Covalent BondingResults from the

sharing of electron pairs between two non metal atoms

Different Covalent Different Covalent BondsBonds

Nonpolar-covalentNonpolar-covalentBonding electrons are shared

equally by the bonded atoms, resulting in a balanced distribution of electrical charge (EN = 0 - 0.4)

Polar-covalentPolar-covalentBonded atoms have an unequal

attraction for the shared electrons (EN = 0.4 – 1.7)

MoleculeA NEUTRAL group of atoms that

are held together by covalent bonds

Diatomic molecules: molecules containing only 2 atoms

Characteristics of a Covalent Bond

• Bond length: The distance between two bonded atoms (average distance between bonded atoms)

• Bond energy: The energy required to break a chemical bond and form neutral isolated atoms

• Happens between nonmetals (nonmetallic)

The Octet Rule

Chemical compounds tend to form so that each atom, by gaining,

losing, or sharing electrons, has an octet of electrons in its

highest occupied energy level

Octet = 8 = s2 + p6

Orbitals Overlap

Electron-Dot Notation

Diagrams that show only valence electrons as dots placed around the

element's symbol

Inner shell electrons are not shown

Lewis StructuresFormulas representing covalent bondsAtomic symbol = nuclei & inner-shell

electronsDot-pairs or dashes = shared

electron pairs in covalent bondsAdjacent dots = unshared electrons

Lewis Structurescontinued…

Structural formula: indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a moleculeSingle bond: a covalent bond in which one pair of electrons is shared between two atoms

How to Write Single-bond Lewis Structures

1. Determine the type and # of atoms in the molecule2. Write the electron-dot notation for each type of atom in the

molecule3. Determine the total number of valence electrons available

in the atom to be combined4. Arrange the atoms to form a skeleton structure for the

molecule. (When carbon present always the central atom, Hydrogen is never central, least-electronegative center when no carbon)

5. Add unshared pairs of electrons to each nonmetal atom (except H) such that each is surrounded by 8 e-

6. Count the e- in the structure to be sure the # of valence e- used equals the # available

*See handout and do some practice……

Multiple Covalent BondsDouble bond = two pairs of electrons are

shared between two atomsTriple bond = three pairs of electrons are

shared between two atomsBoth have greater bond energies and

shorter bond lengths than single bonds

How to Write Multiple Bond Lewis Structures

1. Follow steps 1-3 from single-bond directions

2. Arrange the atoms to form a skeleton structure for the molecule and connect to atoms by e- pair bonds

3. Add unshared pairs of electrons to each nonmetal atom (except hydrogen) such that each is surrounded by eight electrons

4. Count the electrons in the Lewis structure to be sure that the # of valence electrons used equals the # available

5. If too many electrons have been used, subtract one or more lone pairs until the total number of valence e- is correct. Then move one or more lone e- pairs to existing bonds between non-hydrogen atoms until the outer shells of all atoms are completely filled

Bond Length vs Bond Energy

Bond length decreases as the strength of the bond (bond energy) increases

Small bond length = large bond energy

Large bond length = small bond energy

Resonance StructuresResonance = bonding in molecules

of or ions that cannot be correctly represented by a single Lewis structure

A double-headed arrow is place

between a molecule’s resonance structures

Polyatomic IonsA charged group of

COVALENTLY bonded atoms

Network solids (covalent crystals)Network solids (covalent crystals)There are some compounds that do not have molecules, but instead are long chains of covalent bonds (E.g. diamond)

C C C

C C C

C C C C C

C

C CCC

C

This happens in 3 dimensions, creating a crystalBecause there are only covalent bonds, network solids are extraordinarily strong

3-dimensional arrangement of a

molecule’s atoms in space

Molecul

ar

Geometry

Johannes van der WaalsDutch Physicist,

1837-1923Developed

understanding of intermolecular attractions

Won Physics Nobel Prize in 1910

VSEPR TheoryValence-shell, Electron-pair, Valence-shell, Electron-pair,

RepulsionRepulsion

Electron Pairs Repel Electron Pairs Repel Take positions to maximize Take positions to maximize

separation ... separation ...

minimize repulsions minimize repulsions

Explains shape but not how orbitals Explains shape but not how orbitals change when bonding occurschange when bonding occurs

HybridizationMixing of 2 or more atomic

orbitals to produce new hybrid (changed) atomic orbitals

Hybrid orbitals are named for the atomic orbitals that the

electrons used to occupy

Old orbitals energy = New hybrid orbitals energy

Hybridization ExamplePurpose is to form equivalent bonds

CH4 = Methane

Carbon needs to bond equally to 4 Hydrogen's

Intermolecular Forces

The force of attraction between molecules

1.Dipole-Dipole Forces

2.Hydrogen Bonding

3.London Dispersion Forces

Dipole-Dipole• Dipole = opposite charges that are

equal and separated by a short distance (polar molecules exhibit dipoles)

• Dipole-Dipole Forces occur between the molecules

• Dipoles can induce polarity in a nonpolar molecule

Dipole-

Dipole

Hydrogen BondingThe strongest dipole-dipole force

Egg white is clear because of hydrogen bonding. But heat it up and break the bonds, and you end up with a white gelatinous solid.

London Dispersion ForcesConstant motion of electrons may cause an

instant uneven distribution of the electrons and the formation of a now positive pole

Present in ALL atoms and moleculesNoble gas atoms and nonpolar molecules

however can ONLY exhibit these forces

Ionic, H-bonding, Dipole, or Ionic, H-bonding, Dipole, or London?London?

Details Bond Molecule IMFEN = 0 - 0.5 nonpolar nonpolar London

EN = 0.5 - 1.7 polar polar dipole-dipole*EN = 1.7 - 3.2 ionic Ionic ionic*

H + N,O,F polar polar H-bonding*Symmetrical

molecule (any EN) -- nonpolar London

*Since all compounds have London forces. London forces are also present. However, their

affect is minor and overshadowed by the stronger forces present.