Section 8.4 Polar Bonds and Molecules. Objectives After studying this section you should be able to: 1.Differentiate among 4 types of bonds 2.Explain.

Section 8.4

Polar Bonds and Molecules

Objectives

After studying this section you should be able to:1. Differentiate among 4 types of bonds2. Explain the importance of electronegativity in polar

molecules.3. Determine the electronegativity of atoms4. Calculate the difference in electronegativity between

atoms5. Identify and provide examples for three types of

intermolecular forces6. Describe the relationship between bonding and the

physical characteristics of compounds

Bond Polarity

• Non-polar Covalent Bond:

When electrons are shared equally between two atoms.

Formed between diatomic halogen molecules, H2, O2, and N2

Polar Covalent Bond / Polar Bond

• Electrons shared unequally between two atoms.

• The atom with greater electronegativity attracts the electrons, so electrons spend more time with that atom

• That side of the molecule is more negative than the rest.

• The other side of the molecule is positive.

Water is an example of a polar molecule!

Remember electronegativity?

• Definition: The ability of an atom to attract electrons when the atom is in a compound.

Electrons will “side” with the more electronegative atom

But which atom is more

electronegative?

Electronegativity increases

Electronegativity decreases

So, which is more electronegative?

• Hydrogen or Bromine?Bromine!

• Chlorine or Fluorine?Fluorine!

• Oxygen or Lithium?Oxygen!

• Hydrogen or Oxygen?Oxygen!

How do we notate which atom in a compound is more electronegative?

OR

Types of Bonds

• The type of bond that forms between two atoms depends on their differences of electronegativity.

• Types of bonds are: nonpolar covalent, moderately polar covalent, very polar covalent, and ionic.

Use the table below to determine the electronegativity differences:

H 2.1

Li 1.0

Be 1.5

B 2.0

C 2.5

N 3.0

O 3.5

F 4.0

Na 0.9

Mg 1.2

Al 1.5

Si 1.8

P 2.1

S 2.5

Cl 3.0

K 0.8

Ca 1.0

Ga 1.6

Ge 1.8

As 2.0

Se 2.4

Br 2.8

Rb 0.8

Sr 1.0

In 1.7

Sn 1.8

Sb 1.9

Te 2.1

I 2.5

Cs 0.7

Ba 0.9

Tl 1.8

Pb 1.9

Bi 1.9

Calculate the electronegativity differences between:

• Hydrogen & Fluorine:

4 - 2.1 = 1.9

• Potassium & Chlorine:

3.0 – 0.8 = 2.2

• To figure out the type of bond that forms, look at the following table!

Electronegativity and Bond Types

Electronegativity difference 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)

Place the following bonds in order from least to most polar:

A. H – Cl

B. H – S

C. H – Br

D. H – C

Answer: B and D are tied, C, A

Identify the types of bonds:

1. H and Br

2. C and O

3. Cl and F

4. Li and O

5. Br and Br

Answers!

• H and Br form a moderately polar covalent bond

• C and O form a moderately to very polar covalent bond

• Cl and F form a moderately to very polar covalent bond

• Li and O form an ionic bond

• Br and Br form a nonpolar covalent bond

Polar Molecules

• A molecule that has two poles is called a dipole.

• Polar bonds don’t always make the whole molecule polar!

Shape Matters!

CO2 is non-polar even though there are two polar bonds: the shape is linear and the bonds pull in equal and opposite directions, so they cancel each other.

O=C=O

Look at H2O:

Water is a bent molecule. The polar bonds do not pull in equal and opposite directions. Therefore, water IS a POLAR molecule

Attractions between molecules

• Called intermolecular forces

• They are Weaker than ionic or covalent bonds.

• Yet, they are still important and, yes, you must know them!

Types of Intermolecular Forces

1. Van der Waals Forces

2. Hydrogen Bonds

Van der Waals Forces

• Two kinds:

a. Dipole Interactions

b. Dispersion Forces

Dipole Interactions

• Occurs when polar molecules are attracted to one another.

• The negative region of one molecule is attracted to the positive region of the other molecule

Dipole Interactions

Dispersion Forces

• Weakest of all intermolecular forces

• Caused by the movement of electrons

• Because of the constant motion of electrons a molecule can develop a temporary dipole when the electrons are distributed asymmetrically around the nucleus.

Symmetrical Electron Distribution Asymmetrical Electron Distribution

More on dispersion forces

• They grow stronger as the number of electrons increases

• They have an effect on the state of the compound—solid, liquid, or gas

Example: Chlorine

• Has 17 electrons—a relatively low number

• Thus, the strength of the dispersion force between Chlorine molecules are low.

• Because of the weak

dispersion force,

chlorine is a gas.

Example: Bromine

• Has 35 electrons—a high number

• Dispersion force between molecules is strong

• Bromine is a liquid

What about Iodine?

• How many electrons?53

• Is this a high or low number?High

• Look at the periodic table: What state of matter is iodine found in?Solid

Another Intermolecular Force: Hydrogen Bonds

Hydrogen Bonds in Words

• Involves a hydrogen that is covalently bonded to a very electronegative atom (within a compound) and is bonded to an unshared electron pair another electronegative atom.

• Strongest of intermolecular forces• Important in determining the properties of

water and biological molecules such as proteins.

Bonds determine the physical properties of a compound!

Characteristic Ionic Comp. Covalent Comp.

Rep. Unit Formula Unit Molecule

Bond Formation Transfer of e- Sharing e-

Type of Elements Metallic & non-metallic

Nonmetallic

Physical State Solid S, L, or Gas

Melting Point High > 300º High to Low

Solubility in H2O Usually High High to Low

Electrical Conductivity of aqueous solution

Good conductor Poor to nonconducting