Ch3: Bonding and Chemical Interactions

Chemical Bonds

Attractive forces holding molecules together

Form from valence e interactions

Octet Rule

Atoms bond to fill outer shell with 8 e (e configuration of nearest noble gas)

Exception: Incomplete Octet

Stable with

Exception: Expanded Octet

Period 3> hold >8 e

P: 10

S: 12

Cl: 14

Exception: Odd Numbers of e

Cannot distribute for 8 e per atom

Common Elements Always Following Octet Rule

C, N, O, F, Na, Mg

Ionic Bonding

e from atom with low IE (metal) transfer to atom with high EA (nonmetal)

Electrostatic attraction between opposite charges hold ions together

Create lattice structures

Covalent Bonding

e pair shared between 2 atoms (nonmetals) with similar EN

Sharing equality (EN) determines bond polarity

Create molecules

Covalent: Nonpolar Bond

Equal e sharing

∆EN < 0.5

Covalent: Polar Bond

Unequal e sharing

0\.5 < ∆EN < 1.7

Partial charges

Covalent: Coordinate Bond

Both shared e from 1 atom

Between lewis acid and base (unhybridized p orbital form bond)

Ionic Bonds

∆EN > 1.7

Usually between metal and nonmetal

Ionic Physical Properties

High mp and bp

Dissolve in polar solvents

Good conductors

Crystalline lattice solid maximizes attractive forces and minimizes repulsive forces

Covalent Bonds

Energy for e transfer > energy release from bond

Favour sharing e

Covalent Physical Properties

Weak IMF

Low mp and bp

Poor conductors

Covalent Bond Order

\# shared e pairs between atoms

Covalent Bond Length

Average distance between nuclei of bonded atoms

Increase bond order = decrease bond length

Covalent Bond Energy

Energy required to break bond in gaseous state

Increase bond order = increase bond energy

Covalent Bond Polarity

Dipole moment from EN diff

Higher EN atom pulls e density

Bonding e

Involved in covalent bond

In valence shell

Nonbonding e (Lone Pairs)

Not involved in covalent bond

In valence shell

Lewis Structure 1: Backbone

Least EN atom in centre

H (always) and halogens (usually) in terminal position

Lewis Structure 2: Valence e

Sum of valence e = total valence e in molecule

Lewis Structure 3: Bonds and Lone Pairs

Draw single bonds to central atom

Complete octets of outer atoms then central atom

Lewis Structure 4: Multiple Bonds

Central atom less than octet, rewrite double/triple bonds with surrounding atoms using lone pairs

Formal Charge (FC)

Difference between e assigned to atom and normal valence e

Sum = compound charge

FC vs Oxidation Number

FC: Underestimate ∆EN effect

Oxidation Number: Overestimate ∆EN effect (assume more EN atom has 100% of bonding e)

Resonance Structure

Structures with same atom arrangement and differing e placement

Resonance Hybrid

Actual structure = hybrid of all resonance structures

More stable structures contribute more

FC Stability

Small/No FC

Less separation between opposite FC

\-FC on more EN atoms

VSEPR Theory

3D atom arrangement determined by repulsions between bonding and nonbonding e pairs

Nonbonding e exert more repulsion (closer to nucleus)

Linear Bond Angles

180º

Trigonal Planar Bond Angles

120º

Tetrahedral Bond Angles

109\.5º

Trigonal Bipyramidal Bond Angles

90º, 120º, 180º

Octahedral Bond Angles

90º, 180º

Electronic Geometry

Spatial arrangement of all e pairs (bonding and lone)

Molecular Geometry

Spatial arrangement of bonding e pairs only

Determined by coordination number

Molecular Polarity

Nonpolar Bonds Only: Nonpolar molecule

Polar Bonds: Polar or nonpolar molecule, depend on molecular dipole

Molecular Dipole

Sum of bond dipoles

Cancel: Nonpolar molecule

Do not Cancel: Polar molecule

Intermolecular Forces (IMF)

Electrostatic interactions between molecules

Weaker than covalent bonds

Weakest to Strongest IMF

1. London dispersion forces (LDF)
2. Dipole-dipole interactions
3. H bonds

LDF

Rapid polarization of e density in nonpolar molecules form temporary dipoles to other dipoles to form

Temporary attraction

van der Waals force

LDF Strength

Weak from shifting dipoles

Only significant in close proximity

Large molecules easily polarized (stronger LDF)

Between gas, liquid, and solid molecules

Dipole-Dipole

Polar molecule dipoles attract

Longer duration than LDF

Dipole-Dipole Strength

Between liquid and solid molecules

H Bonds

Strong dipole-dipole interaction

IMF and intramolecular bonding

H (naked proton) to NOF

Compounds Forming H Bonds

Water

Alcohols

Amines

Carboxylic acids