CHAPTER 8

CHAPTER 8

Bonding and Molecular Structure

2

Introduction

• Bonds: Attractive forces that hold atoms together in compounds

• Valence Electrons: the outermost electrons– -These e- are involved in bonding

3

Electrons are divided between core and valence electrons

B 1sB 1s22 2s 2s22 2p 2p11

Core = [He]Core = [He] , valence = 2s , valence = 2s22 2p 2p11

Br [Ar] 3dBr [Ar] 3d1010 4s 4s22 4p 4p55

Core = [Ar] 3dCore = [Ar] 3d1010 , valence = 4s , valence = 4s22 4p 4p55

Valence Electrons

4

-The number of valence electrons of a main group atom is the Group number

-For Groups IA-IVA, number of bonding -For Groups IA-IVA, number of bonding (unpaired) (unpaired) electrons is equal to the group numberelectrons is equal to the group number-For Groups VA -VIIA, number of bonding (unpaired) -For Groups VA -VIIA, number of bonding (unpaired)

electrons is equal to 8 - group numberelectrons is equal to 8 - group number

Valence Electrons

5

-Except for H (and sometimes atoms of -Except for H (and sometimes atoms of the 3the 3rdrd group and higher) group and higher)

-The total number of valence electrons -The total number of valence electrons around a given atom in a molecule will around a given atom in a molecule will be eight: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)

Valence Electrons

6

Lewis Dot Formulas of Atoms

H He

Li Be B C N O F Ne

IA IIA IIIA IVA VA VIA VIIA VIIIA

7

Ionic BondingAn ion is an atom or a group of

atoms possessing a net electrical charge

• -cations: positive (+) ions• These atoms have lost 1 or more

electrons1. -anions: negative (-) ions

• These atoms have gained 1 or more electrons

8

Formation of Ionic Compounds• Monatomic ions consist of one atom

– Examples:• Na+, Ca2+, Al3+ - cations• Cl-, O2-, N3- -anions

• Polyatomic ions contain more than one atom

Examples:• NH4

+ - cation• NO2

-,CO32-, SO4

2- - anions

9

Formation of Ionic Compounds• General trend:

– metals become isoelectronic with the preceding noble gas electron configuration

– nonmetals become isoelectronic with the following noble gas electron configuration

10

Formation of Ionic Compounds

• Reaction of Group IA Metals with Group VIIA Nonmetals

point melting C842an with gas solid

solid whiteyellow silver

LiF 2 F Li 2 nometal 17 -G metal 1-G

o

(s)2(g)(s)

11


• 1s 2s 2p Li F

These atoms form ions with these configurations.

Li+ same configuration as [He] F- same configuration as

[Ne]Li + F.

..

.... . Li

+ F[ ]...... ..

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Formation of Ionic Compounds• In general:

the reaction of IA metals and VIIA nonmetals:2 M(s) + X2 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 ns npM M+ __ __ __ __X X-

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reaction of IIA metals with VIIA nonmetals:

Be(s) + F2(g) BeF2(g)

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The valence electrons in these two elements react like:

2s 2p 2s 2p Be [He] Be2+ __ __ __

__F [He] F-

Lewis dot structure representation:

15


The remainder of the IIA metals and VIIA nonmetals react similarly:

M(s) + X2 MX2

M can be any of the metals Be to BaX can be any of the nonmetals F to I

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Formation of Ionic CompoundsFor the reaction of IA metals with

VIA nonmetals:

-2s22(g)(s) O Li2O Li4

Draw the valence electronic configurations for Li, O, and their appropriate ions

17

Formation of Ionic Compounds• Draw the electronic configurations

for Li, O, and their appropriate ionsYou do it!You do it!

2s 2p 2s 2p Li [He] Li1+

O [He] O2-

Draw the Lewis dot formula representation of this reaction

18

Formation of Ionic CompoundsSimple Binary Ionic Compounds Table

• Reacting Groups General Formula Example IA + VIIA MX NaF IIA + VIIA MX2 BaCl2IIIA + VIIA MX3 AlF3 IA + VIA M2X Na2O IIA + VIA MX BaOIIIA + VIA M2X3 Al2S3

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• Reacting Groups General Formula Example IA + VA M3X Na3N IIA + VA M3X2 Mg3P2IIIA + VA MX AlN

-H forms ionic compounds when bound to metals (IA and IIA metals

For example: LiH, KH, CaH2, and BaH2

-When H is bound to nonmetals, the compounds are covalent in nature

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Formation of Covalent Bonds• -potential energy of an H2 molecule as

a function of the distance between the two H atoms

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Covalent Bonding• Atoms share electrons

• If the atoms share: • 2 electrons a single covalent bond is

formed• 4 electrons - a double covalent bond• 6 electrons - a triple covalent bond

Atoms have a lower potential energy when bound…this is a more favorable situation (why?)

22

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

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Writing Lewis Formulas

octet rule: representative elements usually attain stable noble gas electron configurations (8 valence e-) in most compounds

You must distinguish the difference between: – -bonding electrons and nonbonding

electrons -shared (paired) and unshared

(unpaired) electrons

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Formation of Covalent Bonds

• Lewis dot structures:

• 1. H2 molecule formation:

2. HCl molecule formation:

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Lewis Structures• Homonuclear diatomic molecules

– 1. Two atoms of the same element, H2:H HorH H..

2. Fluorine, F2:

3. Nitrogen, N2:

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Lewis Structuresheteronuclear diatomic molecules

1. hydrogen fluoride, HF

2. hydrogen chloride, HCl

3. hydrogen bromide, HBror ··H Cl

··

··H Cl..

·· ····

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Lewis Structures

• Water, H2O

•Ammonia molecule , NH3

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Lewis Structures• Polyatomic ions:• ammonium ion NH4

+

Notice that the N-atom in this molecule has eight electrons around them (H does not)

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Writing Lewis Formulas

• Sulfite ion, SO32-.

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Double and even triple bonds are commonly observed for C, N, P, O, and S

••O OC

•• ••

••

HH22COCO

SOSO33

CC22FF44

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Lewis Structures

• Example: Write Lewis dot and dash formulas for sulfur trioxide, SO3

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Resonance• There are three possible structures for SO3:

O S

O

O·· ···· ·· ··

······

OS

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

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ResonanceResonance is a flawed method of

representing molecules– -There are no single or double bonds in

SO3

SO O

O

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Sulfur Dioxide, SOSulfur Dioxide, SO221. Central atom =1. Central atom =2. Valence electrons = ___2. Valence electrons = ___ or ___ pairsor ___ pairs

4. Form double bond so 4. Form double bond so that S has an octet — but that S has an octet — but note that there are two note that there are two ways of doing this.ways of doing this.

3. Write the Lewis 3. Write the Lewis structurestructure

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Limitations of the Octet Rule • There are some molecules that violate the

octet rule:1. - Be2. - Group IIIA3. -Odd number of total electrons.4. -Central element must have a share of more

than 8 valence electrons to accommodate all of the substituents. (i.e. S and P)

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Limitations of the Octet Rule• Example: Write Lewis formula for

BBr3.

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Sulfur Tetrafluoride, SFSulfur Tetrafluoride, SF44

Central atom = Central atom = Valence electrons = ___ or ___ Valence electrons = ___ or ___

pairs.pairs.Form sigma bonds and Form sigma bonds and

distribute electron pairs.distribute electron pairs.

F

••

••

••

FF

S••

••••

••

•• F

••

••

••

••

•• 5 pairs around the S 5 pairs around the S atom. A common atom. A common occurrence outside the occurrence outside the 2nd period. 2nd period.

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Limitations of the Octet Rule• Example: Write dot structures for

AsF5.

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• 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)

Formal Atomic Charges

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Formal Charge

CO CO22

. .

. .

. .

. .

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Thiocyanate Ion, SCN-

••

•••S NC

•••

••

•••S NC

•••

•••••S NC

•••

Which is the most stable resonance form?Which is the most stable resonance form?

Formal Charge

42

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

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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

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VSEPR Theory1. Two regions of high electron

density around the central atom.

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VSEPR Theory2. Three regions of high electron density around the

central atom.

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VSEPR Theory3. Four regions of high electron density around the

central atom.

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VSEPR Theory4. Five regions of high electron

density around the central atom.

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VSEPR Theory5. Six regions of high electron density around

the central atom.

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VSEPR Theory

1. Electronic geometry(family): locations of regions of electron density around the central atom(s)

2.2. Molecular geometry:Molecular geometry: arrangement of atoms around the central atom(s)

Electron pairs are not used in the molecular geometry determination

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VSEPR TheoryLone pairs (unshared pairs) of electrons require

more volume than shared pairs– -there is an ordering of repulsions of lone electrons

around central atomCriteria 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 Molecular Shapes and BondingBonding

• Symbolism:A = central atomB = bonding pairs around central atomU = lone pairs around central atom

• For example:AB3U designates that there are 3 bonding

pairs and 1 lone pair around the central atom

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Linear Electronic Geometry: AB2

Some examples of molecules with this geometry:BeCl2, BeBr2, BeI2, HgCl2, CdCl2

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Trigonal Planar Electronic Geometry: AB3

Some examples of molecules with this geometry are:

BF3, BCl3

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Tetrahedral Electronic Geometry: AB4

Some examples of molecules with this geometry are: CH4, CF4, CCl4, SiH4, SiF4

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VSEPR Theory• An example of a molecule that has the

same electronic and molecular geometries is methane (CH4)– -Electronic and molecular geometries are

tetrahedral

H

CHH

H

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Tetrahedral Electronic Geometry: AB4

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Tetrahedral Electronic Geometry: AB3U

Some examples of molecules with this geometry are: NH3, NF3, PH3, PCl3, AsH3

– -trigonal pyramidal-electronic and molecular geometries are different.

. .

107.5°

. .

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• Some examples of molecules with this geometry are: H2O, OF2, H2S– -bent

-electronic and molecular geometries are different

Tetrahedral Electronic Geometry: AB2U2

104.5°

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VSEPR Theory• An example of a molecule that has

different electronic and molecular geometries is water (H2O)– -Electronic geometry is tetrahedral– -Molecular geometry is bent or angular

H

CHH

H

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Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and

AB2U3

Some examples of molecules with this geometry are: PF5, AsF5, PCl5

axial

axial

equatorial

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Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3If lone pairs are incorporated into the trigonal

bipyramidal structure, there are three possible new shapes:

1. One lone pair - Seesaw shape2. Two lone pairs - T-shape3. Three lone pairs – linear

The lone pairs occupy equatorial positions first: -they are 120o from each other

-90o from the axial positions– Results in decreased repulsions compared to

lone pair in axial positionaxial

axial

equatorial

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AB2U3

• AB4U molecules have:1. trigonal bipyramid electronic

geometry 2. seesaw shaped molecular geometry 3. polar

• One example of an AB4U molecule is SF4

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AB2U3

H

CHH

H

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Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3

• AB3U2 molecules have: 1. 1. trigonal bipyramid electronic

geometry 2. T-shaped molecular geometry 3. polar

• One example of an AB3U2 molecule is

IF3

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AB2U3

H

CHH

H

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AB2U3• AB2U3 molecules have:

1.trigonal bipyramid electronic geometry

2.linear molecular geometry 3.nonpolar

• One example of an AB3U2 molecule is BrF2

-

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AB2U3

H

CHH

H

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Octahedral Electronic Geometry: AB6, AB5U, and AB4U2

• Some examples of molecules with this geometry are: SF6, SeF6, SCl6, etc.

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If lone pairs are incorporated into the octahedral structure, there are two possible new shapes:1. One lone pair - square pyramidal2. Two lone pairs - square planar

The lone pairs occupy any position because they are all 90o 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

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• AB5U molecules have:1.octahedral electronic geometry 2.Square pyramidal molecular

geometry 3.polar.

• One example of an AB4U molecule is IF5

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• AB4U2 molecules have:1.octahedral electronic geometry 2.square planar molecular geometry 3.and are nonpolar.

• One example of an AB4U2 molecule is

XeF4

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Polarity and Polarity and ElectronegativityElectronegativityFigure 8.11Figure 8.11

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Dipole Moments

• For example, HF and HI:

units Debye 0.38 units Debye 1.91 I-H F-H

--

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Dipole Momentssome “nonpolar molecules” that have

polar bonds

Two conditions to be polar:1. 1. There must be at least one polar bond

present or one lone pair of electrons2. 2. the molecule must be nonsymmetric

Examples: water, CF4, CO2, NH3, NH4+

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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

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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········ ·· N N·· ··or H HorH H..

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Polar and Nonpolar Bonds• Polar covalent bonds: electrons are

not shared equally • -they have different electronegativities

H FElectronegativities: 2.1 4.0

Difference = 1.9 very polar bond

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Polar and Nonpolar Bonds• Compare HF to HI:

H IElectronegativities: 2.1 2.5

Difference = 0.4 slightly polar bond

more complicated geometries exist…

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• Three molecules with polar covalent bonds:

• -Each bond has one atom with a slight negative charge (-)

• -another with a slight positive charge (+)

Bond Polarity

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Polar or Nonpolar?Polar or Nonpolar?AB3 molecules: BF3, Cl2CO, and NH3

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Polar or Nonpolar?Polar or Nonpolar?CO2 and H2O

Which one is polar?

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CHCH44 … CCl … CCl44Polar or Not?Polar or Not?

• Only CH4 and CCl4 are NOT polar. These are the only two molecules that are “symmetrical.”

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Compounds Containing Double Bonds

• Ethene or ethylene, C2H4, is the simplest organic compound containing a double bond.– -has a double bond to obey octet ruleLewis Dot Formula

CCH

HH

HC C

H

H

H

H····

·· ·· ··

··or

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• What is the effect of bonding and structure on molecular properties?

Free rotation Free rotation around C–C single around C–C single bondbond

No rotation No rotation around C=C around C=C double bonddouble bond

and

Double Bonds

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Bond Order # of bonds between similar pairs of

atomsDouble bondDouble bond Single bondSingle bond

Triple Triple bondbond

AcrylonitrileAcrylonitrile

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Consider NO2-:

that typeof bound atoms of # Totaltype-one of bonds of # Total = orderBond

The N—O bond order = 1.5The N—O bond order = 1.5

O O O ON••

••••

••••

••••••••••

••••

••N

Bond Order

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Bond order is proportional to two important bond properties:

(a) bond strength(b) bond length

745 kJ745 kJ

414 kJ414 kJ 110 pm110 pm

123 pm123 pm

Bond Order

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the distance between the nuclei of two bonded atoms

Bond Length

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Bond LengthBond length depends on size of bonded atoms

H—FH—F

H—ClH—Cl

H—IH—I

Bond distances measured Bond distances measured in Angstrom units where 1 in Angstrom units where 1 ÅÅ = 10 = 10-2 -2 pm.pm.

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Bond length depends on bond order

Bond distances measured Bond distances measured in Angstrom units where 1 in Angstrom units where 1 ÅÅ = 10 = 10-2 -2 pm.pm.

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• Measure of the energy required to break a bond

• See Table 9.10• BOND STRENGTH (kJ/mol) H—H 436 KJ C—C 346 KJ C=C 602 KJ CC 835 KJ NN 945 KJ

The GREATER the number of bonds (bond order) the The GREATER the number of bonds (bond order) the HIGHER the bond strength and the SHORTER the HIGHER the bond strength and the SHORTER the bond.bond.

Bond Strength

CHAPTER 8

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