Chapter 8: Covalent Bonding

8.1: Molecular Compounds

Important Terms

• Monatomic: Made of only one

atom

• Covalent bond: SHARING

electrons (like a tug of war for

electrons between atoms)

Molecule

• Formed when a covalent bond

holds a group of atoms together

so their charge is NEUTRAL

Diatomic Molecule

• Molecule consisting of TWO

atoms only.

• Diatomic elements are certain

elements that can NOT exist

alone in nature.

• 7 of them: make a “7” on the P.T.

Molecular Compound

• A compound that consists of

molecules bonded together

(covalent bonds)

• Relatively lower melting and

boiling points than ionic

compounds (ionic bonds)

• Mostly TWO NONMETALS

IONIC vs. MOLECULAR

• High

melting/boiling

• Ionic bonds

• Formed between

metal and

nonmetal

• Lower

melting/boiling

• Covalent bonds

• Formed between

two or more

nonmetals

Molecular Formulas

• Show the number and kind of

each atom in a molecule.

• Example: H2O

–Two hydrogen atoms, one oxygen

atom

• When there is one atom, a “1” is

not written.

More - Molecular Formula

• Not necessarily lowest whole-

number ratios

• Example: Ethane C2H6

Chapter 8

8.2: The Nature of Covalent

Bonding

Ionic Compound Review

• Electrons are shared between cations (metals) and anions (nonmetals) so all atoms involved have 8 valence electrons

Single Covalent Bond

• When ONE pair of electrons is

shared (tug of war)

• Shown by a pair of dots or a dash

Double/Triple Covalent Bond

• Some compounds have to share more

than one pair of electrons to have 8

valence electrons in each atom.

• Example: O2

• Shares TWO sets of electrons (double

bond)

• The extra electrons are called lone or

unshared pairs

Single/Double/Triple Bonds

• Single = longest, weakest bond

• Triple = shortest, strongest bond

Diatomic Elements

• Do not exist alone

• Fluorine, Chlorine,

Bromine, Iodine,

Hydrogen, Nitrogen, and

Oxygen

• Can be represented by

electron dot structures

Lewis (Electron) Dot Structures

• Write dot structures

of individual atoms

• Determine how

many pairs of atoms

they will have to

share

• Use a dash to

represent a bond

Examples

• Ammonia: NH3

• Water: H2O

• Acetylene: C2H2

• Carbon Tetrachloride:

CCl4

• Dihydrogen Selenide:

H2Se

Structural vs. Molecular Formulas

• Structural formulas

• Dot

structure/shows

bonds

Example: Carbon

tetrachloride

• Molecular formula

• Symbols and

numbers

• Example: CaCl2

Bond Dissociation Energies

• Energy required to break a covalent bond

• Expressed in kJ/mol (energy/amount)

• Larger = stronger/more stable bond

Resonance

• Molecules with more than one possible electron dot structure

• Do not switch back and forth

• Molecules exist as a mixture (hybrid) of the resonance forms

• Example:

Exceptions to Octet Rule

• If valence electron total is odd, the octet

rule doesn’t work

• Some atoms do not require all 8 valence

electrons

• These molecules can exist in stable form

• You do not need to know specific

examples for your test.

Octet Exception Example

8.3

Bonding Theories

VSEPR Theory

• Electron dot structures/structural formulas

don’t show 3D structure.

• Based on electrons repelling one another.

• Electron pairs stay as far apart as possible.

No Unshared Pairs

• Linear (180°)

– BeCl2

• Trigonal Planar (120°)

– BCl3

No Unshared Pairs

• Tetrahedral (109.5°)

– CClF3

• Square Planar (90°)

With Unshared Pairs

• Bent (105°)

– H2O

• Pyramidal (107°)

– NH3

Examples

• H2S

• CO2

• PCl3• OCS

• H2CO

• CH4

• N2H4