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

• 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.

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Lewis Structures - ExamplesLewis Structures - Examples

Exceptions to the Octet RuleExceptions to the Octet Rule

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 – Central Atoms - Less than an OctetExceptions – Central Atoms - Less than an Octet

Exceptions to the Octet RuleExceptions to the Octet Rule

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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Exceptions – Central Atoms - Greater than an OctetExceptions – Central Atoms - Greater than an Octet

• There are five fundamental geometries for molecular shape:

Molecular ShapesMolecular Shapes

Figure 9.3

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Molecular Shapes – 3D NotationsMolecular Shapes – 3D Notations

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 CyberChem video

Examples: VSEPR Molecular ShapesExamples: VSEPR Molecular Shapes

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 increase

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

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

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 NCyberChem video

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