1. Polar Bonds and Molecules
Section 8.4
Polar Bonds and Molecules

2. Objectives After studying this section you should be able to:
Differentiate among 4 types of bonds
Explain the importance of electronegativity in polar molecules.
Determine the electronegativity of atoms
Calculate the difference in electronegativity between atoms
Identify and provide examples for three types of intermolecular forces
Describe the relationship between bonding and the physical characteristics of compounds

3. Bond Polarity Non-polar Covalent Bond:
When electrons are shared equally between two atoms.
Formed between diatomic halogen molecules, H2, O2, and N2

4. 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.

5. Water is an example of a polar molecule!

6. Remember electronegativity?
Definition: The ability of an atom to attract electrons when the atom is in a compound.

7. Electrons will “side” with the more electronegative atom
But which atom is more electronegative?

8. Electronegativity increases
Electronegativity decreases

9. So, which is more electronegative?
Hydrogen or Bromine?
Bromine!
Chlorine or Fluorine?
Fluorine!
Oxygen or Lithium?
Oxygen!
Hydrogen or Oxygen?

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

11. 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.

12. Use the table below to determine the electronegativity differences:Li Be B C N O F Na Mg Al Si P S Cl K Ca Ga Ge As Se Br Rb Sr In Sn Sb Te I Cs Ba Tl Pb Bi

13. Calculate the electronegativity differences between:
Hydrogen & Fluorine:
= 1.9
Potassium & Chlorine:
3.0 – 0.8 = 2.2
To figure out the type of bond that forms, look at the following table!

14. 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)

15. Place the following bonds in order from least to most polar:
H – Cl
H – S
H – Br
H – C
Answer: B and D are tied, C, A

16. Identify the types of bonds:
H and Br
C and O
Cl and F
Li and O
Br and Br

17. 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

18. Polar Molecules A molecule that has two poles is called a dipole.
Polar bonds don’t always make the whole molecule polar!

19. 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

20. 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

21. 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!

22. Types of Intermolecular Forces
Van der Waals Forces
Hydrogen Bonds

23. Van der Waals Forces Two kinds: a. Dipole Interactions
b. Dispersion Forces

24. 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

25. Dipole Interactions

26. 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.

27. Symmetrical Electron Distribution
Asymmetrical Electron Distribution

28. 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

29. 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.

30. Example: Bromine Has 35 electrons—a high number
Dispersion force between molecules is strong
Bromine is a liquid

31. 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

32. Another Intermolecular Force: Hydrogen Bonds

33. 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.

34. Bonds determine the physical properties of a compound!

35. 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
Electrical Conductivity of aqueous solution
Good conductor
Poor to nonconducting