Bonding

Ionic Bonding

Resonance Structures

VSEPRBasic Shapes3-D NotationHybridization (Lab)

M olecular G eom etries

Octet Rule Polar M olecules

Lew is Structures Covalent Bonding

T ypes of Bonds

Chemical BondingChemical Bonding

Chlorides of Period 2compound LiCl BeCl2 BCl3 CCl4 NCl3 OCl2 Cl2

melting point

610 415 -107 -23 -40 -121 -102

Chlorides of Period 3compound NaCl MgCl2 AlCl3 SiCl4 PCl3 SCl6 Cl2

melting point

801 714 193 -69 -112 -51 -102

Let’s examine the melting point of compounds across two periods. What is the trend?

lowhigh

Conductivity - high Conductivity - low

Bonding Can we explain the melting point behavior across a period?

• involves the valence electrons or outermost shell (or highest shell) electrons

• for group A elements - the group number tells how many valence electrons

How many valence electrons on N?

Group 5A – 5 valence electrons

Bonding…

• Lewis dot structures show the valence electrons around at atom and for most molecules and compounds a complete octet for the elements

• most monatomic ions have an electron configuration of noble gases

N Al

F + e- F1s22s22p5 1s22s22p6 Ne

Write out the electron configuration for the following atoms and ions:

He

Ne

Ar

Kr

H-

Ca+2

O-2

Br-

Which noble gas is isoelectronic with each ion?

Na Cl

NaCl

+ -

electron transfer

and the formation of ions

This is the formation of an ionic bondionic bond.

ClCl

Cl2 This is the formation of a covalent bondcovalent bond.

sharing of a pair of electrons

and the formation of molecules

What about the distance between the atoms in a bond?

NaCl Na+ Cl- d = 281 pm

Cl2 Cl-Cl d = 199 pm

What property can be used to tell when a bond will ionic or covalent?

ionic

covalentHCl H2O CH4

MgO CaCl2 Na2S

Draw the Lewis dot structures for the following compounds:

Some exceptions to the Octet Rule

BF3

PCl5

SF6

Electronegativity

Chlorides of Period 2compound LiCl BeCl2 BCl3 CCl4 NCl3 OCl2 Cl2

EN 2.2 1.6 1.1 0.6 0 0.6 0

Chlorides of Period 3Compound NaCl MgCl2 AlCl3 SiCl4 PCl3 SCl6 Cl2

EN 2.2 1.9 1.6 1.3 1.0 0.6 0

large difference small difference

The electronegativity difference - EN = ENhigher – EN lower

Using electronegativities to determine bond type

EN > 1.7 ionic bond - transfer

EN < 1.7 covalent bond - sharing

So we have a range of electronegativity difference of 0 to 1.7 for sharing an electron pair.

Is the sharing of electrons in molecules always equal?

X Y EN = 0

X Y EN = 0.3

X Y EN = 0.6

X Y EN = 0.9

X Y EN = 1.2

ENY > ENX

Which element is more electronegative?

non-polar bond

incr

easi

ng p

olar

ity

of b

ond

polar bond

0 < EN < 1.7

Direction of electron migration

BF3 – a planar molecule

Ball & stick

BF

Space-filled

Electrostatic potential maps

top side

negative

positive

Spartan ‘02

2.0

4.0

More sharing examples

O2

N2

O O

N N

O O O O

N N N N N N

double bond (2 pairs)

triple bond (3 pairs)

Share until octet is complete.

octet complete

Bond Energy

F2 single bond BE = 142 kJ/mole

O2 double bond BE = 494

N2 triple bond BE = 942

X2 + energy X + X

incr

easi

ng b

ond

stre

ngth

Is breaking a bond an endothermic or exothermic process?

http://wulfenite.fandm.edu/Data%20/Table_6.html

NH3

NH

H

H

NH

H

HH+

NH4+ NH3 + H+ NH4

+

coordinate covalent bond(the pair of electrons from

the same atom)

normal covalent bond(each atom supplies an

electron)

Some more sharing examples

Type of bond? – I, PC, or NC

TiO2 CH4 NaI

CS2 CO2 KCl

AlCl3 CsF HBr

Using the EN trends to predict bond type

105

Db107

Bh

Increasing EN

Incr

easi

ng E

N

NO RbF FeS H2SModified from http://www.cem.msu.edu/~djm/cem384/ptable.html

Draw the Lewis dot structures

CO2

H3O+ CO

H2CO

NH2-

HCN

(C in center)

Show the direction of electron migration ( ) in the following.

C – H

H – F

C = O

C – Cl

Rank the bond polarity (1-most … 3-least)

As-H N-H P-H

Here is the electrostatic potential map for H2CO.

Show the electron migration on this planar molecule.

C OH

H

How is this molecule different than BF3?

blue – positive red - negative

ionic covalent

valence electrons

Comparison of Bonding Types

sharing of electrons

transfer of electrons

ions molecules

EN > 1.7 EN < 1.7

high mp low mp

molten salts conductive

non-conductive

Bonding spectrum100% covalent 100% ionic

A+ B-A B A B

Increasing EN

Increasing polarity Transfer

• Chemical bond: attractive force holding two or more atoms together.

• Covalent bond results from sharing electrons between the atoms. Usually found between nonmetals.

• Ionic bond results from the transfer of electrons from a metal to a nonmetal.

• Metallic bond: attractive force holding pure metals together.

Chemical Bonds, Lewis Symbols, and the Octet RuleChemical Bonds, Lewis Symbols, and the Octet Rule

Figure 8.3: Ionic BondingFigure 8.3: Ionic Bonding

Figure 8.5: Covalent BondingFigure 8.5: Covalent Bonding

Chemical BondsChemical Bonds

Bond Type Single Double Triple

# of e’s 2 4 6

Notation — =

Bond order 1 2 3

Bond strength

Increases from Single to Triple

Bond length Decreases from Single to Triple

Strengths of Covalent BondsStrengths of Covalent Bonds

Lewis Symbols

Chemical Bonds, Lewis Symbols, and the Octet RuleChemical Bonds, Lewis Symbols, and the Octet Rule

The Octet Rule• All noble gases except He has an s2p6 configuration. • Octet rule: atoms tend to gain, lose, or share electrons

until they are surrounded by 8 valence electrons (4 electron pairs).

• Caution: there are many exceptions to the octet rule.

Chemical Bonds, Lewis Symbols, and the Octet RuleChemical Bonds, Lewis Symbols, and the Octet Rule

Bond Polarity and ElectronegativityBond Polarity and Electronegativity

Electronegativity• Electronegativity: The ability of one atoms in a

molecule to attract electrons to itself.• Pauling set electronegativities on a scale from 0.7 (Cs) to

4.0 (F).• Electronegativity increases

• across a period and

• down a group.

Figure 8.6: Electronegativities of ElementsFigure 8.6: Electronegativities of Elements

Electronegativity

Bond Polarity and ElectronegativityBond Polarity and Electronegativity

Figure 8.7: Electronegativity and Bond Polarity• There is no sharp distinction between bonding types.• The positive end (or pole) in a polar bond is represented

+ and the negative pole -.

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Drawing Lewis StructuresDrawing Lewis Structures

Follow Step by Step Method (See Ng Web-site)

1. Total all valence electrons. [Consider Charge]

2. Write symbols for the atoms and guess skeleton structure [ define a central atom ].

3. Place a pair of electrons in each bond.

4. Complete octets of surrounding atoms. [ H = 2 only ]

5. Place leftover electrons in pairs on the central atom.

6. If there are not enough electrons to give the central atom an octet, look for multiple bonds by transferring electrons until each atom has eight electrons around it.

HyperChemCyberChem (Lewis) video

Exceptions to the Octet RuleExceptions to the Octet Rule

Central Atoms Having Less than an Octet• Relatively rare.• Molecules with less than an octet are typical for

compounds of Groups 1A, 2A, and 3A.

• Most typical example is BF3.

• Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.

Exceptions to the Octet RuleExceptions to the Octet Rule

Central Atoms Having More than an Octet• This is the largest class of exceptions.• Atoms from the 3rd period onwards can accommodate

more than an octet.• Beyond the third period, the d-orbitals are low enough in

energy to participate in bonding and accept the extra electron density.

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• There are five fundamental geometries for molecular shape:

Molecular Shapes: VSEPRMolecular Shapes: VSEPR

Molecular Shapes – 3D NotationsMolecular Shapes – 3D Notations

VSEPR (Ballons)-Movie Clip

Figure 9.3

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Summary of VSEPR Molecular ShapesSummary of VSEPR Molecular Shapese-pairs Notation Name of VSEPR shape Examples

2 AX2 Linear HgCl2 , ZnI2 , CS2 , CO2

3 AX3 Trigonal planar BF3 , GaI3

AX2E Non-linear (Bent) SO2 , SnCl2

4 AX4 Tetrahedral CCl4 , CH4 , BF4-

AX3E (Trigonal) Pyramidal NH3 , OH3-

AX2E2 Non-Linear (Bent) H2O , SeCl2

5 AX5 Trigonal bipyramidal PCl5 , PF5

AX4E Distorted tetrahedral

(see-sawed)

TeCl4 , SF4

AX3E2 T-Shaped ClF3 , BrF3

AX2E3 Linear I3- , ICl2

-

6 AX6 Octahedral SF6 , PF6-

AX5E Square Pyramidal IF5 , BrF5

AX4E2 Square Planar ICl4- , BrF4

-See Ng Web-siteHyperChem CyberChm Gems

Examples: VSEPR Molecular Shapes - IExamples: VSEPR Molecular Shapes - I# electron pairs on Central Atom A

Notation Example Lewis VSEPR & Name of Shape

2AX2

2 bp on A

3

AX3 3 bp on A

AX2E2 bp and 1 lp on A

Examples: VSEPR Molecular Shapes – I – F08Examples: VSEPR Molecular Shapes – I – F08

Examples: VSEPR Molecular Shapes - IIExamples: VSEPR Molecular Shapes - II# electron pairs on Central Atom A

Notation Example Lewis VSEPR & Name of Shape

4

AX4

4 bp on A

AX3E 3 bp and 1 lp on A

AX2E2

2 bp and 2 lp on A

Examples: VSEPR Molecular Shapes – II – F08Examples: VSEPR Molecular Shapes – II – F08

Examples: VSEPR Molecular Shapes - IIIExamples: VSEPR Molecular Shapes - III# electron pairs on Central Atom A

Notation Example LewisVSEPR & Name of Shape

5

AX5 5 bp on A

AX4E4 bp and 1 lp on A

AX3E2

3 bp and 2 lp on A

AX2E3

2 bp and 3 lp on A

Examples: VSEPR Molecular Shapes - IVExamples: VSEPR Molecular Shapes - IV# electron pairs on Central Atom A

Notation Example Lewis VSEPR & Name of Shape

6

AX6

6 bp on A

AX5E 5 bp and 1 lp on A

AX4E2

4 bp and 2 lp on A

The Effect of Nonbonding Electrons

• By experiment, the H-X-H bond angle decreases on moving from C to N to O:

• Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs.

• Therefore, the bond angle decreases as the number of lone pairs increases

VSEPR ModelVSEPR Model

104.5O107O

NHH

HC

H

HHH109.5O

OHH

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Figure 9.10: Shapes of Larger Molecules

• In acetic acid, CH3COOH, there are three central atoms.

VSEPR ModelVSEPR ModelHyperChem

Shapes of Larger Molecules• In glycine (simplest amino acid), NH2CH2CO2H, there are four possible central atoms.

• Draw the Lewis Structure and the 3D VSEPR Molecular Geometry for glycine. Indicate the name of the shape for all possible central atoms, including estimation of bond angles.

• Hint 1: Designate the 2nd carbon in the formula as the central atom in skeleton structure.

• Hint 2: The acid portion of glycine is the same as that of acetic acid.

Lewis-VSEPR HW assigned 10/29/10 . Due 11/1/10.Lewis-VSEPR HW assigned 10/29/10 . Due 11/1/10.Lewis-VSEPR HW assigned 10/29/10 . Due 11/1/10.Lewis-VSEPR HW assigned 10/29/10 . Due 11/1/10.

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Solution Key

Figure 8.10: Drawing Lewis StructuresFigure 8.10: Drawing Lewis Structures

Resonance Structures

Figure 9.12

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Figure 9.11: Molecular Shape and Molecular PolarityFigure 9.11: Molecular Shape and Molecular Polarity

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Figure 9.13: Molecular Shape and Molecular PolarityFigure 9.13: Molecular Shape and Molecular Polarity

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• Lewis structures and VSEPR do not explain why a bond forms.

• How do we account for shape in terms of quantum mechanics?

• What are the orbitals that are involved in bonding?• We use Valence Bond Theory:

• Bonds form when orbitals on atoms overlap.

• There are two electrons of opposite spin in the orbital overlap.

Covalent Bonding and Orbital OverlapCovalent Bonding and Orbital Overlap

Gems - Movie Clip

Figure 9.14: Covalent Bonding and Orbital OverlapFigure 9.14: Covalent Bonding and Orbital Overlap

• To determine the electron pair geometry:• draw the Lewis structure,

• count the total number of electron pairs around the central atom,

• arrange the electron pairs in one of the above geometries to minimize e-e repulsion, and count multiple bonds as one bonding pair.

VSEPR Model (Figure 9.6)VSEPR Model (Figure 9.6)

VSEPR VSEPR ModelModel

Drawing Lewis StructuresDrawing Lewis Structures

Formal Charge• Consider:

• For C: • There are 4 valence electrons (from periodic table).

• In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure.

• Formal charge: 4 - 5 = -1.

C N

Drawing Lewis StructuresDrawing Lewis Structures

Formal Charge• Consider:

• For N:• There are 5 valence electrons.• In the Lewis structure there are 2 nonbonding electrons and 3

from the triple bond. There are 5 electrons from the Lewis structure.

• Formal charge = 5 - 5 = 0.• We write:

C N

C NCyberChm Gems

Ionic Bonding

Resonance Structures

VSEPRBasic Shapes3-D NotationHybridization (Lab)

M olecular G eom etries

Octet Rule Polar M olecules

Lew is Structures Covalent Bonding

T ypes of Bonds

Chemical BondingChemical Bonding

Lewis

AXE notation

VSEPR shapes

Polarity