Honors Chemistry

Chapter 9: Chemical Bonding I

9.1 Lewis Dot Diagrams

• Symbol surrounded by dots for valence e-• Separate dots as much as possible

9.2 Ionic Bonding

• Electrostatic force holding ions together• Ions formed by electron transfer

• Low Eion loses e-, high e.a. gains e-

•Na + Cl Na+ Cl-

•Mg + S Mg2+ S2-

•Li + S 2 Li+ S2-

9.3 Lattice Energy

• Energy that holds ionic compounds together in a crystal lattice

• Transfer of e- requires energy (Eion) and releases energy (e.a.)

• In general, the cation requires more energy than the anion releases, which makes bond formation unstable

• Lattice energy releases additional energy, making bond formation stable

9.4 The Covalent Bond

• Covalent bond = shared pair of electrons•

F + F F – F

• Shared pair – shared electrons, bond• Lone pair – electrons not involved in bond

9.4 Lewis Structures

• Representation of covalent compounds using dots for e- and lines for bonds

• Octet rule• atoms bond in such a way as to gain 8 e- in

valence shell• Exceptions – H and He

• Multiple Bonds• Double bond – share 2 pairs; eg, O2

• Triple bond – share 3 pairs; eg, N2

9.4 Bonding Summary

9.4 Ionic / Covalent Properties

• Intermolecular attractive forces• Ionic – strong, covalent – weak

• Consider phase, density, solubility, conductivity

• Ionic Covalent• Solid Liquid or gas• High density Low density• Usually soluble Often insoluble• Good conductor Poor conductor

9.5 Electronegativity

• Element’s relative attraction for shared e-

9.5 Pauling Electronegativities

9.5 Electronegativity and Atomic Radius

9.5 Bond Character

• Degree of sharing of the bonded e-• Depends on the difference in electronegativity• Small electronegativity difference

• Equal sharing of bonded e-• True covalent bond (nonpolar covalent)

• Moderate electronegativity difference• Unequal sharing of bonded e-• Polar covalent bond

• Large electronegativity difference• Ionic bond

9.5 Bond Character

EN

3.0

2.0

0.0

9.6 Writing Lewis Structures

• Draw a reasonable skeletal structure for the compound

• Count the total valence electrons available• Draw single bonds between all atoms and

use remaining electrons to fulfill the octet rule

• If there are not enough electrons to fulfill the octet rule, form double or triple bonds

• Draw Lewis structures for NH3, O3, CO32-

9.7 Formal Charges

• Difference between the number of e- an atom has in a Lewis structure and the number of e- it has as a free element

• Assigning formal charges• Atom counts all its nonbonding e-’s• Atom counts 1 e- from each bond

• Total formal charge must add up to the total charge of the molecule / ion

9.7 Formal Charges

• The most plausible structures have:• The fewest formal charges• Formal charges of smallest magnitude• Negative formal charges on the most

electronegative elements• No adjacent charges of the same type

9.8 Resonance

• Some molecules have more than one plausible Lewis structure

OO O

OO O

• Resonance –use of both structures to represent a molecule

• Reality is that bonds are delocalizedO

O O

9.8 Resonance

• Draw resonance structures for• N2O• HSCN• NO3

-

• CO32-

• Extreme resonance – C6H6

9.9 Exceptions to the Octet Rule

• Incomplete Octets• Not enough electrons to make an octet• Usually occurs with Groups IIA and IIIA• Examples: BeH2, BF3

• Consider resonance in BF3

• BF3 + NH3 F3BNH3

• Coordinate covalent bond• One atom donates both shared electrons

9.9 Exceptions to the Octet Rule

• Odd Electron Molecules• Often called radicals• Examples: NO, NO2

• Single e- goes on element with lower EN• Very reactive• Tend to form dimers

9.9 Exceptions to the Octet Rule

• Expanded Octets• More than 8 e- on central atom• Requires the atom to have a d orbital• Can’t happen with periods 1 and 2• Examples: SF6, XeF4, ClO4-