1. Intermolecular Forces Chapter 12 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. The Kinetic Molecular Theory  The Kinetic Molecular Theory explains the forces between molecules and the energy that they possess. This theory has 3 basic assumptions.  Matter is composed of small particles (molecules). The measure of space that the molecules occupy (volume) is derived from the space in between the molecules and not the space the molecules contain themselves.  The molecules are in constant motion. This motion is different for the 3 states of matter. Gas - The kinetic energy of the molecule is greater than the attractive force between them, thus they are much farther apart and move freely of each other. Lack strong forces of attraction between molecules. Liquid - Molecules will flow or glide over one another, but stay toward the bottom of the container. Motion is a bit more random than that of a solid. Significant forces of attraction. Solid - Molecules are held close to each other by their attractions of charge. They will bend and/or vibrate, but will stay in close proximity. Vibrate around fixed position. Significant forces of attraction.  When the molecules collide with each other, or with the walls of a container, there is no loss of energy. Gas Liquid Solid From: http://www.psinvention.com/ki netic.htm

3. A phase is a homogeneous part of the system in contact with other parts of the system but separated from them by a well-defined boundary. 2 Phases Solid phase - ice Liquid phase - water 12.1

4. Observable Properties  Takes shape and volume of its container  Low density  Very compressible  Takes shape of its container, has definite volume  High density  Difficult to compress  Has definite shape and volume  Density slightly higher than for liquid  Highly incompressible Gas Liquid Solid

5. Intermolecular Forces Intramolecular Forces  Forces of attraction between molecules  Much weaker than intramolecular forces  Boiling points and melting points of a substance generally increase with the strength of intermolecular forces  Hold atoms together in a molecule

6. Intermolecular Forces 12.2 Intermolecular vs Intramolecular 41 kJ to vaporize 1 mole of water (inter) 930 kJ to break all O-H bonds in 1 mole of water (intra) Generally, intermolecular forces are much weaker than intramolecular forces. “Measure” of intermolecular force boiling point melting point H vap H fus H sub

7. Types of Intermolecular Forces  Dipole-dipole  Ion-dipole  Dipole-induced dipole  Ion-induced dipole  Dispersion forces (London forces) (Induced dipole-induced dipole)

8. Intermolecular Forces Dipole-Dipole Forces - Attractive forces between polar molecules (molecules with dipole moments) Orientation of Polar Molecules in a Solid 12.2 The larger the dipole moment  greater the force

9. Intermolecular Forces Ion-Dipole Forces - Attractive forces between an ion and a polar molecule 12.2 Ion-Dipole Interaction Strength of interaction is dependent on: charge and size of the ion magnitude of dipole moment size of molecule

10. 12.2

11. Intermolecular Forces Dispersion Forces (London forces) (Induced dipole-induced dipole) - Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules 12.2 ion-induced dipole interaction - Attractive interaction between an ion and induced dipole dipole-induced dipole interaction - Attractive interaction between a polar molecule and induced dipole Attractive interaction in non-polar substances Also present in polar molecules

12. Induced dipole: dipole in a non-polar molecule or atom Caused by separation of + and – charges in an atom or non- polar molecule due to proximity of an ion or a polar molecule What affects if dipole moment is induced? Charge on ion Strength of dipole Polarizability of atom or molecule

13. Induced Dipoles Interacting With Each Other 12.2

14.  Only type of intermolecular forces present in non-polar molecules are dispersion forces  All molecules (both polar and non-polar) have dispersion forces **However, while polar molecules may have other types of intermolecular forces present, non-polar molecules have ONLY dispersion forces

15. Polarizability – ability to distort electron cloud (change shape of electron cloud) The more electrons in a molecule  the more polarizable it is (easier to distort electron cloud)  larger molar mass  stronger dispersion forces The more electrons in a molecule  the higher the melting point

16. S O O What type(s) of intermolecular forces exist between each of the following molecules? HBr HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules. CH 4 CH 4 is nonpolar: dispersion forces. SO 2 SO 2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO 2 molecules. 12.2

17. Types of Intermolecular Forces  Dipole-dipole Hydrogen bonding *** (see next slide) Strong type of dipole-dipole interaction  Dipole-induced dipole  Dispersion forces (Induced dipole-induced dipole) These types of intermolecular forces are referred to as van der Waals forces

18. Other types of Intermolecular Forces that are NOT van der Waals forces  Ion-dipole  Hydrogen bonding Placed in a separate category since only certain elements can take part in formation of hydrogen bonding

19. Hydrogen bonding  Special type of dipole-dipole interaction between hydrogen atom in a polar bond (N-H, O-H, or F-H) and an electronegative O, N, or F atom  Strong type of intermolecular attraction  However, hydrogen bond is still much weaker than covalent (intramolecular) bonds

20. 12.2 A H … B A H … A or A & B are N, O, or F Hydrogen bonding

21. Hydrogen Bond 12.2

22.  Boiling points of NH 3, H 2 O, HF are much higher than expected Why?  due to hydrogen bonds

23. Boiling points  Usually increase as molar mass increases, for similar compounds composed of elements in the same group on periodic table Why?  More electrons in a molecule  increase in dispersion forces  increase in boiling point

24. Figure 12.6 in textbook (see next slide)  Hydrogen compounds of Group 4A (increase molar mass  increase boiling point) CH 4 (lightest compound)  lowest boiling point SnH 4 (heaviest compound)  highest boiling point  Hydrogen compounds of Groups 5A, 6A, 7A (NH 3, H 2 O, HF) do NOT follow this trend In these groups, the lightest compound  heaviest boiling point  this is due to hydrogen bonding (NH 3, H 2 O, HF have stronger intermolecular attractions (hydrogen bonding) than other molecules in the same groups)

25. Decreasing molar mass Decreasing boiling point 12.2