molecular geometry and bonding theories ap chemistry – ch 9 mr. christopherson

Molecular Geometry and Bonding Theories

AP Chemistry – Ch 9Mr. Christopherson

http://www.chem1.com/acad/webtext/chembond/index.html

Bonding Theories & Geometry

• Molecular Geometry (shapes)• VSEPR Theory• Lewis Structures• Molecular Polarity (dipoles)• Covalent Bonds• Hybridization• Ionic Bonds

molecular formula

structural formula

molecular shape

ball-and-stick model

CH4 C

H

H

HH

H

H

H

H

109.5o

C

tetrahedrontetrahedralshape ofmethane

CH

H

H

H

109.5o109.5o

TetrahedronTetrahedron

CentralAtom

CentralAtom

SubstituentsSubstituents

Methane, CH4Methane, CH4

Tetrahedralgeometry

Tetrahedralgeometry

Methane, CH4Methane, CH4

Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

Methane & Carbon Tetrachloridemolecular formula

structural formula

molecular shape

ball-and-stick model

CH4 C

H

H

HH

H

H

H

H

109.5o

C

CCl4

space-filling model

C

Cl

Cl

ClCl

Molecular Geometry

H

H

H

H

109.5o

C

Linear Trigonal planar

Tetrahedral

Trigonal pyramidalBent

109.5o

107.3o104.5o

H2O CH4 AsCl3 AsF5 BeH2 BF3 CO2

180o

A Lone PairA Lone Pear

C109.5o

H

HHH

N107o HH

H

..

O104.5o H

H

..

..

CH4, methane NH3, ammonia H2O, water

..

O

O

O

lone pairelectrons

OOO

O3, ozone

Molecular ShapesThree atoms (AB2) Four atoms (AB3)

Five atoms (AB4)

Six atoms (AB5)

Seven atoms (AB6)

• Linear (180o)• Bent

• Trigonal planar (120o)• Trigonal pyramidal• T-shaped

• Tetrahedral (109.47o)• Square planar• Seesaw

• Trigonal bipyramidal (BeABe, 120o) & (BeABa, 90o)• Square pyramidal

• Octahedral

B BA

B

B

A

B

linear trigonal planarB

A

BB

B

tetrahedral

A Be

Be

Be

Ba

Ba

TrigonalbipyramidalB

B

B

B

B

BA

Bailar, Moeller, Kleinberg, Guss, Castellion, Metz, Chemistry, 1984, page 313.

Bonding and Shape of Molecules

Number of Bonds

Number of Unshared Pairs Shape Examples

2

3

4

3

2

0

0

0

1

2

Linear

Trigonal planar

Tetrahedral

Pyramidal

Bent

BeCl2

BF3

CH4, SiCl4

NH3, PCl3

H2O, H2S, SCl2

-Be-

B

C

N

:

O

:

:

CovalentStructure

http://upload.wikimedia.org/wikipedia/commons/1/11/Sp2-Orbital.svg

http://upload.wikimedia.org/wikipedia/commons/9/9f/Sp3-Orbital.svg

Molecular ShapesAB2

Linear

AB3

Trigonal planar AB4

Tetrahedral

AB5

Trigonal bipyramidal

AB6

Octahedral

AB2EAngular or Bent AB3E

Trigonalpyramidal

AB2E2

Angular or Bent

AB4EIrregular tetrahedral(see saw)

AB3E2

T-shaped

AB2E3

Linear

AB6ESquare pyramidal

AB5E2

Square planar

Valence

Shell

Electron

Pair

Repulsion

Theory

Planar triangular

Tetrahedral


Octahedral

Valence

Shell

Electron

Pair

Repulsion

Theory

Planar triangular

Tetrahedral


Octahedral

http://upload.wikimedia.org/wikipedia/commons/6/6c/AX3E0-3D-balls.png

http://upload.wikimedia.org/wikipedia/commons/1/1a/Octahedral-3D-balls.png

http://upload.wikimedia.org/wikipedia/commons/3/3a/Trigonal-bipyramidal-3D-balls.png

......

The VSEPR Model

O OC

Linear

The Shapes of Some Simple ABn Molecules

O OS

BentO O

S

O

Trigonalplanar

FF

F

N

Trigonalpyramidal

T-shaped Squareplanar

F FCl

F

F F

Xe

F FF

F

FP

F

F

Trigonalbipyramidal

Octahedral

FF

F

S

F

F

F

SF6

SO2

Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 305

Molecular ShapesAB2

Linear

AB3

Trigonal planarAB4

Tetrahedral

AB5


AB6

Octahedral

AB2EAngular or Bent

AB3ETrigonal

pyramidal

AB2E2

Angular or Bent

AB4EIrregular tetrahedral(see saw)

AB3E2

T-shaped

AB2E3

Linear

AB5ESquare pyramidal

AB4E2

Square planar

Geometry of Covalent Molecules ABn, and ABnEm

AB2

AB2EAB2E2

AB2E3

AB3

AB3E

AB3E2

AB4

AB4E

AB4E2

AB5

AB5EAB6

222233

34

4

45

56

012301

20

1

20

10

LinearTrigonal planarTetrahedralTrigonal bipyramidalTrigonal planarTetrahedral

Triangular bipyramidalTetrahedral

Triangular bipyramidal

OctahedralTriangular bipyramidal

OctahedralOctahedral

LinearAngular, or bentAngular, or bentLinearTrigonal planarTriangular pyramidal

T-shapedTetrahedral

Irregular tetrahedral (or “see-saw”)Square planarTriangular bipyramidal

Square pyramidalOctahedral

CdBr2

SnCl2, PbI2

OH2, OF2, SCl2, TeI2

XeF2

BCl3, BF3, GaI3

NH3, NF3, PCl3, AsBr3

ClF3, BrF3

CH4, SiCl4, SnBr4, ZrI4

SF4, SeCl4, TeBr4

XeF4

PF5, PCl5(g), SbF5

ClF3, BrF3, IF5

SF6, SeF6, Te(OH)6, MoF6

TypeFormula

Shared Electron

Pairs

Unshared Electron

Pairs

IdealGeometry

ObservedMolecular Shape Examples

Bailar, Moeller, Kleinberg, Guss, Castellion, Metz, Chemistry, 1984, page 317.

Predicting the Geometry of Molecules

• Lewis electron-pair approach predicts number and types of bonds between the atoms in a substance and indicates which atoms have lone pairs of electrons but gives no information about the actual arrangement of atoms in space

• Valence-shell electron-pair repulsion (VSEPR) model predicts the shapes of many molecules and polyatomic ions but provides no information about bond lengths or the presence of multiple bonds

Introduction to Lewis Structures

Lewis dot symbols

1. Used for predicting the number of bonds formed

by most elements in their compounds

2. Consists of the chemical symbol for an element surrounded by dots that represent its valence

electrons

3. A single electron is represented as a single dot

Lewis Structures

1) Count up total number of valence electrons

2) Connect all atoms with single bonds

- “multiple” atoms usually on outside

- “single” atoms usually in center;

C always in center,

H always on outside.

3) Complete octets on exterior atoms (not H, though)

4) Check

- valence electrons math with Step 1

- all atoms (except H) have an octet; if not, try multiple bonds

- any extra electrons? Put on central atom

Molecules with Expanded Valence ShellsAtoms that have expanded octets have AB5 (trigonal bipyramidal)

or AB6 (octahedral) electron domain geometries.

Trigonal bipyramidal structures have a plane containing three electron pairs.

• The fourth and fifth electron pairs are located above and below this plane.• In this structure two trigonal pyramids share a base.

For octahedral structures, there is a plane containing four electron pairs.

• Similarly, the fifth and sixth electron pairs are located above and below this plane.• Two square pyramids share a base.

F

F

FP

F

F

FF

F

S

F

F

F

Trigonal Bipyramid F

F

FP

F

F• The three electron pairs in the plane are called equatorial.

• The two electron pairs above and below this plane are called axial.

• The axial electron pairs are 180o apart and 90o from to the equatorial electrons.

• The equatorial electron pairs are 120o apart.

• To minimize electron-electron repulsions, nonbonding pairs are always placed in equatorial positions, and bonding pairs in either axial or equatorial positions.

Octahedron• The four electron pairs in the plane are 90o to each other.

• The remaining two electron pairs are 180o apart and 90o from the electrons in the plane.

• Because of the symmetry of the system, each position is equivalent.

• The equatorial electron pairs are 120o apart.

• If we have five bonding pairs and one nonbonding pair, it doesn’t matter where the nonbonding pair is placed.

The molecular geometry is square pyramidal.

• If two nonbonding pairs are present, the repulsions are minimized by pointing them toward opposite sides of the octahedron.

The molecular geometry is square planar.

FF

F

S

F

F

F

F F

Xe

F F

Electron-Domain GeometriesNumber of Electron Domains

Arrangement ofElectron Domains

Electron-DomainGeometry

Predicted Bond Angles

2

3

4

5

6

Linear

Trigonalplanar

Tetrahedral

Trigonal-bipyramidal

Octahedral

180o

120o

109.5o

120o

90o

90o

A Be

Be

Be

Ba

Ba

B

B

B

B

B

BA

B

B

A

B

B

A

BB

B

B BA

Number of electron domains

Electron-domain geometry

Predicted bond angles

TetrahedralTrigonalplanar Tetrahedral

109.5o 120o 109.5o

C C OH H

H

H O

4 3 4

Acetic Acid, CH3COOH


Hybridization of central atom sp3 sp2 none

http://upload.wikimedia.org/wikipedia/commons/e/ef/Acetic-acid-3D-balls.png

Intermolecular Forces

Ion-ion (ionic bonds)

Ion-dipole

Dipole-dipole

Hydrogen bonding

London dispersion forces

+ −− +

+ − + −

HO

H

HO

H

H

O

H

−+

London Dispersion Forces

• London dispersion forces are created when on molecule with a temporarily dipole causes another to become temporarily polar.

+

−

+

− +

−

Molecular Polarity

Molecular Structure

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Electronegativity

+ – 0 0

H Cl H H

Na

Ionic vs. Covalent• Ionic compounds form repeating units.• Covalent compounds form distinct molecules.• Consider adding to NaCl(s) vs. H2O(s):

HO

H Cl

Cl

Cl

ClH

OH

H

O H

• NaCl: atoms of Cl and Na can add individually forming a compound with million of atoms.

• H2O: O and H cannot add individually, instead molecules of H2O form the basic unit.

Na Na

Na

Cl ClNa Na

Holding it togetherQ: Consider a glass of water.

Why do molecules of water stay together?A: There must be attractive forces.

Intramolecular forces occur between atoms

Intermolecular forces occur between molecules

• Intermolecular forces are not considered in ionic bonding because there are no molecules.

• The type of intramolecular bond determines the type of intermolecular force.

Intramolecular forces are much

stronger

I’m not stealing, I’m sharing unequally• We described ionic bonds as stealing electrons• In fact, all bonds share – equally or unequally.• Note how bonding electrons spend their time:

• Bonding electrons are shared in each compound, but are NOT always shared equally.

• The greek symbol indicates “partial charge”.

H2 HCl LiCl

+ –0 0 + –

covalent (non-polar) polar covalent ionic

H H H Cl [Li]+[ Cl ]–

Dipole Moment

• Direction of the polar bond in a molecule.• Arrow points toward the more

electronegative atom.

H Cl+ -


Dipole-induced dipole attraction

Dipole-induced dipole attraction

The attraction between a dipole and an induced

dipole.

The attraction between a dipole and an induced

dipole.

Oxygen, O2Oxygen, O2

Oxygen, O2Oxygen, O2

NonpolarNonpolar

Water, H2OWater, H2O

Water, H2OWater, H2O

d-d-++

d-d-++

d-d-++

++--

d-d-++

++ --

++ --d-d-++

++ --

++

induced dipole

induced dipole

DipoleDipole

d-d-

++ --d-d-++

d-d-++

++ --

d-d-++

++--

d-d-++

++--

++d-d-

++--

++d-d-

induced dipole

induced dipoleDipoleDipole

PolarPolar NonpolarNonpolar


Determining Molecular Polarity

• Depends on:– dipole moments– molecular shape


+–

+ –

+ –

+ –

H Cl

+ –

BF3

F

F F

B


• Nonpolar Molecules– Dipole moments are symmetrical and cancel

out.


H2OH H

O


• Polar Molecules– Dipole moments are asymmetrical and don’t

cancel .

netdipolemoment


CHCl3

H

Cl ClCl


• Therefore, polar molecules have...– asymmetrical shape (lone pairs) or – asymmetrical atoms

netdipolemoment


Dipole Moment

Nonpolar

Polar

....

H H

O

C OO

Bond dipoles

Overall dipole moment = 0

Bond dipoles

Overall dipole moment

The overall dipole moment of a moleculeis the sum of its bond dipoles. In CO2 thebond dipoles are equal in magnitude butexactly opposite each other. The overall dipole moment is zero.

In H2O the bond dipoles are also equal inmagnitude but do not exactly oppose eachother. The molecule has a nonzero overall dipole moment.

221

dqqk

F Coulomb’s lawm = Q r Dipole moment, m


......

Polar Bonds

H Cl

Polar

A molecule has a zero dipole moment because their dipoles cancel one another.

H HO

PolarF F

B

F

Nonpolar

HH

H

N

Polar

Polar Nonpolar

F FCl

F

F F

Xe

F FCl

ClC

Cl

Nonpolar Polar

Cl

HC

Cl

H

H

HF HCl HBr HI

Mark Wirtz, Edward Ehrat, David L. Cedeno*

How is the electron density distributed in these different molecules?

Based on your comparison of the electron density distributions, which molecule should have the most polar bond, and which one the least polar?

Arrange the molecules in increasing order of polarity.

CH3Cl CH2Cl2 CHCl3 CCl4


Describe how is the electron density distributed in these different molecules? Based on your comparison of the electron density distributions, which molecule(s) should be the most polar, and which one(s) the least polar?

Arrange the molecules in increasing order of polarity.

Benzene NO3-

Nitrobenzene


2s 2p (x, y, z) carbon


How does H2 form?

+ +

The nuclei repel

But they are attracted to electrons

They share the electrons

Hydrogen Bond Formation

0.74 A

- 436

0

H – H distance

Ene

rgy

(KJ/

mol

)


no interaction

increasedattraction

balanced attraction& repulsion

increasedrepulsion

Potential Energy Diagram - Attraction vs. Repulsion

(internuclear distance)

Covalent bonds

• Nonmetals hold onto their valence electrons.• They can’t give away electrons to bond.• Still want noble gas configuration.

• Get it by sharing valence electrons with each other.

• By sharing both atoms get to count the electrons toward noble gas configuration.

1s22s22p63s23p6…eight valence electrons (stable octet)

Covalent bonding

• Fluorine has seven valence electrons

F

•A second atom also has seven

F

By sharing electrons…both end with full orbitals

8 Valence electrons

8 Valence electrons

Single Covalent Bond

• A sharing of two valence electrons.• Only nonmetals and Hydrogen.• Different from an ionic bond because they

actually form molecules.• Two specific atoms are joined.• In an ionic solid you can’t tell which atom

the electrons moved from or to.

Sigma bonding orbitals

• From s orbitals on separate atoms

+ +

s orbital s orbital

+ ++ +

Sigma bondingmolecular orbital

Sigma bonding orbitals

• From p orbitals on separate atoms

p orbital p orbital

Sigma bondingmolecular orbital

Pi bonding orbitals

• P orbitals on separate atoms

Pi bondingmolecular orbital

Sigma and pi bonds

• All single bonds are sigma bonds• A double bond is one sigma and one pi

bond• A triple bond is one sigma and two pi

bonds.

Atomic Orbitals and Bonding• Bonds between atoms are formed by electron pairs in

overlapping atomic orbitals

1s 1s:E 1s• Example: H2 (H-H)

– Use 1s orbitals for bonding

• Example: H2O– From VSEPR: bent, 104.5°

angle between H atoms– Use two 2p orbitals for bonding?

E2s

2p

90° How do we explain the structure predicted by VSEPR

using atomic orbitals?2p

2p

1s

1s

Overlapping OrbitalsDraw orbital diagrams for F + F, H + O, Li + F

1s 2s 2p

1s2s2p

1s 2s 2p

1s

1s

F2

H2O

1s 2s 1s2s2p

LiF is ionic (metal + non-metal)

FLi

electron transfer

1+ 1-

e-

3p+

lithium atomLi

e-

loss of one valence

electron

e-

e-

lithium ionLi+

3p+e-

e-

9p+

fluorine atomF

e-

e-

e-

e-

e-

e-

e- e-

e-

gain of one valence

electron

fluoride ionF1-

10p+e-

e-

e-

e-e-

e-

e-

e-

e-

e-

e-

Formation of Cation

3p+

lithium atomLi

e-

loss of one valence

electron

e-

e-

lithium ionLi+

3p+e-

e-

e-

Formation of Anion

9p+

fluorine atomF

e-

e-

e-

e-

e-

e-

e- e-

e-

e-

gain of one valence

electron

fluoride ionF1-

10p+e-

e-

e-

e-e-

e-

e-

e-

e-

e-

Formation of Ionic Bond

fluoride ionF1-

9p+e-

e-

e-

e-e-

e-

e-

e-

e-

e-

lithium ionLi+

3p+e-

e-

http://en.wikipedia.org/wiki/Ionic_bond

Be

H

H

BeH2

s p

First, the formation of BeH2 using pure s and p orbitals.

The formation of BeH2 using hybridized orbitals.

atomic orbitals atomic orbitals

Be

s p

Be H

H

s p

atomic orbitals

hybrid orbitals

No overlap = no bond!

sp p

Be HH

All hybridized bonds have equal strength and have orbitals with identical energies.

BeH2Be

Be = 1s22s2

Hybrid OrbitalsGround-state Be atom

1s 2s 2p

1s 2s 2p

Be atom with one electron “promoted”

s

px py pz

sp

hybrid orbitals

Ene

rgy

hybridize

s orbital p orbital

two sp hybrid orbitals sp hybrid orbitals shown together(large lobes only)

1s sp 2p

Be atom of BeH2 orbital diagram

H HBe

n = 1

n = 2

Hybrid Orbitals

2s 2p

Ground-state B atom

s

px py pzEne

rgy

sp2 2p

B atom of BH3 orbital diagram

hybridize

s orbital

2s 2p

B atom with one electron “promoted”

sp2

hybrid orbitals

p orbitals sp2 hybrid orbitals shown together

(large lobes only)three sps hybrid orbitals

H

H

HB

Hybridization…the blending of orbitals

Valence bond theory is based on two assumptions:

1. The strength of a covalent bond is proportional to the amount of overlap between atomic orbitals; the greater

the overlap, the more stable the bond.

2. An atom can use different combinations of atomic orbitals to maximize the overlap of orbitals used by bonded

atoms.

We have studied electron configuration notation and the sharing of electrons in the formation of covalentbonds.

Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it.

Lets look at amolecule of methane, CH4.

What is the expected orbital notation of carbonin its ground state?

(Hint: How many unpaired electrons does this carbon atom have available for bonding?)

Can you see a problem with this?

Carbon ground state configuration

1s2s

2p

You should conclude that carbon only has TWO electrons available for bonding. That is not enough!

How does carbon overcome this problem so thatit may form four bonds?

The first thought that chemists had was that carbon promotes one of its 2s electrons…

…to the empty 2p orbital.

Carbon’s Empty Orbital

1s2s

2p

1s2s

2p

1s2s

2p

Non-hybridized orbital hybridized orbital

However, they quickly recognized a problem with such an arrangement…

Three of the carbon-hydrogen bonds would involvean electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.

A Problem Arises

1s2s

2p

1s 1s1s 1s

Unequal bond energy

But what about the fourth bond…?

The fourth bond is between a 2s electron from thecarbon and the lone 1s hydrogen electron.

Such a bond would have slightly less energy than the other bonds in a methane molecule.

Unequal bond energy #2

1s2s

2p

1s 1s1s 1s

This bond would be slightly different in character than the other three bonds in methane.

This difference would be measurable to a chemistby determining the bond length and bond energy.

But is this what they observe?

The simple answer is, “No”.

Chemists have proposed an explanation – they call ithybridization.

Hybridization is the combining of two or more orbitalsof nearly equal energy within the same atom into orbitals of equal energy.

Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane.

Enter Hybridization

http://upload.wikimedia.org/wikipedia/commons/9/9f/Sp3-Orbital.svg

In the case of methane, they call the hybridization sp3, meaning that an s orbital is combined with threep orbitals to create four equal hybrid orbitals.

These new orbitals have slightly MORE energy thanthe 2s orbital…

… and slightly LESS energy than the 2p orbitals.

sp3 Hybrid Orbitals

http://upload.wikimedia.org/wikipedia/en/b/bb/Ch4_hybridization.svg

http://upload.wikimedia.org/wikipedia/commons/5/55/Ch4-structure.png

s

px py pz

Carbon 1s22s22p2

Carbon could only make two bondsif no hybridization occurs. However,carbon can make four equivalent bonds.

sp3

hybrid orbitals

Ene

rgy

sp3

C atom of CH4 orbital diagram

B

A

BB

B


Hybridization of s and p Orbitals• The combination of an ns and an np orbital

gives rise to two equivalent sp hybrids oriented at 180º.

• Combination of an ns and two or three np orbitals produces three equivalent sp2 hybrids or four equivalent sp3 hybrids.


Hybridization of s and p Orbitals

• Both promotion and hybridization require an input of energy; the overall process of forming a compound with hybrid orbitals will be energetically favorable only if the amount of energy released by the formation of covalent bonds is greater than the amount of energy used to form the hybrid orbitals.


Hybridization Involving d Orbitals

3s 3p 3d 3s 3p 3d

promote

five sp3d orbitals 3dF

F

FP

F

F

A Be

Be

Be

Ba

Ba


hybridize

degenerateorbitals

(all EQUAL)

unhybridized P atomP = [Ne]3s23p3

vacant d orbitals

Pure atomicorbitals of

central atom

Hybridizationof the central

atom

Numberof hybridorbitals

Shape of hybridorbitals

s,p

s,p,p

s,p,p,p

s,p,p,p,d

s,p,p,p,d,d

sp

sp2

sp3

sp3d

sp3d2

2

3

4

5

6

Linear

Trigonal Planar

Tetrahedral

TrigonalBipyramidal

Octahedral

Hybridization Animation, by Raymond Chang

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/hybrv18.swf

Hybridization Animation, by Raymond Chang

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/hybrv18.swf

Bonding

• Single bonds– Overlap of bonding orbitals on bond axis– Termed “sigma” or σ bonds

• Double bonds– Sharing of electrons between 2 p orbitals

perpendicular to the bonding atoms– Termed “pi” or π bonds

2p 2p

Bond Axis of σ bond

One π bond

Multiple Bonds

2s 2p 2s 2p sp2 2p

promote hybridize

C C

H

H H

H

C2H4, ethene

one s bond and one p bond

H

H

CC

H

H s

ss

s

s

H

H

CC

H

H

Two lobes ofone p bond

Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 325-326

Multiple Bonds

2s 2p 2s 2p sp2 2p

promote hybridize

C C

H

H H

HC2H4, ethene

one s bond and one p bond

H

H

CC

H

H s

ss

s

s

H

H

CC

H

H

Two lobes ofone p bond

Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 325-326

C C

H

H

sp2

sp2

sp2

H

H

sp2

sp2

sp2

p p

p p

p bond

Internuclear axis

p p

s bonds

H

CHC

H

C

H

C

H C

H

C


C6H6 = benzene

2p atomic orbitals


s bonds and p bonds

H

C

HCH

C

H

C

H C

H

C


s bondsH

CH

C

HC

H

C

H

C

H

C

HC

H

C


s bondsH

C

H

C

HC

H

C

H

C

H

C

HC

H

C


N NO

O

O

O

dinitrogen tetraoxide

N2O4

NO

ON

O

O

hn 2 NO2nitrogen dioxide

(free radical)

colorless red-brown

1s 1s

s1s

s*1s

1s 1s

s1s

s*1s

H2 molecule

He2 molecule

H atom H atom

He atom He atom

Ene

rgy

Ene

rgy

Energy-level diagram for (a) the H2 molecule and (b) the hypothetical He2 molecule

(a)

(b)


Bond Order

Bond order = ½ (# or bonding electrons - # of antibonding electrons)

• A bond order of 1 represents a single bond,

• A bond order of 2 represents a double bond,

• A bond order of 3 represents a triple bond.

• A bond order of 0 means no bond exists.

Because MO theory also treats molecules with an odd number of electrons, Bond orders of 1/2 , 3/2 , or 5/2 are possible.

1s2 1s2

s1s

s*1s

2s1 2s1

s2s

s*2s

Energy-level diagram for the Li2 molecule

Li = 1s22s1

Li Li

Li2

Ene

rgy


2s 2s

s2s

s*2s

2p 2p

s2p

s*2p

p2p

p*2p

Energy-level diagram for molecular orbitals of second-row homonuclear diatomic molecules.


s2s

s*2s

s2p

p2pEnergy of molecular orbitals

O2, F2, Ne2 B2, C2, N2

Increasing 2s – 2p interaction


s2s

*s 2s

p2p

s2p

*p 2p

*s 2p

Large 2s – 2p interaction Small 2s – 2p interaction

s2s

*s 2s

s2p

p2p

*p 2p

*s 2p

C2 N2B2 F2 Ne2O2

Bond order

Bond enthalpy (kJ/mol)

Bond length (angstrom)

Magnetic behavior

1 2 3 2 1 0

290 620 941 495 155 -----

1.59 1.31 1.10 1.21 1.43 -----

Paramagnetic Diamagnetic Diamagnetic Paramagnetic Diamagnetic _____


s2s

s*2s

p2px

p*2px

p2py

p*2py

s2p

s*2p

C2

Arrange the atomic and molecular orbitals in order of increasing energy.How many orbitals are per molecule?

Can you distinguish the bonding from the antibonding MOs?


Magnetic Properties of a Sample

PARAMAGNETISM – molecules with one or more unpaired electrons are attracted into a magnetic field. (appears to weigh MORE in a magnetic field)

DIAMAGNETISM – substances with no unpaired electrons are weakly repelled from

a magnetic field. (appears to weigh LESS in a magnetic field)

Experiment for determining the magnetic properties of a sample

The sample is first weighed in the absence of a magnetic field.

When a field is applied, a diamagneticsample tends to move out of the fieldand appears to have a lower mass.

A paramagnetic sample is drawninto the field and thus appears to gain mass.

Paramagnetism is a much stronger effect than is diamagnetism.

sample N SN SN S


Experiment for determining the magnetic properties of a sample

N S N S

The sample is first weighed in the absence of a magnetic field.

When a field is applied, a diamagneticsample tends to move out of the fieldand appears to have a lower mass.

A paramagnetic sample is drawninto the field and thus appears to gain mass.

Paramagnetism is a much stronger effect than is diamagnetism.

sample


lone Pair single bond double bond triple bond

Electron Domains

Determine the shape of the BCl3 molecule:

B

Cl

ClCl

: :

::

:: :

:

:

There are 3 electron domains aboutthe central atom: no lone pairs andthree single bonds. Three electrondomains arrange themselves in atrigonal plane, with 120o angles.We predict a trigonal planar geometry.

:

B

Cl

ClCl

: :

::

:: :

:

Electron-domain geometry: trigonal planar

Molecular geometry (shape):

trigonal planar

One s orbital

Two p orbitals

Three sp2 hybrid orbitals

sp2 hybrid orbitalsshown together(large lobes only)

Hybridize

Ammonia, NH3Ammonia, NH3

Triangular pyramidalTriangular pyramidal


Introduction toBonding


http://www.nisd.net/communicationsarts/pages/chem

Chemical bond — the force that holds atoms together in a chemical compound

Covalent bonding — electrons are shared between atoms in a molecule or polyatomic ion

Ionic bonding — positively and negatively charged ions are held together by electrostatic forces

Ionic compounds — dissolve in water to form aqueous solutions that conduct electricity

Covalent compounds — dissolve to form solutions that do not conduct electricity

Vocabulary Chemical Bond

–attractive force between atoms or ions that binds them together as a unit

–bonds form in order to…• decrease potential energy (PE)• increase stability


Vocabulary

CHEMICAL FORMULA

molecularformula

formulaunit

IONIC COVALENT

CO2NaClCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Vocabulary

COMPOUND

ternarycompound

binarycompound

2 elementsmore than 2

elements

NaNO3NaClCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Vocabulary

ION

polyatomicIon

monatomicIon

1 atom 2 or more atoms

NO3-Na+


IONIC COVALENT

Bond Formation

Type of Structure

Solubility in Water

Electrical Conductivity

OtherProperties

e- are transferred from metal to nonmetal

high

yes (solution or liquid)

yes

e- are shared between two nonmetals

low

no

usually not

MeltingPoint

crystal lattice true molecules

Types of Bonds

Physical State

solid liquid or gas

odorousCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

“electron sea”

METALLICBond Formation

Type of Structure

Solubility in Water

Electrical Conductivity

OtherProperties

MeltingPoint

Types of Bonds

Physical State

e- are delocalized among metal atoms

very high

yes (any form)

no

malleable, ductile, lustrous

solid


Lattice Energies in Ionic Solids

Ionic compounds

1. Usually rigid, brittle, crystalline substances with flat surfaces that intersect at characteristic angles

2. Not easily deformed

3. Melt at relatively high temperatures

4. Properties result from the regular arrangement of the ions in the crystalline lattice and from the strong electrostatic attractive forces between ions with opposite charges

Metallic Bonding - “Electron Sea”

Types of Bonds

Bond Polarity Most bonds are

a blend of ionic and covalent characteristics.

Difference in electronegativity determines bond type.


Ionic

Polar-covalent

Nonpolar-covalent

3.3

1.7

0.3

0

100%

50%

5%

0%

Diff

eren

ce in

ele

ctro

nega

tiviti

es Percentage ionic character

Types of Chemical Bonds

Copyright © 2006 Pearson Education Inc., publishing as Benjamin Cummings

Bond Polarity

Electronegativity–Attraction an atom has for a shared

pair of electrons.–higher e-neg atom -

– lower e-neg atom +


Ionic bonding: Li + ClIonic bonding (stealing/transfer of electrons) can be represented in three different ways

Li + Cl [Li]+[Cl]–

1e-

3p+

4n02e- 17p+

18n08e-8e-2e 3p+

4n02e-1e- 17p+

18n07e- 8e- 2e-

Li Cl [ Cl ]–[Li]+

lithium atom chlorine atom lithium ion chlorine ionchloride ion

Ionic bonding: Mg + O

Mg + O [Mg]2+[O]2–

12p+

12n02e- 8e- 2e-

1e-

[ O ]2–[Mg]2+

6e- 2e-

8n0 8p+

1e-

8e- 2e-

8n0 8p+ 12p+

12n02e- 8e-

OMg

F

4.0

Ar

--

Kr

3.0

Xe

2.6

Rn

2.4

Bond Polarity Electronegativity Trend

– Increases up and to the right.

Be

1.5

Al

1.5

Si

1.8

Ti

1.5

V

1.6

Cr

1.6

Mn

1.5

Fe

1.8

Co

1.8

Ni

1.8

Cu

1.9

Zn

1.7

Ga

1.6

Ge

1.8

Nb

1.6

Mo

1.8

Tc

1.9

Ag

1.9

Cd

1.7

In

1.7

Sn

1.8

Sb

1.9

Ta

1.5

W

1.7

Re

1.9

Hg

1.9

Tl

1.8

Pb

1.8

Bi

1.9

N

3.0

O

3.5

F

4.0

Cl

3.0

C

2.5

S

2.5

Br

2.8

I

2.5

Na

0.9

K

0.8

Rb

0.8

Cs

0.7

Ba

0.9

Fr

0.7

Ra

0.9

H

2.1

B

2.0

P

2.1

As

2.0

Se

2.4

Ru

2.2

Rh

2.2

Pd

2.2

Te

2.1

Os

2.2

Ir

2.2

Pt

2.2

Au

2.4

Po

2.0

At

2.2

Actinides: 1.3 - 1.5

Li

1.0

Ca

1.0

Sc

1.3

Sr

1.0

Y

1.2

Zr

1.4

Hf

1.3

Mg

1.2

La

1.1

Ac

1.1

Lanthanides: 1.1 - 1.3

*

*y

y

Ne

--

He

--

Bond Polarity Electronegativity Trend

– Increases up and to the right.

1

2

3

4

5

6

1

2

3

4

5

6

7

1A

2A

3B 4B 5B 6B 7B 1B 2B

3A 4A 5A 6A 7A

8A

8B

Nonpolar Covalent Bond–electrons are shared equally–symmetrical electron density–usually identical atoms

Bond Polarity

+ -

Bond Polarity Polar Covalent Bond

–electrons are shared unequally–asymmetrical e- density– results in partial charges (dipole)


Nonpolar

Polar

Ionic

Bond Polarity


Bond PolarityExamples:

Cl2

HCl

NaCl

3.0 - 3.0 = 0.0 Nonpolar

3.0 - 2.1 = 0.9 Polar

3.0 - 0.9 = 2.1 Ionic

Ionic

Polar-covalent

Nonpolar-covalent

3.3

1.7

0.3

0

100%

50%

5%

0%

Diff

eren

ce in

ele

ctro

nega

tiviti

es Percentage ionic character

Mg

Write the electron dot diagram for

Na Mg C O F Ne He

1s22s22p63s1

1s22s22p63s2

1s22s22p2

1s22s22p4

1s22s22p5

1s22s22p6

1s2

Na

C

O

F

He

Ne

Ionic Bonding

Na Cl

transfer of electron

+ -

NaCl

Ca +2

P -3Ca

+2

P

All the electrons must be accounted for!

+2

Ionic Bonding

Ca -3

Ionic Bonding

Ca3P2

Formula Unit

Ca2+

Ca2+

Ca2+

P3-

P 3-

Ca2+

P3- Ca2+

P3- Ca2+

Metals are Malleable Hammered into shape (bend). Ductile - drawn into wires. Electrons allow atoms to slide by.

+ + + ++ + + +

+ + + +

Ionic solids are brittle

+ - + -+- +-

+ - + -+- +-

Force + - + -

Strong repulsion breaks crystal apart.

How does H2 form?

+ +

The nuclei repel

But they are attracted to electrons

They share the electrons

Hydrogen Bond Formation

0.74 A

- 436

0

H – H distance

Ene

rgy

(KJ/

mol

)


no interaction

increasedattraction

balanced attraction& repulsion

increasedrepulsion

Potential Energy Diagram - Attraction vs. Repulsion

(internuclear distance)

Covalent bonds Nonmetals hold onto their valence

electrons. They can’t give away electrons to bond. Still want noble gas configuration. Get it by sharing valence electrons with

each other. By sharing both atoms get to count the

electrons toward noble gas configuration.

Covalent bonding

F F

Fluorine has seven valence electrons A second F atom also has seven By sharing electrons Both end with full orbitals (stable octets)

8 Valence electrons

8 Valence electrons

Single Covalent Bond A sharing of two valence electrons. Only nonmetals and Hydrogen. Different from an ionic bond because

they actually form molecules. Two specific atoms are joined. In an ionic solid you can’t tell which

atom the electrons moved from or to.

How to show how they formed It’s like a jigsaw puzzle. I have to tell you what the final formula

is. You put the pieces together to end up

with the right formula. For example - show how water is

formed with covalent bonds.

Water

H

O

Each hydrogen has 1 valence electron

Each hydrogen wants 1 more

The oxygen has 6 valence electrons

The oxygen wants 2 more

They share to make each other happy

Water Put the pieces together The first hydrogen is happy The oxygen still wants one more

H O

Water The second hydrogen attaches Every atom has full energy levels A pair of electrons is a single bond

H OH H

H O

Lewis Structures

1) Count up total number of valence electrons

2) Connect all atoms with single bonds

- “multiple” atoms usually on outside - “single” atoms usually in center;

C always in center,H always on outside.

3) Complete octets on exterior atoms (not H, though)4) Check - valence electrons math with Step 1 - all atoms (except H) have an octet; if not, try multiple bonds - any extra electrons? Put on central atom

Multiple Bonds Sometimes atoms share more than one

pair of valence electrons. A double bond is when atoms share two

pair (4) of electrons. A triple bond is when atoms share three

pair (6) of electrons.

Carbon dioxide CO2 - Carbon is central

atom ( I have to tell you)

Carbon has 4 valence electrons

Wants 4 more Oxygen has 6 valence

electrons Wants 2 more

O

C

Carbon dioxide Attaching 1 oxygen leaves the oxygen 1

short and the carbon 3 short

OC

Carbon dioxide Attaching the second oxygen leaves

both oxygen 1 short and the carbon 2 short

OCO

The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in

the bond8 valence electrons

8 valence electrons

8 valence electrons

Carbon dioxide

OCO

Formation of Multiple Covalent Bonds

By combining more than one unpaired electron at a time, a double bond is formed.

Both oxygen atoms end up with eight valence electrons.

Ox x

x

xOx

x

x

x

x

xO

x

x O

How to draw them Add up all the valence electrons. Count up the total number of electrons

to make all atoms happy. Subtract. Divide by 2 Tells you how many bonds - draw them. Fill in the rest of the valence electrons

to fill atoms up.

Examples NH3

N - has 5 valence electrons wants 8

H - has 1 valence electrons wants 2

NH3 has 5+3(1) = 8

NH3 wants 8+3(2) = 14

(14-8)/2= 3 bonds 4 atoms with 3 bonds

N

H

N HHH

Examples Draw in the bonds All 8 electrons are accounted for Everything is full

Examples HCN C is central atom N - has 5 valence electrons wants 8 C - has 4 valence electrons wants 8 H - has 1 valence electrons wants 2

HCN has 5 + 4 + 1 = 10

HCN wants 8 + 8 + 2 = 18

(18 - 10) / 2= 4 bonds 3 atoms with 4 bonds -will require

multiple bonds - not to H

HCN Put in single bonds Need 2 more bonds Must go between C and N

NH C

HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add

NH C

HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add Must go on N to fill octet

NH C

Another way of indicating bonds

Often use a line to indicate a bond Called a structural formula Each line is 2 valence electrons

H HO =H HO

Structural Examples

H C N

C OH

H

C has 8 electrons because each line is 2 electrons

Ditto for N

Ditto for C here Ditto for O

Coordinate Covalent Bond When one atom donates both electrons

in a covalent bond. Carbon monoxide CO

OC

How do we know if Have to draw the diagram and see what

happens. Often happens with polyatomic ions and

acids.

Resonance When more than one dot diagram with

the same connections are possible.

NO2-

Which one is it? Does it go back and forth. It is a mixture of both, like a mule.

NO3-

VSEPR Valence Shell Electron Pair Repulsion. Predicts three dimensional geometry of

molecules. Name tells you the theory. Valence shell - outside electrons. Electron Pair repulsion - electron pairs

try to get as far away as possible. Can determine the angles of bonds.

VSEPR Based on the number of pairs of

valence electrons both bonded and unbonded.

Unbonded pair are called lone pair.

CH4 - draw the structural formula

Has 4 + 4(1) = 8 wants 8 + 4(2) = 16 (16-8)/2 = 4 bonds

VSEPR

Single bonds fill all atoms.

There are 4 pairs of electrons pushing away.

The furthest they can get away is 109.5º.

C HH

H

H

4 atoms bonded Basic shape is

tetrahedral. A pyramid with a

triangular base. Same shape for

everything with 4 pairs. CH H

H

H

109.5º

3 bonded - 1 lone pair

N HH

H

NH HH

<109.5º

Still basic tetrahedral but you can’t see the electron pair.

Shape is calledtrigonal pyramidal.

2 bonded - 2 lone pair

OH

H

O HH

<109.5º

Still basic tetrahedral but you can’t see the 2 lone pair.

Shape is calledbent.

3 atoms no lone pair

CH

HO

The farthest you can the electron pair apart is 120º

3 atoms no lone pair

CH

HO

The farthest you can the electron pair apart is 120º.

Shape is flat and called trigonal planar.

C

H

H O

120º

2 atoms no lone pair With three atoms the farthest they can

get apart is 180º. Shape called linear.

C OO180º

Hybrid Orbitals

Combines bonding with geometry

Hybridization The mixing of several atomic orbitals to form the

same number of hybrid orbitals. All the hybrid orbitals that form are the same

(degenerate = equal energy). sp3 - one s and three p orbitals mix to form four

sp3 orbitals. sp2 - one s and two p orbitals mix to form three sp2

orbitals leaving one p orbital. sp - one s and one p orbitals mix to form four sp

orbitals leaving two p orbitals.

Hybridization We blend the s and p-orbitals of the

valence electrons and end up with the tetrahedral geometry. We combine one s orbital and three p-orbitals.

sp3 hybridization has tetrahedral geometry.

sp3 geometry

109.5º

This leads to tetrahedral shape.

Every molecule with a total of 4 atoms

and lone pair is sp3 hybridized.

Gives us trigonal pyramidal and bent shapes also.

How we get to hybridization

We know the geometry from experiment. We know the orbitals of the atom hybridizing

atomic orbitals can explain the geometry. So if the geometry requires a tetrahedral

shape, it is sp3 hybridized. –This includes bent and trigonal

pyramidal molecules because one of the sp3 lobes holds the lone pair.

sp2 hybridization

C2H4

double bond acts as one pair trigonal planar Have to end up with three blended orbitals

–use one s and two p orbitals to make three sp2 orbitals.

– leaves one p orbital perpendicular

Where is the P orbital? Perpendicular The overlap of

orbitals makes a sigma bond

(s bond)

Two types of Bonds Sigma bonds from overlap of orbitals

between the atoms Pi bond (p bond) above and below atoms Between adjacent p orbitals. The two bonds of

a double bond

CCH

H

H

H

H

H

HB

sp2 hybridization

when three things come off atom trigonal planar 120º one p bond

B

B

A

B

trigonal planar

hybridize

s orbital

p orbitals three sps hybrid orbitals

What about two when two things come off one s orbital and one p orbital hybridize linear

sp hybridization end up with two lobes 180º

apart. p orbitals are at right

angles makes room for two p

bonds and two sigma bonds.

a triple bond or two double bonds

CO2

C can make two s and two p O can make one s and one p

CO O

Polar Bonds When the atoms in a bond are the

same, the electrons are shared equally. This is a nonpolar covalent bond. When two different atoms are

connected, the atoms may not be shared equally.

This is a polar covalent bond. How do we measure how strong the

atoms pull on electrons?

Electronegativity

A measure of how strongly the atoms attract electrons in a bond.

The bigger the electronegativity difference the more polar the bond.

0.0 - 0.5 Covalent nonpolar

0.5 - 1.0 Covalent moderately polar

1.0 -2.0 Covalent polar

>2.0 Ionic

How to show a bond is polar Isn’t a whole charge just a partial charge + d means a partially positive - d means a partially negative

The Cl pulls harder on the electrons The electrons spend more time near the Cl

H Cl+d -d

Polar Molecules

Molecules with ‘ends’

Polar Molecules Molecules with a positive and a

negative end Requires two things to be true

The molecule must contain polar bonds This can be determined from

differences in electronegativity.

Symmetry can not cancel out the effects of the polar bonds. Must determine geometry first.

Is it polar?

......

H Cl

Polar

H HO

PolarF F

B

F

Nonpolar

HH

H

N

Polar

Polar Nonpolar

F FCl

F

F F

Xe

F F

Nonpolar Polar

Cl

ClC

Cl

Cl

HC

Cl

H

H

XeF4

CCl4 CH3Cl

HCl H2O BF3 NH3

Bond Dissociation Energy The energy required to break a bond C - H + 393 kJ C + H We get the Bond dissociation energy

back when the atoms are put back together

If we add up the BDE of the reactants and subtract the BDE of the products we can determine the energy of the reaction (DH)

Find the energy change for the reaction

CH4 + 2O2 CO2 + 2H2O

For the reactants we need to break 4 C-H bonds at 393 kJ/mol and 2 O=O bonds at 495 kJ/mol= 2562 kJ/mol

For the products we form 2 C=O at 736 kJ/mol and 4 O-H bonds at 464 kJ/mol

= 3328 kJ/mol reactants - products = 2562-3328 = -766kJ

Intermolecular Forces

What holds molecules

to each other?

Intermolecular Forces They are what make solid and liquid

molecular compounds possible. The weakest are called van derWaal’s

forces - there are two kinds Dispersion forces Dipole Interactions

–depend on the number of electrons –more electrons stronger forces–bigger molecules

Depend on the number of electrons More electrons stronger forces Bigger molecules more electrons

fluorine (F2) is a gas

bromine (Br2) is a liquid

iodine (I2) is a solid

Dipole interactions

Dipole interactions

Occur when polar molecules are attracted to each other.

Slightly stronger than dispersion forces. Opposites attract but not completely

hooked like in ionic solids.

Dipole interactions Occur when polar molecules are

attracted to each other. Slightly stronger than dispersion forces. Opposites attract but not completely

hooked like in ionic solids.

H Fd+ d-

H Fd+ d-

Dipole Interactionsd

+ d

-

d+ d-

d+ d -

d+ d-

d+ d -

d+

d-

d + d

-d+

d-

Hydrogen bonding Are the attractive force caused by

hydrogen bonded to F, O, or N. F, O, and N are very electronegative so

it is a very strong dipole. The hydrogen partially share with the

lone pair in the molecule next to it. The strongest of the intermolecular

forces.

Hydrogen Bonding

HH

Od+

d-

d+

H HOd

+

d-

d+

Hydrogen bonding

HH

O H HO

HH

O

H

H

OH

HO

H

HO HH

O

Resources - Bonding

Objectives

Episode 8 – Chemical Bonds

Episode 9 – Molecular Architecture

http://www.chem1.com/acad/webtext/chembond/index.html

http://www.backflip.com/perl/go.pl?url=18424627

http://www.backflip.com/perl/go.pl?url=18424628

molecular geometry and bonding theories ap chemistry – ch 9 mr. christopherson

Documents

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o h h h h n

c h h hh h h h h

molecular geometry h

o b e ab

tetrahedral ab

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