1. What the force?

2. Intramolecular Forces:
The attractive forces between atoms and ions within a molecule
e.g Ionic bond, covalent bonds (atoms lose, gain or share electrons to become more stable)
Strong
Intermolecular Forces:
The attractive forces between molecules
E.g. Van der Waals forces (London dispersion forces, dipole-dipole forces), hydrogen bonds
Weak (in comparison to intramolecular forces)
I.e. much less energy to melt H2O (inter) than for it to decompose it into H2 and O2 (intra)

3. Intermolecular Forces
If covalent bonds were the only forces at work, most molecular compounds would be gases as there would be no attraction between molecules strong enough to group the molecules as liquids or solids

4. van der Waals Forces
Dipole-dipole
London Dispersion

5. Dipole-dipole
Forces of attraction between oppositely charged ends of polar molecules.

6. London Dispersion
Attractive forces between all molecules, including nonpolar molecules
Result of temporary displacements of the electron cloud around atoms in a molecule (extremely short-lived dipoles)
Therefore weaker than dipole-dipole

7. Hydrogen Bonding
Strong dipole-dipole force between the positive H atom of one molecule and a highly electronegative atom of another molecule

8. Electronegativity and Polarity

9. Electronegativity: Attracting Electrons
When two atoms form a bond, each atom attracts the other atom’s electrons in addition to its own.
Electronegativity of an atom is a measure of an atom’s ability to attract electrons in a chemical bond.
As you move from left to right on the periodic table, the EN increases. As you move from top to bottom on the periodic table, the EN decreases.

10. Nonpolar Covalent Bond
When electrons are shared between 2 atoms, a covalent bond is formed.
If the atoms are identical, e.g. Cl2, the electrons are shared equally (nonpolar)

11. Polar Covalent Bond
If the electrons are shared between 2 different atoms, e.g. HBr, the sharing is unequal
The bonding electrons spend more time near the more electronegative atom H Br

12. Electronegativity Values

13. Electronegativity Differences
The absolute value of the difference in electronegativities of two bonded atoms provides a measure of polarity of a bond.
The greater the difference, the more polar the bond.
0 to 0.4
Nonpolar covalent
0.41 to 1.66
Polar covalent
> 1.67
Ionic
Electronegativity Difference

14. The ionic range is from 1.7 to 3.3.
The covalent range is from 0 to 0.5.
The polar covalent range is from 0.5 to 1.7.

15. Predicting Bond Type Using EN
You can use the differenced between EN to decide whether the bond between two atoms is ionic or covalent.
ΔEN = larger EN – smaller EN
For example, K-F
EN for K is 0.8
EN for F is 4.0
4.0 – 0.8 = 3.2
According to the diagram, KF is ionic.

16. What type of bond is N-O? EN for N is 3.0 EN for O is 3.5
3.5 – 3.0 = 0.5
According to the diagram, the bond is mostly covalent.

17. Polar Covalent Bonds (The in-between bonds)
When two bonding atoms have an EN difference that is greater than 0.5, but less than 1.7, they are considered to be polar covalent bonds.
The difference is not great enough for the more EN atom to take the electrons from the less EN atom.
The difference is great enough though, for the bonding electrons to spend more time near the more EN atom.

18. Look at the H-O bond in water.
EN for O is 3.5
EN for H is 2.1
Difference is 1.4 and this falls in the polar covalent region.

19. Which bond is more polar?
B – Cl or C – Cl?
The difference in electronegativities between chlorine and boron is 1.0 (3.0 – 2.0 = 1.0).
The difference in electronegativities between chlorine and carbon is 0.5 (3.0 – 2.5 = 0.5).
The B – Cl bond is more polar and the chlorine atom will hold the partial negative charge because of its higher electronegativity.

20. Which bond is more polar?
P – F or P – Cl?
The difference in electronegativities between fluorine and phosphorus is 1.9 (4.0 – 2.1 = 1.9).
The difference in electronegativities between chlorine and phosphorus is 0.9 (3.0 – 2.1 = 0.9).
The P – F bond is more polar and the fluorine atom will hold the partial negative charge because of its higher electronegativity.

21. Polar Molecules
Note: Not all molecules with polar bonds are polar molecules
Can you suggest a reason for their state at room temperature based on polarity?

22. Metallic Bonding
When metals combine with other metals, they do not have an EN greater than 1.7, does that mean that they are covalent?
Sodium metal is soft enough to be cut with a butter knife, copper and gold can be drawn into thin wires and hammered into sheets (ionic bonds are brittle), does this mean they have covalent bonds?
No. Metals do not have enough valance electrons to achieve stable configurations by sharing electrons.

23. The force that holds metal atoms together is called a metallic bond.
Unlike covalent and ionic bonding, metallic bonding does not have a particular orientation in space. Electrons are free to move and the metal ions are not rigidly held in formation. This is why a hammer can pound metal – the atoms can slide past each other.