1. Intermolecular Forces Chemistry 11 Ms. McGrath

2. Intermolecular Forces The forces that bond atoms to each other within a molecule are called intramolecular forces. Covalent bonds are an example of intramolecular forces. In comparison, the forces that bond molecules to other molecules are called intermolecular forces.

3. Intermolecular Forces Because pure covalent compounds have low melting and boiling points, intermolecular forces must be very weak compared with the intramolecular forces (ie., it does not take very much energy to break the bonds that hold the molecules to each other).

4. Intermolecular Forces Intermolecular forces were studied extensively by Johannes van der Waals (1837 – 1923) and they are therefore often referred to as van der Waals forces.

5. Intermolecular Forces We will be looking at 4 Intermolecular forces: 1.Dipole – Dipole 2. Ion – Dipole 3.Dispersion Forces 4.Hydrogen Bonding

6. Intermolecular Forces Dipole – Dipole Forces  Dipole means polar molecule.  Dipoles will change their direction so that their oppositely charged ends are near to one another.  The electrostatic attraction between the oppositely charged ends of the polar molecules are called dipole-dipole forces.

7. Intermolecular Forces Dipole – Dipole Forces  As a result of these dipole – dipole forces of attraction, polar molecules will tend to attract one another more at room temperature than similarly sized non-polar molecules would.  The strength of these dipole-dipole forces and the amount of energy required to separate them is indicated by the temperature at which the substance changes phase – its melting point and its boiling point.

8. Intermolecular Forces Ion – Dipole Forces  The force of attraction between an ion and a dipole (polar molecule).  NaCl disolves in water because the polar charges on the water molecules are strong enough to overcome the forces that bind the ions together.

9. Intermolecular Forces Ion – Dipole Forces  In aqueous solution, ionic solids dissolve at their negative ends and positive ends as the water molecules become oriented with the corresponding oppositely charged ions that make up the ionic compound, pulling them away from the solid into solution.

10. Intermolecular Forces Dispersion (London) Forces  The weakest intermolecular force between non polar molecules.  It is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles.

11. Intermolecular Forces Dispersion (London) Forces  The shared pairs of electrons in a covalent bond are constantly vibrating.  The bond vibrations, which are part of the normal condition of a non-polar molecule, cause momentary, uneven distribution of charge.

12. Intermolecular Forces Dispersion (London) Forces  In other words, a non-polar molecule becomes slightly polar for an instant, and continues to do this on a random but on-going basic. At the instant that one non-polar molecule is in a slightly polar condition, it is capable of inducing a dipole in a nearby molecule.

13. Intermolecular Forces Dispersion (London) Forces  An intermolecular force of attraction results.  The degree of force is related to the number of electrons as well as the size and shape of the molecule – larger molecules tend to display stronger London dispersion forces.

14. Intermolecular Forces Hydrogen Bonding  Is the attractive force between the hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different molecule.  The molecule involved with hydrogen bonding should have at least one lone pair. More lone pairs means stronger hydrogen bonds.

15. Intermolecular Forces Interesting facts… “Hydrogen Bonding”  If a compound experiences strong intermolecular forces such as hydrogen bonding, it will have a higher melting and boiling point. This is because hydrogen bonding helps to stick the molecules together, making it harder to pull them apart.

16. Intermolecular Forces Interesting facts… “Like Dissolves Like”  Ionic solutes dissolve in polar solvents (ex: NaCl and H 2 O)  Non polar solutes dissolve in non polar solvents (ex: solid I 2 and liquid Br 2 )