COVALENT BONDING

LET’SFIRSTREVIEWIONICBONDING

In an IONIC bond,

electrons are lost or gained,

resulting in the formation of IONS

in ionic compounds.

FK

FK

FK

FK

FK

FK

FK+

_

The compound potassium fluoride

consists of potassium (K+) ions

and fluoride (F-) ions

Ionic Bonds are/have:

1. Metals and Nonmetals

2. Solids

3. High melting points

4. Large differences in electronegativity

COVALENT BONDS

Ms. Thompson

Covalent Bonding

“tug of war”

Share electrons

Molecules of 2 or more atoms

2 nonmetal molecules

Groups 4A, 5A, 6A, and 7A

Covalent Molecules

Properties

Two nonmetals

Low melting points

Low conductivity

Nonelectrolytes

Prefixes for Covalent Molecules

Prefix Number

Mono- 1

Di- 2

Tri- 3

Tetra- 4

Penta- 5

Hexa- 6

Hepta- 7

Octa- 8

Nona- 9

Deca- 10

Naming Molecular Compounds

First nonmetal is written as:

PREFIX + element name

If there is one atom, do NOT use “mono”

Second nonmetal is written as:

PREFIX + element name root + “ide”

If there is one atom, USE “mono”

Example:

IF5 Iodine Pentafluoride

H2O Dihydrogen Monoxide

SBr6 Sulfur Hexabromide

Examples

CO

CO2

SO2

CF4

Carbon monoxide

Carbon dioxide

Sulfur dioxide

Carbon tetrafluoride

So

what

are

covalent

bonds?

In covalent bonding,

atoms still want to achieve

a noble gas configuration

(the octet rule).

In covalent bonding,

atoms still want to achieve

a noble gas configuration

(the octet rule).

But rather than losing or gaining

electrons,

atoms now share an electron pair.

In covalent bonding,

atoms still want to achieve

a noble gas configuration

(the octet rule).

But rather than losing or gaining

electrons,

atoms now share an electron pair.

The shared electron pair

is called a bonding pair

Types of Covalent Bonds

Single: two atoms share one pair of electrons

Ex. H – H

Double: two shared pairs of electrons

Ex. O = O

Triple: three shared pairs of electrons

Ex. N = N

Cl2

Chlorine

forms

a

covalent

bond

with

itself

ClClHow

will

two

chlorine

atoms

react?

ClClEach chlorine atom wants to

gain one electron to achieve an octet

ClClNeither atom will give up an electron –chlorine is highly electronegative.

What’s the solution – what can they

do to achieve an octet?

ClCl

Cl Cl

Cl Cl

Cl Cl

Cl Cloctet

Cl Cl

circle the electrons for

each atom that completes

their octets

octet

Cl Cl

circle the electrons for

each atom that completes

their octets

The octet is achieved by

each atom sharing the

electron pair in the middle

Cl Cl

circle the electrons for

each atom that completes

their octets

The octet is achieved by

each atom sharing the

electron pair in the middle

Cl Cl

circle the electrons for

each atom that completes

their octets

This is the bonding pair

Cl Cl

circle the electrons for

each atom that completes

their octets

It is a single bonding pair

Cl Cl

circle the electrons for

each atom that completes

their octets

It is called a SINGLE BOND

O2

Oxygen is also one of the diatomic molecules

How will two oxygen atoms bond?

OO

OOEach atom has two unpaired electrons

OO

OO

OO

OO

OO

OO

Oxygen atoms are highly electronegative.

So both atoms want to gain two electrons.

OO

Oxygen atoms are highly electronegative.

So both atoms want to gain two electrons.

OO

OO

OO

OO

OO

OOBoth electron pairs are shared.

6 valence electrons

plus 2 shared electrons

= full octet

OO

6 valence electrons

plus 2 shared electrons

= full octet

OO

two bonding pairs,

OO

making a double bond

OO=For convenience, the double bond

can be shown as two dashes.

OO

OO=This is the oxygen molecule,

O2

Drawing the Lewis Structure

1. Sum the total number of valence electrons

2. The atom usually written first in the chemical formula is

the central atom in the Lewis structure

3. Complete the octet bonded to the central atom

4. If there are not enough electrons to give the central atom

an octet try multiple bonds

VSEPRValence Shell Electron Pair Repulsion Theory

States the shape of molecules is based upon the concept that electrons, being of like charge, will repel themselves to the greatest possible distances.

•Electron – pairs repel•Shows three-dimensional shape

•Electron pairs are as far apart as possible•Produces molecular geometry of shapes

Shapes of Molecules

AB2 must be either linear or bent: Examples of Linear molecules

Linear - No non-bonding electronsLinear Molecules have a bond angle = 180°

Water is a bent molecule

with bond angles of 104.5°

**Notice – the bond angle

decreases as the number of

non-bonding pairs

increases**

AB2E or

AB2E2 -

classification

H2O

Linear Molecules have a bond angle = 180°Bent molecules have a bond angle ≠ 180°

AB B

AB3 most common shapes place the B atoms at the

corners of an equilateral triangle:

Trigonal Planar

The A atom lies in the

same plane as the B

atoms (Flat)

Bond angle = 120°

No non-bonding electrons

The A atom lies above the plane of the B atom.

Pyramid with an equilateral triangle as the base.

Trigonal Pyramidal

AB4 is Tetrahedral

The carbon has 4 valence electrons and thus needs 4 more electrons from

four hydrogen atoms to complete its octet. The hydrogen atoms are as far

apart as possible at 109° bond angle. This is tetrahedral geometry. The

molecule is three dimensional.

H

CH

HH

Six pairs of electrons around the central atom

are based on the Octahedron structure.AB6

The central atom can be visualized as

being at the centre of an octahedron,

with the six electrons pointing to the

six vertices – all bond angles are 90°

Octahedral Square Pyramidal

BrF5

Square Planar

XeF4

Should be less than 90º90°

SF6

Further Examples:

Tutorial Questions :Draw Lewis structures and the molecular geometry of the following

molecules:

H3O+, NH4

+, CS2, SCl2

Shape Bonding-

pairs

Non-

bonding

pairs

Bond angle Examples

Linear 2 0 180 BeCl2, CO2, HCN,

C2H2

Trigonal

planar

3 0 120 BF3, SO3, NO3-,

CO32-, C2H4

Tetrahedral 4 0 109.5 NH4+, SO4

2-, PO43-,

Ni(CO)4, CH4

Trigonal

pyramidal

3 1 107 PH3, SO32-, NH3

Bent 2 2 104.5 H2S, SO2, H2O

Predicting the Shape of the Molecule

5. Sum the Number of Electron Domains around the Central

Atom in the Lewis Structure; Single = Double = Triple

Bonds = Non-Bonding Lone Pair of Electrons = One

Electron Domain

6. From the Total Number of Electron Domains, Predict the

Geometry and Bond Angle(s); 2 (Linear = 180º); 3 (Trigonal

Planar = 120º); 4 (Tetrahedral = 109.5º); 5 (Trigonal

Bipyramidal = 120º and 90º); 6 (Octahedral = 90º)

7. Lone Pair Electron Domains exert a greater repulsive force

than Bonding Domains. Electron Domains of Multiple

Bonds exert a greater repulsive force than Single Bonds.

Thus they tend to compress the bond angle.

Review

Ionic vs. covalent

Naming

Lewis dot structures (ionic and covalent)

Shapes/bond angles

Nonpolar vs. Polar Bonds

Nonpolar Polar Bond formed between two

similar atoms

Electrons are shared equally

Usually a diatomic molecule or between C & H

Ex. H2, N2, O2, F2, Cl2, I2, Br2

(Diatomic Molecules) CH4, C2H6, C3H8

Bond formed between two atoms of two differentelements

Electrons are shared unequally

Ex. HCl, H2O, CO

Intermolecular Forces

Dispersion Forces (aka van der Waals forces):

Weakest

Between two nonpolar molecules

Usually a diatomic molecule or

between C & H

Ex. F2, Cl2, C2H6, CH4

Intermolecular Forces Dipole-Dipole Forces:

Between two polar molecules

Slightly negative region is weakly attracted to a slightly positive region

Ex. HCl

Intermolecular Forces

Hydrogen Bonding:

A hydrogen bonds to a very electronegative atom (O, N, or F)

Ex. H2O

Ranking based on Polarity

Polarity increases with:

Increased intermolecular forces

Polarity also increases with:

Increased boiling point

Increased viscosity

Increased electronegativity

Decreased temperature

**Therefore, when ranking molecules by increasing boiling point, you are really ranking the molecules by increasing polarity or intermolecular forces.

Types of Intermolecular Forces Practice Wkst Answers1. Hydrogen

2. Dispersion

3. Hydrogen

4. Dispersion

5. Dipole-Dipole

6. Dispersion

7. Dispersion

8. Dipole-Dipole

9. Hydrogen

10. Dispersion

11. Dipole-Dipole

12. Dipole-Dipole

13. Hydrogen

14. Dispersion

15. Dipole-Dipole

Polarity Practice Worksheet

Polarity Practice Ranking by Inc Polarity

1. CHBr3

2. H2O

3. HI

4. C2HBr

5. CH3OH

1. NF3 < PF3 < SF2 < LiOH

2. N2H2 < C2H5OH , CH3OH < Ni(OH)3

3. B2F4 < H2C2O4 < CF2O < CuCl2

4. PH3 < NF3 < PF3 < NH3

5. H2 < H2S < H2O < HF

Intermolecular Forces Wkst

1. Dispersion

2. Hydrogen

3. Dispersion

4. Dipole-Dipole

5. Dispersion

6. Dipole-Dipole

7. Dipole- Dipole

8. Hydrogen

11. C2H6 < C2H5 < C2H5OH

12. H2 <H2S < H2O

13. BI3 < BBr3 < BCl3

14. CH4 < CH3OCH3 < CH4O < CaCO3

Electronegativity &Nonpolar vs. Polar Bonds

ElectronegativityDifference Range

Type of Bond Example

0.0 – 0.4 Nonpolar Covalent H – H (0.0)

0.4 – 1.0 Moderately Polar Covalent

H – Cl (0.9)

1.0 – 2.0 Very Polar Covalent H – F (1.9)

≥ 2.0 Ionic Na+ Cl- (2.1)