chapter 8: covalent bonding - wunder...

Chapter 8

Chapter 8: Covalent Bonding

Chapter 8

• Two methods to gain or lose valence electrons:

– Transfer of Electrons = Ionic Bonding – Sharing of Electrons = Covalent Bonding

Bonding

Ionic Bonding - attracted

to each other, but not fully

committed

Covalent Bonding - fully

committed, and shares

everything

Chapter 8

• Prefixes used to denote number of atoms of element

• The suffix -ide is added to the more electronegative

element

• For the least electronegative (on the left), mono- is not

used to denote one atom.

Naming Covalent Compounds

Numerical Prefixes

mono- 1 hexa- 6

di- 2 hepta- 7

tri- 3 octa- 8

tetra- 4 nona- 9

penta- 5 deca- 10

Chapter 8

Molecules

• Molecule = covalently bonded atoms

• Diatomic molecule = two atoms

• Molecular compounds – lower melting and

boiling point than ionic compounds

• Molecular formula

– Doesn’t have to be lowest whole-number ratio

– Does represent structure

Chapter 8

Molecular Representations

Figure 8.5, Pg. 215, Text

Chapter 8

Lewis dot structures of covalent

compounds

• In covalent compounds atoms share electrons.

We can use Lewis structures to help visualize

the molecules.

Lewis structures

•Multiple bonds must be considered.

• Will help determine molecular geometry.

• Will help explain polyatomic ions.

Chapter 8

Types of electrons

• Bonding pairs

• Two electrons that are shared between two

atoms. A covalent bond.

• Unshared pairs

• A pair of electrons that are not shared

between two atoms. Lone pairs or nonbonding

electrons.

H Cl

oo

oo

oo

oo

Bonding pair

Unshared

pair

Chapter 8

Single covalent bonds

H H

H

C H H

H

Do atoms (except H) have octets?

F F

Chapter 8

Covalent Bonding

• When two similar atoms bond,

none of them wants to lose or

gain electrons

– Share pairs of electrons to each

obtain noble gas e-

configuration.

• Each pair of shared electrons

= one covalent bond

• Unshared Pairs = Pairs of e-

not shared by all atoms

– Show unshared pairs as dots

Visual, Pg. 219, Text

N H

H H

Chapter 8

Sharing of Electrons

• Water - H2O - Covalent Bonding – H = 1e-, O = 6e-

– Each H shares their electron

with O O = 8e- = [Ne]

– O shares 1e- with each H

H = 2e- = [He]

– Use dashes or dots to show

covalent bonds

Visual, Pg. 218, Text

H H O

Chapter 8

Multiple Covalent Bonds

• Elements can share more

than two electrons - creates

double and triple bonds

– Carbon Dioxide (CO2) - 4

electrons shared between the

carbon and the oxygens

(double bond)

O C O

O C O

Chapter 8

– Triple bond = Six electrons shared between the carbon atoms

– Ethyne (C2H2), a.k.a acetylene

– Bond length decreases from single to triple bonds

– Bond strength increases from single to triple bonds

Triple Covalent Bonds

C C H H C C H H

Chapter 8

Coordinate Covalent Bonds

• Coordinate covalent bond = one atom

contributes both bonding electrons

• Polyatomic ions contain coordinate covalent

bonds

C O O = 8e-

C = 6e-

O must contribute one of its pairs C O O = 8e-

C = 8e-

Chapter 8

O

OO

Resonance Structures

• Structures with multiple bonds can have similar

structures with the multiple bonds between different pairs

of atoms = Resonance Structures

• Example: Ozone has two identical bonds whereas the

Lewis Structure requires one single (longer) and one

double bond (shorter).

Chapter 8

Resonance Structures

Resonance Structures

Chapter 8

• Sometimes we can have two or more

equivalent Lewis structures for a molecule.

O - S = O O = S - O

• They both - satisfy the octet rule

- have the same number of bonds

- have the same types of bonds

• Which is right?

Resonance structures

Chapter 8

• They both are!

O - S = O O = S - O

O S O

This results in an average of 1.5 bonds between

each S and O.

Resonance structures

Chapter 8

• Three classes of exceptions to the octet rule for

molecules

• Odd number of electrons

• One atom has less than an octet

• One atom has more than an octet

• Odd Number of Electrons

• Few examples.

• Molecules such as ClO2, NO, and NO2 have an odd

number of electrons.

N O N O

Exceptions to Octet Rule

Chapter 8

• Less than an Octet

• Relatively rare.

• Typical for compounds of Groups 13.

• Most typical BF3.

• More than an Octet

• Atoms from the 3rd period onwards can accommodate

more than an octet.

• The d-orbitals are low enough in energy to participate in

bonding and accept extra electron density.

Exceptions to Octet Rule

Chapter 8

Atoms with fewer than eight electrons

• Beryllium and boron will both form

compounds where they have less than

8 electrons around them.

:Cl:Be:Cl: : :

: : :F:B:F:

:F:

: :

: :

: :

Chapter 8

Species with an odd

total number of electrons

• Example - NO

• Nitrogen monoxide is an example of a

compound with an odd number of electrons.

• It has a total of 11 valence electrons: six

from oxygen and 5 from nitrogen.

• The best Lewis structure for NO is:

:N::O:

: .

Chapter 8

Drawing Lewis structures

• Write the symbols for the elements in the correct

structural order.

• Calculate the number of valence electrons for all atoms

in the compound.

• Put a pair of electrons between each symbol, the bond

between each.

• Beginning with the outer atoms, place pairs of electrons

around atoms until each has eight (except for hydrogen).

• If an atom other than hydrogen has less than eight

electrons, move unshared pairs to form multiple bonds.

Chapter 8

Lewis Dot Structures: SO42-

• 1. Write a possible

– arrangement.

• 2. Total the electrons. • 6 from S, 4 x 6 from O

• add 2 for charge

– total = 32

• 3. Spread the electrons

– around.

S O

O

O

O

- - |

|

S O

O

O

O

Chapter 8

Lewis structures

• Example CO2

• Step 1

– Draw any possible structures

C-O-O O-C-O

You may want to use lines for bonds.

Each line represents 2 electrons.

Chapter 8

Lewis structures

• Step 2

– Determine the total number of valence electrons.

– CO2 1 carbon x 4 electrons = 4

2 oxygen x 6 electrons = 12

Total electrons = 16

Chapter 8

Lewis structures

• Step 3

– Try to satisfy the octet rule for each atom

• - all electrons must be in pairs

• - make multiple bonds as required

Try the C-O-O structure

No matter what you

try, there is no way

satisfy the octet for

all of the atoms.

C O O

Chapter 8

Lewis structures

O=C=O

This arrangement needs

too many electrons.

How about making some double bonds?

That works!

O C O

= is a double bond,

the same as 4 electrons

Chapter 8

Ammonia, NH3

H

H N H

Step 1 Step 2 3 e- from H

5 e- from N

8 e- total

Step 3 N has octet

H has 2 electrons

(all it can hold)

H

H N H

Chapter 8

Electronegativity

• Electronegativity = Ability to attract

electrons in a chemical bond

– Decrease as you move down a group

– Increase as you move from left to right

across a period.

Decreasing

Electro-

negativity

Increasing

Electro-

negativity

Chapter 8

Polar Covalent Bonds

• Sharing of electrons in a covalent bond does not imply equal sharing of those electrons.

• In some covalent bonds - electrons located closer to one atom than the other

• Unequal sharing polar bonds.

• Sharing based on electronegativity of elements in bond

Chapter 8

Electronegativities

Chapter 8

Electronegativity and Polarity

• Difference in electronegativity between atoms is

a gauge of bond polarity

Difference Type of Bond Electrons

0 – 0.4 Nonpolar Covalent Equal sharing

0.4 – 1.0 Moderately Polar

Covalent

Unequal sharing

1.0 – 2.0 Very Polar

Covalent

Unequal sharing

3 Ionic Transfer

Nonpolar and polar covalent bonds

• Nonpolar: When two atoms share a pair of

electrons equally.

• H H Cl Cl

• Polar: A covalent bond in which the electron

pair in not shared equally.

H Cl

• Note: A line can be used to represent a

shared pair of electrons.

d+ d-

oo

oo

oo

oo

oo

oo

oo o

o

oo

oo

oo

oo

Chapter 8

• The positive end (or pole) in a polar

bond is represented d+ and the

negative pole d- = dipoles (partial

charges)

• Polar molecules placed between

electric (+/-) plates – become

aligned with plates

Polar Dipoles

H H O

δ+ δ+

δ-

Chapter 8

Attractions between Molecules

• Intermolecular attractions weaker than ionic or

covalent bonds

• these are attractions BETWEEN molecules not

inside molecules like ionic or covalent bonds

(intramolecular)

Chapter 8

Types of Intermolecular Attractions

– van der Waals forces

• Dipole interaction

– δ+ end of one molecule attracted to δ- end of another

• Dispersion forces

– Caused by movements of e-

– Hydrogen Bonds

• Hydrogen bonded with very electronegative element is also

weakly bonded to an unshared pair on another

electronegative atom

Chapter 8

Characteristics of Ionic and

Covalent Compounds Characteristic Ionic Compound Covalent Compound

Representative Unit Formula Unit Molecule

Bond Formation Transfer of e- Sharing of e-

Type of elements Metal + Nonmetal Nonmetals

Physical State Solid Solid, liquid, or gas

Melting Point High (usually above

300°C)

Low (usually below

300°C)

Solubility in water High High to low

Electrical conductivity

of aqueous solution

Good conductor Poor to

nonconducting

Table 8.4, Pg 244, Text

Chemistry

Molecular Geometries

Adapted from

John D. Bookstaver

St. Charles Community College

St. Peters, MO

2006, Prentice-Hall, Inc.

Molecular Shapes

• The shape of a

molecule plays an

important role in its

reactivity.

• By noting the number

of bonding and

nonbonding electron

pairs we can easily

predict the shape of

the molecule.

What Determines the Shape of a Molecule?

• Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.

• By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

Valence Shell Electron Pair Repulsion

Theory (VSEPR)

“The best

arrangement of a

given number of

electron domains is

the one that

minimizes the

repulsions among

them.”

Molecular Geometries

• The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

Linear Electron Domain

• In this domain, there is only one molecular

geometry: linear.

• NOTE: If there are only two atoms in the

molecule, the molecule will be linear no

matter what the electron domain is.

Trigonal Planar Electron Domain

• There are two molecular geometries:

– Trigonal planar, if all the electron domains are

bonding

– Bent, if one of the domains is a nonbonding pair.

Polarity

• In Chapter 8 we

discussed bond dipoles.

• But just because a

molecule possesses

polar bonds does not

mean the molecule as a

whole will be polar.

Polarity

By adding the

individual bond

dipoles, one can

determine the

overall dipole

moment for the

molecule.

Overlap and Bonding

• We think of covalent bonds forming through the

sharing of electrons by adjacent atoms.

• In such an approach this can only occur when

orbitals on the two atoms overlap.