Covalent Bonding

Ch. 16

The Nature of CovalentBonding

16-1

Skip pgs 444 - 451

Covalent Bonds• Covalent (molecular) bond = the attraction

of two atoms for a shared pair of electrons– Neither atom will have an ionic charge– Usually between 2 nonmetals (some involve

metalloids)!

• Covalent compound = a compound whose atoms are held together by covalent bonds

• Molecule = an uncharged group of two or more atoms held together by covalent bonds

Single Covalent Bonds

• Single Covalent Bond = 2 atoms share one pair of electrons.

– H2, F2, H2O

• Structural Formula = chemical formula that shows the arrangement of atoms.

– H + H H H H - H

– F + F F F F - F

– H + O O H O - H

H H H

F2:

H2O:

H

CH4 C + 4H H C HH

H *H C H

H

NH3 N + 3H H N HH

H N HH

Double and TripleCovalent Bonds

• Double Covalent bonds = bonds that involve 2 shared pairs of electrons.

– O2 O O

• Triple Covalent bonds = bonds that involve 3 shared pairs of electrons.

– N2 N N

O2 O + O O O O O

N2 N + N N N N N

Bonding Theories

16-2

Skip pgs. 452–454 + 457-459

VSEPR Theory

• VSEPR Theory states that because electron pairs repel, molecular shape adjusts so that valence-electron pairs are as far apart as possible.– Ex: H2O bond is NOT linear!

2H + O O H H

VSEPR Geometrics

A = Central Atom, X = Attached Species, E = Lone Pair of e-’s on A

Total # of Attached Species

Species

Type

Molecular Geometry

Example

2 AX2 Linear CO2

4 AX4

AX3E

AX2E2

Tetrahedral

Pyramidal

Bent

CH4

NH3

H2O

A

AA

AA A

107° 105°

Triatomic120°

Triagonal Pyramidal Bent

Linear Example

Tetrahedral Example

Pyramidal Example

Bent Example

Polar Bonds +Molecules

16-3 Part I

Bond Polarity

• Bonding pairs of electrons are pulled, as in a tug-of-war, between nuclei of atoms sharing electrons.

• If bonding pairs are shared equally it is a nonpolar covalent bond.– Atoms will have equal electronegativities (pg. 405)– Ex: N2, O2, H2, Cl2, CO2

• If bonding pairs are shared unequally it is a polar covalent bond.– Atoms have unequal electronegativity.– H2O, HCl, CO

Polar Molecules• Polar molecule = one end

of molecule is slightly negative and other end is slightly positive.

Ex: HCl

Electronegativity:

H = 2.1, Cl = 3.0

Difference = 0.9

Ex: H2O

H = 2.1, O = 3.5

Difference = 1.4

Electronegativity Differences + Bond Types

Electronegativity Difference

Type of Bond Example

0.0 – 0.3 Nonpolar Covalent H – H (0.0)

0.4 – 1.0 Moderate Polar

Covalent

∂+ ∂-

H – Cl (0.9)

1.1 – 2.0 Very Polar Covalent

∂+ ∂-

H – F (1.9)

> 2.0 Ionic Na+Cl- (2.1)

• The polarity of a molecule depends on the shape + orientation of the bonds.– Ex: CO2 polarity

cancels out since it is linear = nonpolar molecule

– Ex: H2O poles add up due to its bent shape = polar molecule

• Polar:

F

H C H

H

H N H

H

O

H H

• Nonpolar:

H

H C H

H

O C O

H H

C C

H H

Attractions Between Molecules

16-3 Part II

Attractions between Molecules• van der Waals forces = weakest attractions

(ionic + covalent are stronger); consist of dispersion forces, dipole interactions and hydrogen bonds.1) Dispersion forces = weakest of all molecular

interactions; caused by motion of electrons.• Increases as the number of electrons increases.

• Halogen diatomic molecules (F2, Cl2, Br2, I2)

• Fluorine + Chlorine have weak dispersion forces (less electrons); thus are gases at STP.

• Bromine (more electrons) is a liquid at STP, and Iodine (most electrons) is a solid at STP.

2) Dipole Interactions = occurs when polar molecules are attracted to one another.– The slightly

negative region of a polar molecule is attracted to the slightly positive region of another polar molecule

– When placed in an electric field, dipole molecules become oriented with respect to (-) and (+) charge

3) Hydrogen bonds = attractive forces in which a hydrogen covalently bonded to a very electronegative atoms is also weakly bonded to an unshared electron pair of another electronegative atom.-Strongest of intermolecular (van der Waals) forces

Van Der Waals Forces

Summary

Comparing Ionic + Molecular Properties

Characteristic Ionic Cmpd Covalent Cmpd

Representative Unit Formula Unit Molecule

Bond Formation Transfer e-’s Share pairs e-’s

Type of Element Metal + nonmetal Nonmetal (possible metalloid)

Physical State Solid S, L, or G

Melting Point High (>300°C) Low (< 300°C)

Solubility in Water High High to Low

Electrical Conductivity as aqueous soln

Good Poor to none

Network Solids

• Most molecules are easy to break; however, a few molecular solids are very stable.

• Network Solids = solids in which all atoms are covalently bonded to each other.– Solid does not “melt” until 1000°C or higher, in

which it vaporizes without melting at all.– Ex: Diamond; made of carbon, each carbon

bonded to 4 other carbons

Diamond +Silicon carbide (SiC)