Covalent Bonding

Unit 6

Chapter 6

Haves and Have-Nots

In the late 1800’s and early 1900’s, scientists did not know very much about the reasons that certain chemicals reacted with others.

Knowledge of this reactivity, however, was shown to help in industrial and medicinal applications,

but there was no real understanding, only a repertoire of techniques.

Ionic Bonding is Like Magnets

We knew about electrons and could determine that cations lost electrons and anions gained electrons.

We understood why ions were attracted to each other

because we had physical theories of magnetism that could be used as models.

A Wrench Thrown in the Gears

Positive and negative charges attracted each other – this gave us ionic bonding.

Non-ionic bonding, when a nonmetal bonds with another nonmetal, was very different.

We had very little understanding of this phenomenon.

Gilbert the Octopus

Gilbert Newton Lewis studied the behavior

of many ionic and non-ionic compounds and

Came up with the Octet Rule in 1902.

He envisioned atoms as cubes. Gilbert Newton Lewis

(1875 – 1946)

Thinking Cubed

In Lewis’ model, electrons were stationary & formed the corners of the cubes.

A graduate student saw Lewis’ work &

Suggested that the corners were shared

This led to bonding. Lewis’ Original Diagrams

The Dot Before the Dot Com

In 1916, Lewis published his paper on “The Atom and the Molecule”

which discussed non-ionic bonding. In his paper, Lewis discussed ways to predict

non-ionic bonding using diagrams. These diagrams would later become known

as Lewis Dot Structures.

Languish and Langmuir

Unfortunately, Lewis was a poor communicator.

His ideas did not get very far. During WWI, he met Irving

Langmuir Who built on Lewis’ work and

published a paper in 1921 called “The Arrangement of Electrons in

Atoms and Molecules.”

Irving Langmuir(1881 – 1957)

I dunno Lewis…I’ll call it “Mine!”

In his paper, Langmuir coined the term “covalence”

to describe the sharing of electrons in non-ionic bonding.

He promoted the Octet Rule and Covalent Bonding so well that the theory was often known as

the Lewis-Langmuir theory (or simply Langmuir’s).

Lewis Dot Structures

We can imagine an atom like a square. According to the Octet Rule, an atom is stable

when it has 8 electrons surrounding it. Since there are four sides of a square, each

side can hold 2 electrons.

Ne

Lewis Dot Structures Continued

Number of dots surrounding an element is determined by the # of valence electrons.

Follow Hund’s Rule (from Quantum Mechanics) and put a dot on each side of the

Atomic symbol until we have to pair them up.

H B

C

N

O

F

NeBe

What’s it all mean?

Unpaired electrons are free to participate in bonding.

Pairs of electrons do not participate in bonding and are called Lone Pairs.

(Lone Pear)

Bonding with Valence Electrons

Covalent means “with valence” Bonds form between unpaired valence

electrons of adjacent atoms Atoms will only make as many bonds as there

are unpaired electrons

N HH

H

Already Paired UpDo not Bond!

Redox Redux

An Oxidation State is the charge an atom would have if the bonds within a molecule were completely ionic.

Many reactions are driven forward by a change in oxidation state.

When an atom’s oxidation state is increased (made more positive), it is oxidized.

When an atom’s oxidation state is decreased (made more negative), it is reduced.

IUPAC Rules!

Rules for determining Oxidation State:

1. Atoms in their elemental state have an oxidation state of 0.

2. Any simple monatomic ion has an oxidation state equivalent to the charge of the ion.

3. Hydrogen is (almost) always +1 and oxygen is (almost) always -2.

4. The sum of the oxidation states must equal the charge of the molecule/ion (0 in a neutral atom)

A Charge is a TypeOf Oxidation State!

In Practice

A monatomic ion’s charge is its oxidation state. Hydrogen is +1 unless it is bonded to an active metal

e.g. LiH, H = -1 Oxygen is -2 unless it is bonded to itself

e.g. peroxides, H2O2, O = -1 Representative elements typically acquire the same

oxidation as if they were ions. e.g. Alkali Metals = +1, Halogens = -1

If there is uncertainty, the most electronegative element gets the negative charge e.g. FCl, F = -1, Cl = +1

This is NOTCommon!

Oxidation Station

The sum of the oxidation states must equal the charge of the molecule/ion.

SCl4 S = ? Cl = -1 Charge = 0 1 sulfur plus 4 chlorines = 0 S + 4(-1) = 0 S = +4

PO4-3

P = ? O = -2 Charge = -3 P + 4(-2) = -3 P = +5

Make It So!

Determine the oxidation state of sulfur in:

Substance Equation Oxidation of S

1 S8 8(S) = 0 0 (elemental state)

2 H2S 2(+1) + S = 0 -2

3 SO2 S + 2(-2) = 0 +4

4 SO3-2 S + 3(-2) = -2 +4

5 H2SO4 2(+1) + S + 4(-2) = 0 +6

A Charge is a type of Oxidation State, butAn Oxidation State is not necessarily a Charge!