Transcript CHAPTER 8
CHAPTER 8
Bonding and Molecular Structure
Introduction
• Bonds: Attractive forces that hold atoms together in compounds • Valence Electrons: the outermost electrons – -These e are involved in bonding 2
Valence Electrons
Electrons are divided between core and valence electrons B 1s 2 2s 2 2p 1 Core = [He] , valence = 2s 2 2p 1 Br [Ar] 3d 10 4s 2 Core = [Ar] 3d 10 4p 5 , valence = 4s 2 4p 5 3
Valence Electrons
-The number of valence electrons of a main group atom is the Group number -For Groups IA-IVA, number of bonding (unpaired) electrons is equal to the group number -For Groups VA -VIIA, number of bonding (unpaired) electrons is equal to 8 - group number 4
Valence Electrons
-Except for H (and sometimes atoms of the 3 rd group and higher) -The total number of valence electrons around a given atom in a molecule will be eight:
OCTET RULE
- (with the exception of hydrogen) atoms in molecules prefer to be surrounded by 8 electrons (or have 4 bonds = 8 electrons) 5
Lewis Dot Formulas of Atoms
IA IIA IIIA IVA VA VIA VIIA VIIIA
H He Li Be B C N O F Ne
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Ionic Bonding
An
ion
is an atom or a group of atoms possessing a net electrical charge • -cations: positive (+) ions • These atoms have lost 1 or more electrons 1.
anions: negative (-) ions • These atoms have gained 1 or more electrons 7
Formation of Ionic Compounds
• • Monatomic – ions consist of one atom Examples: • • Na + , Ca 2+ , Al Cl , O 2 , N 3 3+ - cations -anions Polyatomic ions contain more than one atom Examples: • NH 4 + - cation • NO 2 ,CO 3 2 , SO 4 2 - anions 8
Formation of Ionic Compounds
• General trend: – metals become
isoelectronic
preceding noble gas electron configuration with the – nonmetals become
isoelectronic
with the following noble gas electron configuration 9
Formation of Ionic Compounds
• Reaction of Group IA Metals with Group VIIA Nonmetals
G 1 metal G 17 nometal 2 Li
(s)
F
2(g)
2 LiF
(s)
silver yellow white solid solid gas with an 842
o
C melting point
10
•
Formation of Ionic Compounds
Li 1s F 2s . 2p .
Li + F These atoms form ions with these configurations .
. ..
.
same configuration as [He] same configuration as [Ne] Li + [ ] 11
Formation of Ionic Compounds
• In general: the reaction of IA metals and VIIA nonmetals: 2 M (s) + X 2 2 MX (s) – – where M is the metals Li to Cs and X is the nonmetals F to I Electronically it looks like: ns np M M + X X ns np __ __ __ __ 12
Formation of Ionic Compounds
reaction of IIA metals with VIIA nonmetals: Be (s) + F 2(g) BeF 2(g) 13
Formation of Ionic Compounds
The valence electrons in these two elements react like: Be [He] F [He] 2s 2p __ __ __ 2s 2p Be 2+ __ __ __ __ F Lewis dot structure representation: 14
Formation of Ionic Compounds
The remainder of the IIA metals and VIIA nonmetals react similarly: M (s) + X 2 MX 2 M can be any of the metals Be to Ba X can be any of the nonmetals F to I 15
Formation of Ionic Compounds
For the reaction of IA metals with VIA nonmetals: 4 Li (s) O 2(g) 2 Li 2 O 2 Draw the valence electronic configurations for Li, O, and their appropriate ions 16
Formation of Ionic Compounds
• Draw the electronic configurations for Li, O, and their appropriate ions
You do it!
2s Li [He] 2p O [He] 2s 2p Li 1+ O 2 Draw the Lewis dot formula representation of this reaction 17
Formation of Ionic Compounds
• Simple Binary Ionic Compounds Table
Reacting Groups General Formula Example
IA + VIIA IIA + VIIA IIIA + VIIA IA + VIA IIA + VIA IIIA + VIA MX MX 2 MX 3 M 2 X MX M 2 X 3 NaF BaCl 2 AlF 3 Na 2 O BaO Al 2 S 3 18
Formation of Ionic Compounds
•
Reacting Groups General Formula
IA + VA IIA + VA IIIA + VA M 3 X M 3 X 2 MX
Example
Na 3 N Mg 3 P 2 AlN H forms ionic compounds when bound to metals (IA and IIA metals For example: LiH, KH, CaH 2 , and BaH 2 When H is bound to nonmetals, the compounds are covalent in nature 19
Formation of Covalent Bonds
• -potential energy of an H atoms 2 molecule as a function of the distance between the two H 20
Covalent Bonding
• Atoms share electrons • • • • If the atoms share: 2 4 6 electrons a single electrons - a electrons - a triple covalent bond is formed double covalent bond covalent bond Atoms have a lower potential energy when bound…this is a more favorable situation (why?) 21
Writing Lewis Formulas:
• • • • • • 1. Add the number of valence electrons for all the atoms that are present in the molecule 2. Add or subtract electrons based on the molecule’s (or ion’s) charge 3. Identify the central atom and draw a skeletal structure: – -the one that requires the most e- to complete octet – -the less electronegative 4. Place a bond between each atom (1 bond = 2 e-) 5. Fill in octet of outer atoms first 6. Finish by completing the octet of central atom – – if you run out of e- then multiple bonds must be created between the central atom and atoms bound to it 22
Writing Lewis Formulas
octet rule: representative elements usually attain stable noble gas electron configurations (8 valence e compounds ) in most You must distinguish the difference between: – -bonding electrons and nonbonding electrons -shared (paired) and unshared (unpaired) electrons 23
Formation of Covalent Bonds
• Lewis dot structures: • 1. H 2 molecule formation: 2. HCl molecule formation: 24
Lewis Structures
• Homonuclear diatomic molecules – 1. Two atoms of the same element , H 2 : or H H 2. Fluorine, F 2 : 3. Nitrogen, N 2 : 25
Lewis Structures
heteronuclear diatomic molecules 1. hydrogen fluoride, HF 2. hydrogen chloride, HCl ·· H Cl ·· · · or ·· H Cl · · ·· 3. hydrogen bromide, HBr 26
Lewis Structures
• Water, H 2 O • Ammonia molecule , NH 3 27
Lewis Structures
• • Polyatomic ions: ammonium ion NH 4 + Notice that the N-atom in this molecule has eight electrons around them (H does not) 28
Writing Lewis Formulas
• Sulfite ion, SO 3 2 .
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Double and even triple bonds are commonly observed for C, N, P, O, and S
H 2 CO SO 3 C 2 F 4
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Lewis Structures
• Example: Write Lewis dot and dash formulas for sulfur trioxide, SO 3 31
Resonance
• There are three possible structures for SO 3 : · · O ·· S · · O ·· · · ·· O ·· · · · · ·· O ·· S · · O · · ·· O ·· · · · · ·· O ·· S · · O ·· · · O ·· · · -Two or more Lewis formulas are necessary to show the bonding in a molecule use equivalent resonance structures to show the molecule’s structure -Double-headed arrows are used to indicate resonance formulas 32
Resonance
Resonance is a flawed method of representing molecules – -There are no single or double bonds in SO 3 O S O O 33
Sulfur Dioxide, SO
2 1. Central atom = 2. Valence electrons = ___ or ___ pairs 3. Write the Lewis structure 4. Form double bond so that S has an octet — but note that there are two ways of doing this.
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Limitations of the Octet Rule
• There are some molecules that violate the octet rule: 1.
- Be 2.
3.
4.
- Group IIIA -Odd number of total electrons.
-Central element must have a share of more than 8 valence electrons to accommodate all of the substituents. (i.e. S and P) 35
Limitations of the Octet Rule
• Example: Write Lewis formula for BBr 3 .
36
Sulfur Tetrafluoride, SF
4 Central atom = Valence electrons = ___ or ___ pairs.
Form sigma bonds and distribute electron pairs.
5 pairs around the S atom. A common occurrence outside the 2nd period.
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Limitations of the Octet Rule
• Example: Write dot structures for AsF 5 .
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Formal Atomic Charges
• Atoms in molecules often bear a charge (+ or -) • The predominant resonance structure of a molecule is the one with charges on atoms as close to 0 as possible • • • • Formal charge = Group number – 1/2 (# of bonding electrons) - (# of Lone electrons) = Group number – (# of bonds) – (# of Lone electrons) 39
Formal Charge
CO 2 . .
. .
. .
. .
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Formal Charge
Thiocyanate Ion, SCN • • • • •
S
•
C N
• • • • • •
S C
• •
N
• • • •
S C
• • •
N
• • •
Which is the most stable resonance form?
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Theories of Covalent Bonding
• • Valence Shell Electron Pair Repulsion Theory – Commonly designated as VSEPR – Principal originator • R. J. Gillespie in the 1950’s Valence Bond Theory (Chapter 9) – – Involves the use of hybridized atomic orbitals Principal originator • L. Pauling in the 1930’s & 40’s 42
VSEPR Theory
electron densities around the central atom are arranged as far apart as possible to minimize repulsions (why?) • • Five basic molecular shapes: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral 43
VSEPR Theory
1.
Two regions of high electron density around the central atom.
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VSEPR Theory
2.
Three regions of high electron density around the central atom.
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VSEPR Theory
3.
Four regions of high electron density around the central atom.
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VSEPR Theory
4.
Five regions of high electron density around the central atom.
47
VSEPR Theory
5.
Six regions of high electron density around the central atom.
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VSEPR Theory
1.
2.
Electronic geometry (family): locations of regions of electron density around the central atom(s) Molecular geometry: arrangement of atoms around the central atom(s) Electron pairs are not used in the molecular geometry determination 49
VSEPR Theory
Lone pairs (unshared pairs) of electrons require more volume than shared pairs – -there is an ordering of repulsions of lone electrons around central atom Criteria for the ordering of the repulsions: 1. Lone pair to lone pair is the strongest repulsion.
2. Lone pair to bonding pair is intermediate repulsion.
3. Bonding pair to bonding pair is weakest repulsion.
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Molecular Shapes and Bonding
• • Symbolism: A = central atom B = bonding pairs around central atom U = lone pairs around central atom For example: AB 3 U designates that there are 3 bonding pairs and 1 lone pair around the central atom 51
Linear Electronic Geometry: AB 2
Some examples of molecules with this geometry: BeCl 2 , BeBr 2 , BeI 2 , HgCl 2 , CdCl 2 52
Trigonal Planar Electronic Geometry: AB 3
Some examples of molecules with this geometry are: BF 3 , BCl 3 53
Tetrahedral Electronic Geometry: AB 4
Some examples of molecules with this geometry are: CH 4 , CF 4 , CCl 4 , SiH 4 , SiF 4 54
VSEPR Theory
• An example of a molecule that has the same electronic and molecular geometries is methane (CH 4 ) – -Electronic and molecular geometries are tetrahedral H H H C H 55
Tetrahedral Electronic Geometry: AB 4
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Tetrahedral Electronic Geometry: AB 3 U
Some examples of molecules with this geometry are: NH 3 , NF 3 , PH 3 , PCl 3 , AsH 3 – -trigonal pyramidal -electronic and molecular geometries are different.
. .
. .
107.5° 57
•
Tetrahedral Electronic Geometry: AB 2 U 2
Some examples of molecules with this geometry are: H 2 O , OF 2 , H 2 S – -bent -electronic and molecular geometries are different 104.5° 58
•
VSEPR Theory
An example of a molecule that has different electronic and molecular geometries is water (H 2 O) – -Electronic geometry is tetrahedral – -Molecular geometry is bent or angular H H H C H 59
Trigonal Bipyramidal Electronic Geometry: AB 5 , AB 4 U, AB 3 U 2 , and AB 2 U 3
Some examples of molecules with this geometry are: PF 5 , AsF 5 , PCl 5 axial equatorial axial 60
Trigonal Bipyramidal Electronic Geometry: AB 5 , AB 4 U, AB 3 U 2 , and AB 2 U 3
If lone pairs are incorporated into the trigonal bipyramidal structure, there are three possible new shapes: 1.
One lone pair - Seesaw shape 2.
3.
Two lone pairs - T-shape Three lone pairs – linear The lone pairs occupy equatorial positions first: -they are 120 o from each other – -90 o from the axial positions Results in decreased repulsions compared to lone pair in axial position axial equatorial 61 axial
Trigonal Bipyramidal Electronic Geometry: AB 5 , AB 4 U, AB 3 U2, and AB 2 U 3
• • AB 1.
2.
3.
4 U molecules have: trigonal bipyramid electronic geometry seesaw shaped molecular geometry polar One example of an AB 4 U molecule is SF 4 62
Trigonal Bipyramidal Electronic Geometry: AB 5 , AB 4 U, AB 3 U2, and AB 2 U 3
H H H C H 63
Trigonal Bipyramidal Electronic Geometry: AB 5 , AB 4 U , AB 3 U 2 , and AB 2 U 3
• • AB 1.
2.
3.
3 U 2 molecules have: 1.
trigonal bipyramid electronic geometry T-shaped molecular geometry polar One example of an AB 3 U 2 IF 3 molecule is 64
Trigonal Bipyramidal Electronic Geometry: AB 5 , AB 4 U , AB 3 U2, and AB 2 U 3
H H H C H 65
Trigonal Bipyramidal Electronic Geometry: AB 5 , AB 4 U, AB 3 U2 , and AB 2 U 3
• • AB 2 U 3 molecules have: 1.
trigonal bipyramid electronic geometry 2.
linear molecular geometry 3.
nonpolar One example of an AB 3 U 2 BrF 2 molecule is 66
Trigonal Bipyramidal Electronic Geometry: AB 5 , AB 4 U, AB 3 U2 , and AB 2 U 3
H H H C H 67
Octahedral Electronic Geometry: AB 6 , AB 5 U, and AB 4 U 2
• Some examples of molecules with this geometry are: SF 6 , SeF 6 , SCl 6 , etc.
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Octahedral Electronic Geometry: AB 6 , AB 5 U, and AB 4 U 2
If lone pairs are incorporated into the octahedral structure, there are two possible new shapes: 1.
One lone pair - square pyramidal 2.
Two lone pairs - square planar The lone pairs occupy any position because they are all 90 o from all bonds positions: – – -Additional lone pairs occupy the position 180º from the first set of lone pairs -This results in decreased repulsions compared to lone pairs in the other positions 69
Octahedral Electronic Geometry: AB 6 , AB 5 U, and AB 4 U 2
• AB 5 U molecules have: 1.
octahedral electronic geometry 2.
Square pyramidal molecular geometry 3.
polar .
• One example of an AB 4 U molecule is IF 5 70
Octahedral Electronic Geometry: AB 6 , AB 5 U , and AB 4 U 2
• AB 4 U 2 molecules have: 1.
octahedral electronic geometry 2.
square planar molecular geometry 3.and are nonpolar .
• One example of an AB 4 U 2 XeF 4 molecule is 71
Polarity and Electronegativity
Figure 8.11
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Dipole Moments
• For example, HF and HI: H F H I 1.91
Debye units 0.38
Debye units 73
Dipole Moments
some “nonpolar molecules” that have polar bonds Two conditions to be polar: 1.
2.
1. There must be at least one polar bond present or one lone pair of electrons 2. the molecule must be nonsymmetric Examples: water, CF 4 , CO 2 , NH 3 , NH 4 + 74
Polar Molecules
• Molecular geometry affects molecular polarity – -they either cancel or reinforce each other A B A linear molecule nonpolar A B A angular molecule polar 75
Polar and Nonpolar Bonds
• Covalent bonds in which the electrons are shared equally are designated as nonpolar covalent bonds – -Nonpolar covalent bonds have a symmetrical charge distribution (electron distribution) · · N · · · · · · N · · or · · N N · · or H H 76
Polar and Nonpolar Bonds
• • Polar covalent bonds: electrons are not shared equally -they have different electronegativities H Electronegativities: Difference = 1.9 very polar bond F 2.1 4.0
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Polar and Nonpolar Bonds
• Compare HF to HI: Electronegativities: H 2.1 2.5
Difference = 0.4 slightly polar bond I more complicated geometries exist… 78
Bond Polarity
• • • Three molecules with polar covalent bonds: -Each bond has one atom with a slight negative charge ( ) -another with a slight positive charge (+ ) 79
Polar or Nonpolar?
AB 3 molecules: BF 3 , Cl 2 CO, and NH 3 80
Polar or Nonpolar?
CO 2 and H 2 O Which one is polar?
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CH 4 … CCl 4 Polar or Not?
• Only CH 4 and CCl 4 are NOT polar. These are the only two molecules that are “symmetrical.” 82
Compounds Containing Double Bonds
• Ethene or ethylene, C 2 H 4 , is the simplest organic compound containing a double bond.
– -has a double bond to obey octet rule H H · · C Lewis Dot Formula H C H or H H C H C H 83
Double Bonds
• What is the effect of bonding and structure on molecular properties? s and p
Free rotation around C –C single bond No rotation around C=C double bond
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Bond Order
# of bonds between similar pairs of atoms
Double bond Single bond Acrylonitrile Triple bond
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Bond Order
Consider NO 2 :
•• •• ••• ••• •• •• The N —O bond order = 1.5
Bond
Total # of
order
= Total # of bonds atoms of bound one of type that type 86
Bond Order
Bond order is proportional to two important bond properties: (a) bond strength (b) bond length
414 kJ 110 pm 123 pm 745 kJ
87
Bond Length
the distance between the nuclei of two bonded atoms 88
Bond Length atoms
H —F H —Cl Bond distances measured in Angstrom units where 1 Å = 10 -2 pm.
H —I
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Bond length depends on bond order
Bond distances measured in Angstrom units where 1 Å = 10 -2 pm.
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Bond Strength
• Measure of the energy required to break a bond • • See Table 9.10
BOND H —H C —C C=C C C N N STRENGTH (kJ/mol) 436 KJ 346 KJ 602 KJ 835 KJ 945 KJ
the bond strength and the SHORTER the bond.