COVALENT BONDINGChapter 8

What is a covalent bond?

Why do atoms bond? Atoms bond to become more stable. Atoms are most stable when they have 8

valence electrons. (i.e. the same electron configuration as the closest noble gas.)

A chemical bond that results from two atoms sharing electrons is called a covalent bond.

Fluorine

Fluorine, and some other common elements, usually exist as a molecule of F2.

Fluorine’s electron configuration is….

Fluorine’s Lewis structure is…

Lewis Structures

1. Add up all the valence electrons involved.2. Determine which atom is the central atom.3. Form single bonds between the central

atom and all of the surrounding atoms. (each single bond uses 2 valence electrons)

4. Place electrons around the surrounding atoms until all have eight or until we have run out of valence electrons.

5. If there are any remaining valence electrons place them on the central atom.

Draw the lewis structures for the following molecules.

PH3

HCl

PBr3

CCl4

SiH4

NH3

NF3

Single Covalent Bonds

When only one pair of electrons are shared (like in F2) it is called a single bond.

Single bonds are also called sigma (σ) bonds Group 17:

The halogens all have seven valence electrons and will therefore be able to form one covalent bond.

Group 16: The elements in group 16 have six valence electrons and can

therefore make two single covalent bonds. Group 15:

The elements in group 15 have five valence electrons and can therefore make three single covalent bonds.

Group 14: The elements in group 14 have four valence electrons and can

therefore make four single covalent bonds.

Double and Triple bonds

Some elements are able to share more than one pair of electrons in order to obtain eight valence electrons.

A double covalent bond forms when two atoms share two pairs of electrons.

A triple covalent bond forms when two atoms share three pairs of electrons.

One bond in a multiple bond molecule is still called a sigma bond and the second and/or third are called pi (π) bonds.

Lewis Structures of Polyatomic Ions

To draw a lewis structure of a polyatomic ion we first figure out how many valence electrons we have, then add or subtract electrons to account for the charge.

SO42-

Bond Strength

The strength of a covalent bond depends on the distance between the bonded nuclei.

F2 Single covalent bond Bond length 1.43 x 10-10 m

O2

Double covalent bond Bond length 1.21 x 10-10 m

N2

Triple covalent bond Bond length 1.10 x 10-10 m

Exceptions to the octet rule

A small group of molecules have have an odd number of valence electrons and be unable to form an octet around each atom.

NO2

An other group of molecules have expanded octets.

SF6

Naming Molecules

1. The first element in the formula is always named first, using the entire name of the element.

2. The second element in the formula is named using its root and adding the ending –ide.

3. Prefixes are used to indicate the number of atoms of each element that are present in the compound.

NO2 – Nitrogen Dioxide N2O2 – Dinitrogen Dioxide

Resonance Structures

Sometimes its possible to have more than one correct lewis structure for a molecule.

In these cases all of the possible lewis structures are called resonance structures.

NO3-

SO2

O3

Naming Practice

CO2

SO2

NF3

CCl4

What is the formula of Diarsneic trioxide?

Naming Acids

Binary acids (meaning molecules with just hydrogen bonded to another single atom)

The first word has the prefix – hydro – the rest of the first word consists of a from of the name of the second atom followed by the prefix – ic.

HCl – hydrochloric acid HF – hydrofluoric acid HI – hydroiodic acid

Naming Oxyacids

Oxyacids are acids that contain hydrogen bonded to a polyatomic anion.

The first word of the name of these acids consists of the root of the polyatomic ion. If the polyatomic ion’s name ends in –ate you replace it with the ending –ic. If the polyatomic’s name ends in –ite you replace it with –ous.

HClO3 – Chloric acid HClO2 – Chlourous acid HNO3 – Nitric acid HNO2 – Nitrous acid

Naming Acids Practice

HI

HClO3

H3PO4

H2SO4

H2S

Molecular Shapes

The molecular geometry, or shape, of a molecule can be determined once a lewis structure is drawn.

The model used to determine a molecules shape is the VSEPR model

V = Valence S = Shell E = Electron P = Pair R = Repulsion

Bond Angle

Each bond, or set of nonbonding electrons, repels the other electrons in a molecule.

Molecules form the shape that results in the smallest interactions between the electrons in the molecule.

The shape of a molecule is dependant on what the eight, or sometimes more, electrons around the central atom are doing.

The VSEPR Model

A covalent bond forms between two atoms when a pair of electrons occupies the space between the atoms. This is a bonding pair of electrons The region these electrons occupy is called

the electron domain. A nonbonding pair of electrons (Lone

pair) reside on only one atom. Example NH3: How many bonding pairs?

How many nonbonding pairs?

VSEPR predicts that the best arrangement of electron domains is the one that minimizes the repulsion among them. The arrangement of electron domains about

the central atom of an Abn molecule is it’s electron domain geometry.

There are five different electron domain geometries. Linear, Trigonal planar, Tetrahedral, trigonal

bipyramidal, and octahedral.

Linear Example CO2

Trigonal PlanarExample BH3

Tetrahedral Example CH4

Trigonal BipyramidExample PCl5

OctahedralExample SF6

Effect of Nonbonding Electrons on Molecular Geometry

Consider CH4, NH3 and H2O

All three have the same electron domain geometry but have different molecular geometry and bond angles.

Electron domains for nonbonding electron pairs exert greater repulsive forces on other electron domains. This is why the bond angles decrease as the number of nonbonding electron domains increases.

Electronegativity and Polarity

The type of bond formed during a reaction is related to each atoms attraction for electrons.

Electron affinity is a measure of the tendency of an atom to accept and electron.

Electron affinity increases as you move from left to right on the periodic table and from bottom to top.

The difference between electronegativity determines what type of bond is formed.

Bond Type by ElectronegativityElectronegativity Difference Bond Type

> 1.7 Ionic Bond

0.4 – 1.7 Polar Covalent Bond

< 0.4 Nonpolar Covalent Bond

Shaken not stirred James Bond

Polar Covalent Bonds

Because not all atoms attract electrons equally sometimes we have one atom in a bond being “greedy”.

HCl Electronegativity of Cl – 3.16 Electronegativity of H – 2.20 Difference – 0.96