ionic Compounds
•The Li+ ion contains two electrons, same as the helium
atom.
– Li+ ions are isoelectronic 等電子離子 with helium.
•The F- ion contains ten electrons, same as the neon
atom.
– F- ions are isoelectronic with neon.
•Isoelectronic species contain the same number of
electrons.
•Most ionic compounds formed by reactions between
representative metals 典型金屬and representative
nonmetals 典型非金屬
14
Formation of
Ionic Compounds
• The reaction of potassium with bromine is a second
example of a group IA metal with a Group VIIA non
metal.
– Write the reaction equation.
1A metal
2K(s) +
7A nonmetal
Br2(g)
 2KBr(s)
ionic solid
15
Formation of
Ionic Compounds
• We look at the electronic structures of K and Br.
4s
4p
K [Ar] 
Br [Ar]     and the d electrons
The atoms form ions with these electronic structures.
4s
K+
Br-
4p
same configuration as [Ar]

  
same configuration as [Kr]
16
Formation of
Ionic Compounds
• Write the Lewis dot formula representation for the
reaction of K and Br.
17
Formation of
Ionic Compounds
• There is a general trend evident in the formation of
these ions.
• Cations become isoelectronic with the preceding
noble gas. 之前的鈍氣
• Anions become isoelectronic with the following
noble gas. 之後的鈍氣
18
Formation of
Ionic Compounds
• In general for the reaction of IA metals and VIIA
nonmetals, the reaction equation is:
2 M(s) + X2  2 M+ X-(s)
– where M is the metals Li to Cs
– and X is the nonmetals F to I.
Electronically this is occurring.
ns
M 
X 
np
  
ns
 M+
 X- 
np
  
19
Formation of
Ionic Compounds
• Next we examine the reaction of IIA metals with VIIA
nonmetals.
• This reaction forms mostly ionic compounds.
– Notable exceptions are BeCl2, BeBr2, and BeI2
which are covalent compounds.
• One example is the reaction of Be and F2.
Be(s) + F2(g) BeF2(g)
20
Formation of
Ionic Compounds
• The valence electrons in these two elements are
reacting in this fashion.
2s
2p
2s
2p
Be [He] 
 Be2+
F [He]      F-    
Next, draw the Lewis dot formula representation of this
reaction.
21
Formation of
Ionic Compounds
• The remainder of the IIA metals and VIIA
nonmetals react similarly.
• Symbolically this can be represented as:
M(s) + X2  M2+ X2M can be any of the metals Be to Ba.
X can be any of the nonmetals F to Cl.
22
Formation of
Ionic Compounds
• For the reaction of IA metals with VIA nonmetals, a
good example is the reaction of lithium with oxygen.
• The reaction equation is:
23
Formation of
Ionic Compounds
• Draw the electronic configurations for Li, O, and
their appropriate ions.
2s
2p
Li [He] 
O [He]    
2s
2p
 Li1+
 O2-    
Draw the Lewis dot formula representation of this
reaction.
24
Formation of
Ionic Compounds
• The remainder of the IA metals and VIA
nonmetals behave similarly.
• Symbolically this can be represented as:
2 M (s) + X  M21+ X-
M can be any of the metals Li to Cs.
X can be any of the nonmetals O to Te.
25
Formation of
Ionic Compounds
• The reaction of IIA metals and VA nonmetals also
follows the trends that we have established in this
chapter.
• The reaction of calcium with nitrogen is a good
example.
• The reaction equation is:
3Ca(s) + N2(g)  Ca3N2 (s)
26
Formation of
Ionic Compounds
• Draw the electronic representation of Ca, N, and their
ions.
4s 4p
4s
4p
Ca [Ar]

 Ca2+
2s 2p
2s
2p
N [He]
     N3-    
• Draw the Lewis dot representation of this reaction.
27
Formation of
Ionic Compounds
• Other IIA and VA elements behave similarly.
• Symbolically, this reaction can be represented as:
3 M(s) + 2 X(g)  M32+ X23M can be the IIA elements Be to Ba.
X can be the VA elements N to As.
28
Formation of
Ionic Compounds
• d-transition Metal Ions
鈧
– The outermost s electrons and energy level lower d
electrons are always the first ones lost when transition
metals form simple ions.
3d
4s
3d
4s
Sc[Ar] 

 Sc3+[Ar]
Zn[Ar]       Zn2+[Ar]     
3e- lost
2e- lost
– Most other 3d-transition metals can form at least two cations
in their compounds.
3d
4s
3d
4s
Co[Ar]      
 Co2+[Ar]     
Co[Ar]      
Co3+[Ar]     
2e- lost
3e- lost
29
Formation of Ionic Compounds
3A
H, a nonmetal, forms ionic compounds with IA and IIA metals for example, LiH,
KH, CaH2, and BaH2. Other hydrogen compounds are covalent.
Group IA and IIA can form peroxide (contain O22- ion ) or superoxide (contain
O2- ion). The peroxide and superoxide ions contain atoms that are covalently 30
bonded to one another
Group IA and IIA can form peroxide (contain O22- ion ) or superoxide (contain
O2- ion). The peroxide and superoxide ions contain atoms that are covalently
bonded to one another
31
Formation of
Ionic Compounds
• Ionic compounds form extended three dimensional
arrays of oppositely charged ions.
• Ionic compounds have high melting points because
the coulomb force庫侖力, which holds ionic
compounds together, is strong.
32
Formation of
Ionic Compounds
• Coulomb’s Law describes the attraction of positive ions for
negative ions due to the opposite charges.

q q 
F


2
d
where
F  force of attraction between ions
q  magnitude of charge on ions
d  distance between center of ions
33
Formation of
Ionic Compounds
• Small ions with high ionic charges have large Coulombic
forces of attraction.
• Large ions with small ionic charges have small Coulombic
forces of attraction.
• Use this information, plus the periodicity rules from Chapter
6, to arrange these compounds in order of increasing
attractions among ions
KCl, Al2O3, CaO
3
2
2
1
Al O  Ca O  K Cl
23
2-
-
34
Covalent Bonding 共價鍵結
• Covalent bonds are formed when atoms share electrons. It
Occurs when the electronegativity difference between
elements (atoms) is zero or relativity small (電負度幾乎沒差)
• The bonds between atoms within a molecule
(intramolecular bonds 分子內鍵結) are relatively strong, but
the force of attraction between molecules (intermolecular
forces 分子間鍵結) are relatively weak
lower melting and boiling points than ionic compound (較離
子化合物的熔點及沸點低)
– If the atoms share 2 electrons a single covalent bond is formed
(若原子共享2個電子則形成單一共價鍵)
– If the atoms share 4 electrons a double covalent bond is
formed (若原子共享4個電子則形成二個共價鍵)
– If the atoms share 6 electrons a triple covalent bond is formed
(若原子共享6個電子則形成三個共價鍵)
• The attraction between the electrons is electrostatic in
nature
– The atoms have a lower potential energy when bound.
35
Formation of
Covalent Bonds
•Representation of the
formation of an H2 molecule
from H atoms.
The electron of each atom is attracted by the
positively charged nucleus of the other atom, so the
electron density begins to shift (blue arrows) (電子受
到帶正電的原子核的吸引, 電子團開始變化 )
The electron clouds of the two atoms repel one
another, and so do the nuclei of the two atoms (Red
arrows) (兩個原子亦會有排斥力)
The two 1s orbitals overlap
36
Formation of Covalent Bonds
• This figure shows the potential energy of an H2
molecule as a function of the distance between the two
H atoms.
太靠近,產生斥力
太遠,引力太小,無鍵結
斥力與引力達成
平衡,穩定的排列
For any covalent bond there is an
internuclear distance where the attractive
and repulsive forces balance. This distance
is the bond length (彼此間的距離稱鍵長). The
energy difference is the bond energy (能量
的差異稱為鍵能)
37
Bond dissociation energy 鍵離解能
Bond energy 鍵能
38
39
Formation of
Covalent Bonds
We can use Lewis dot formulas to show covalent bond
formation.
1.H molecule formation representation.
H + H
 H:H or H-H
2. HF molecule formation

H + F :

 or HF
 H:F:


:F + F:
 
 or F
 :F:F:
2
 
3. F2 molecule formation
40
H O

H
Dot formula

or H O
H
Dash formula線結構式
• CO2
• NH4+
N: 5 electrons
H: 1 electrons
Total 9 electrons




• H2O
Writing Lewis Formulas:
The Octet Rule八隅體法則
•N - A = S rule
Simple mathematical relationship to help us write Lewis dot formulas.
–N = number of electrons needed to achieve a noble gas
configuration. (要達到鈍氣組態的電子數目,通常為8)
• N usually has a value of 8 for representative elements.
• N has a value of 2 for H atoms.
–A = number of electrons available in valence shells of the
atoms. (原子的價電子數)
• A is equal to the periodic group number for each element.
• A is equal to 8 for the noble gases.
–S = number of electrons shared in bonds. (形成鍵結可共享的電
子數)
–A-S = number of electrons in unshared, lone, pairs. (不共享的電
子數,又稱孤電子對)
42
Lewis Formulas for Molecules and
Polyatomic Ions
•First, we explore Lewis dot formulas of homonuclear
diatomic molecules.
–Two atoms of the same element.
1.Hydrogen molecule, H2.
N=2x2=4 e- needed
A=2x1=2 e- available
S=N-A=2 e- shared
H:H or H−H


or N  N

N





N
F − F
 




2. Fluorine, F2


F F or
 
3. Nitrogen, N2
N=2x8=16 e- needed
A=2x7=14 e- available
S=N-A=2 e- shared
N=2x8=16 e- needed
A=2x5=10 e- available
S=N-A=6 e- shared
2 molecule with 6 electrons
3 covalent bond
43
Lewis Formulas for Molecules and
Polyatomic Ions
•Next, look at heteronuclear diatomic molecules.
–Two atoms of different elements.
•Hydrogen halides 鹵化氫 are good examples.
1. hydrogen fluoride, HF
or


H− F




H F







2. hydrogen chloride, HCl


H Cl or H − Cl


3. hydrogen bromide, HBr


H Br or H − Br


N=1x2+1x8=10 e- needed
A=1x1+1x7=8 e- available
S=N-A=2 e- shared
44
Lewis Formulas for Molecules and
Polyatomic Ions
• Now we will look at a series of slightly more
complicated heteronuclear molecules.
• Water, H2O
or

H O
H
N=2x2+1x8=12 e- needed
A=2x1+1x6=8 e- available
S=N-A=4 e- shared



H 
O

H
• Ammonia molecule , NH3


H 
N H

H
or
 H
H N
H
N=3x2+1x8=14 e- needed
A=3x1+1x5=8 e- available
S=N-A=6 e- shared
45
Lewis Formulas for Molecules and
Polyatomic Ions
• Lewis formulas can also be drawn for
molecular ions.
• One example is the ammonium ion , NH4+.
N=1x8+4x2=16 e- need
1C atom 4H atom
A=1x5+4x1-1=8e- available
1C atom 4H atom (+1charge)
S=N-A=16-8=8e- share
•Notice that the atoms other than H in these
molecules have eight electrons around them.
46
A Guild to Writing Lewis Formulas
1. Select a reasonable (symmetrical) “skeleton” for the molecule
or polyatomic ion 選擇最合理分子當作骨架
–The least electronegative element is usually the central
element, except the H 電負度最小者通常為中心分子,氫除外
–Carbon bonds to two, three or four atoms, but never more
than four 碳可與2, 3及4個原子鍵結
–Oxygen atoms do not bond to each other except in (氧原子不
會鍵結在一起,除非)
•O2 and O3
•hydrogen peroxide, H2O2, and the peroxide contain the O22group
•Superoxide, which contain the O2- group
–In ternary oxoacids, hydrogen usually bond to an O atom,
not to the central atom, HNO2 三元含氧酸中,氫通常與氧鍵結
–For those have more than one central atom, the most
symmetrical skeletons possible are used, such as C2H4,
P2O74-
A Guild to Writing Lewis Formulas
2. Calculate N, the number of valence shell electrons needed by all
atoms in the molecule or iron to achieve noble gas configurations
PF3 N=1x8(P atom)+3x8(F atoms) = 32e- need
CH3OH N=1x8(C atom)+4x2(H atoms)+1x8 (O atom)
= 24e- need
NO3- N=1x8(N atom)+3x8(O atoms) = 32e- need
Calculate A, the number of electrons available in the outer shells of
all the atoms
PF3 A=1x5(P atom)+3x7(F atoms) = 26e- available
CH3OH A=1x4(C atom)+4x1(H atoms)+1x6 (O atom)
= 14e- available
NO3- A=1x5(N atom)+3x6(O atoms)+1 (for 1-charge)
= 24e- available
A Guild to Writing Lewis Formulas
Calculate S, total number of electrons shared in the molecule or
ion, using the relationship S=N-A
PF3
S=N-A = 32-26=6 e- shared (3 pairs of e-shared)
CH3OH S=24-14= 10 e- shared (5 pairs of e-shared)
NO3- S= 32-24=8 e- hared (4 pairs of e-shared)
3. Place the S electron in to the skeleton as shared pairs. Use
double and triple bonds only when necessary.
S=6
(3 pairs of e-shared)
S=10
(5 pairs of e-shared)
S=8
(4 pairs of e-shared)
4. Check the additional electrons into the skeleton as unshared
pairs to fill the octet of every A group element (except H,
shared only 2e-)
PF3
CH3OH
NO3-
Writing Lewis Formulas:
The Octet Rule
• The octet rule states that representative elements
usually attain stable noble gas electron
configurations in most of their compounds.
• Lewis dot formulas are based on the octet rule.
• We need to distinguish between bonding (or shared)
electrons and nonbonding (or unshared or lone
pairs) of electrons.
52
Writing Lewis Formulas:
The Octet Rule
• Example 7-2: Write Lewis dot and dash formulas for
hydrogen cyanide, HCN.
N = 2 (H) + 8 (C) + 8 (N) = 18
A = 1 (H) + 4 (C) + 5 (N) = 10
S= 8
A-S = 2
• This molecule has 8 electrons in shared pairs and 2
electrons in lone pairs.
or
C
C
N

N





H C
53
Writing Lewis Formulas:
The Octet Rule
• Example 7-3: Write Lewis dot and dash
formulas for the sulfite ion, SO32-.
N = 8 (S)+3 x 8 (O)
= 32
A = 6 (S)+3 x 6 (O)+ 2 (- charge)= 26
S = 6
A-S = 20
• Thus this polyatomic ion has 6 electrons in
shared pairs and 20 electrons in lone pairs.
• Which atom is the central atom in this ion?
54
Writing Lewis Formulas:
The Octet Rule
• What kind of covalent bonds, single, double, or triple,
must this ion have so that the six shared electrons are
used to attach the three O atoms to the S atom?

O


or
2


O
O
 S 


 
 

2


O S O
 

O


55
Example 7-1 Writing Lewis Formula
Write the Lewis formula for the nitrogen molecule, N2, carbon disulfide, CS2
and the carbonate ion, CO32-
N=1x8+3x8=32 e- needed
A=1x4+3x6+2=24 e- available
S=N-A=8 e- shared
Four pairs are shared

O

C
or
S =C= S

O
2-







C S


CO32-
S
N N
or
or
Exercise 29, 30, 38
–
(b)
N=1x8+2x8=24 e- needed
A=1x4+2x6=16 e- available
S=N-A=8 e- shared
3 molecule with 8 electrons
2x2 covalent bond
N




(b) CS2
N


(a) N2
N=2x8=16 e- needed
A=2x5=10 e- available
S=N-A=6 e- shared
2 molecule with 6electrons
3 covalent bond
C
2-
甲烷
乙烷
乙烯
hydrocarbon
乙炔
甲醛
氯仿,三氯甲烷
乙醇,酒精
Formal Charge 形式電荷
• Calculation of a formal charge on a molecule is a mechanism
for determining correct Lewis structures
• The formal charge is the hypothetical charge on an atom in
a molecule or polyatomic ion.
• The best Lewis structures will have formal charges on the
atoms that are zero or nearly zero.
形式電荷：基於共用電子對由兩鍵結原子平分共用的原則
下，此原子的價電子數與在自由原子狀態的價電子數兩者
的差值。
58
Formal Charge
Rules for Assigning Formal Charge
1. Formal Charge = group number – (number of
bonds + number of unshared e-)
2. An atom that has the same number of bonds as
its periodic group number has a formal charge of
0.
3. a. The formal charges of all atoms must sum to 0
in molecules.
b. The formal charges must sum to the ion’s
charge for a polyatomic ion.
♦ 形式電荷的定義:
(路易士結構中的) 價電子數 = 孤電子數 + ½鍵結電子數
形式電荷 =自由原子的價電子數  價電子數
=自由原子的價電子數  孤電子數  ½鍵結電子數59
In NH3
N atom: group 5A, has 3 bonds and 2 unshared eFC =(group number) –[(number of bonds)+(number of unshared e-)]
= 5-(3+2)=0
H: FC= 1-(1+0)=0
the formal charge of N and H are both zero in NH3, so the sum of
the formal charged is 0+3(0)=0
In NH4+
the atom has four bonds and no shared efor N FC=5-(4+0)=1
for H FC=1-(1+0)=0
NH4+=1+4(0)=1
60
Formal Charge
Consider nitrosyl chloride, NOCl
i
Cl
N
O
Cl N O
ii
7 – (2+4) = +1
5 – (3+2) = 0
6 – (1+6) = -1
Cl
N
O
Cl N O
7 – (1+6) = 0
5 – (3+2) = 0
6 – (2+4) = 0
(ii) is a preferable Lewis formula because
it has smaller charges than (i)
61
Writing Lewis Formulas:
Limitations of the Octet Rule
• There are some molecules that violate違背the octet rule.
–For these molecules the N - A = S rule does not apply:
1. The covalent compounds of Be.
– Use 4 electrons as N number
2. The covalent compounds of the IIIA Group.
– Use 6 electrons as N number
3. Species which contain an odd number of electrons, such as
NO2.
– NO with 11 valence shell electrons and NO2 with 17.
4. Species in which the central element must have a share of
more than 8 valence electrons to accommodate all of the
substituents.
5. Compounds of the d- and f-transition metals.
62
Writing Lewis Formulas:
Limitations of the Octet Rule
• In those cases where the octet rule does not apply, the
substituents attached to the central atom nearly always
attain noble gas configurations.
• The central atom does not have a noble gas configuration
but may have fewer than 8 (exceptions 1, 2, & 3) or more
than 8 (exceptions 4 & 5).
63
Example 7-1 Writing Lewis Formula
Write the Lewis formula for the beryllium chloride, BeCl2, boron
trichloride, BCl3 and the phosphorus pentafluoride, PF5
 
  



B

Cl

or
F








F
P
Cl
Cl



 
 


Cl
Cl B Cl
or






Cl Be Cl or Cl Be Cl


N=5x8+1x8=48 e- needed
A=5x7+1x5=40 e- available
S=N-A=8 e- shared
5 F bond to P. This requires the
sharing a minimum 10 e-.
Increase S from 8 e- to 10 e-.


(b) PF5
N=3x8+1x6=30 e- needed
A=3x7+1x3=24 e- available
S=N-A=6 e- shared



(b) BCl3
N=2x8+1x4=20 e- needed
A=2x7+1x2=16 e- available
S=N-A=4 e- shared


(a) BeCl2
P
64
Example 7-1 Writing Lewis Formula
Write the Lewis formula for the sulfir tetrafluoride, SF4, the triiodide
ion, I3-
(a) SF4
N=1x8+4x8=40 e- needed
A=1x6+4x7=34 e- available
S=N-A=6 e- shared
Requires a minimum 8e-
(b) I3N=3x8=24 e- needed
A=3x7+1x1=22 e- available
S=N-A=2 e- shared
Exercise 62
65
Resonance 共振
• Resonance is a flawed method of representing
molecules.
– There are no single or double bonds in SO3.
• In fact, all of the bonds in SO3 are equivalent.
• The best Lewis formula of SO3 that can be drawn is:
共價鍵結構的混成體,稱之為共振
66
The typical C-O single
bond length is 1.43Å
The typical C=O double
bond length is 1.22Å
The C-O bond in the CO32ion is at 1.29Å
Delocalization of bonding
electron 非定域電子鍵
67
Polar and Nonpolar Covalent Bonds 極性
及非極性共價鍵
• Covalent bonds in which the electrons are shared
equally are designated as nonpolar covalent bonds.
– Nonpolar covalent bonds have a symmetrical
charge distribution (對稱的電荷分佈)
• To be nonpolar the two atoms involved in the bond
must be the same element to share equally.
• Some examples of nonpolar covalent bonds.
– H2
– N2
68
Polar and Nonpolar Covalent Bonds
• Covalent bonds in which the electrons are not shared equally
are designated as polar covalent bonds
– Polar covalent bonds have an asymmetrical charge
distribution
• To be a polar covalent bond the two atoms involved in the
bond must have different electronegativities (具不同的電負
度)
• Some examples of polar covalent bonds.
• HF
H
F
Electroneg ativities
2.1
4.0


1.9
Difference  1.9 very polar bond
69
Polar and Nonpolar Covalent Bonds
• Shown below is an electron density map of HF
– Blue areas indicate low electron density
– Red areas indicate high electron density
• Polar molecules have a separation of centers of negative and
positive charge, an asymmetric charge distribution.
70
Polar and Nonpolar Covalent Bonds
• Shown below are electron density maps of the hydrogen
halides.
– Notice that the charge separation decreases as we move
from HF to HI.
71
Polar and Nonpolar Covalent Bonds
• Polar molecules can be attracted by magnetic and electric
fields.
72
Dipole Moments電偶極矩
• Molecules whose centers of positive and negative charge do not
coincide同位,重疊, have an asymmetric charge distribution, and
are polar.
– These molecules have a dipole moment.
• The dipole moment has the symbol .
•  is the product of the distance, d, separating charges of equal
magnitude and opposite sign, and the magnitude of the charge, q.
73
Dipole Moments 電偶極矩
• Molecules that have a small separation of charge have
a small .
• Molecules that have a large separation of charge have
a large .
• For example, HF and HI:

 H - F


-
1.91 Debye units
 H -I

-
0.38 Debye units
74
Dipole Moments
•
There are some nonpolar molecules that have
polar bonds.
• There are two conditions that must be true for a
molecule to be polar.
– There must be at least one polar bond present
or one lone pair of electrons.
– The polar bonds, if there are more than one,
and lone pairs must be arranged so that their
dipole moments do not cancel one another.
75
The Continuous Range of Bonding Types
• Covalent and ionic bonding represent two extremes.
1. In pure covalent bonds electrons are equally shared by
the atoms.
2. In pure ionic bonds electrons are completely lost or
gained by one of the atoms.
• Most compounds fall somewhere between these two
extremes.
76
Continuous Range of
Bonding Types
• All bonds have some ionic and some covalent character.
– For example, HI is about 17% ionic
• The greater the electronegativity differences the more polar
the bond.
77