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Chemistry 100 – Chapter 11 Intermolecular Forces.

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Presentation on theme: "Chemistry 100 – Chapter 11 Intermolecular Forces."— Presentation transcript:


2 Chemistry 100 – Chapter 11 Intermolecular Forces

3 Intramolecular and Intermolecular Forces  We have just been discussing the covalent bond - the force that holds atoms together making molecules.  We have also talked about the ionic bond. These are intramolecular forces.  There are also forces that cause molecules to attract each other. These are called intermolecular forces.

4 Intermolecular Forces  Salt (NaCl) is a solid because of the strong electrostatic attraction of the Na + and the Cl – ions: the ionic bond  Q: Why is Cl 2 a gas, Br 2 a liquid and I 2 a solid? A: Intermolecular forces  Q: Why is molasses “thick” while water has low viscosity? Same answer.  Q: Why is it possible to float a needle on water?

5 Intermolecular Forces (cont’d)  Q: Why is water a liquid but H 2 S is a gas at 25ºC?  Q: Why are real gases not ideal? A: van der Waals forces (same thing, different name)

6 States of Matter  Gas: No defined shape or volume, compressible, rapid diffusion, flows readily  Liquid: Takes shape of container, virtually incompressible, low diffusion, flows readily  Solid: Has its own shape and volume, virtually incompressible, extremely slow diffusion, does not flow  Liquids and solids are called condensed phases because particles are close together


8 Intermolecular Forces  Much weaker than chemical bonds. Covalent bonds 200 kJ/mol and more.  Intermolecular forces less than 50 kJ/mol  When a liquid vaporizes, the intermolecular forces must be overcome. But no covalent bonds are broken.

9 Ion - dipole forces  Dissolve an ionic compound in water  Charged ions interact with the dipole of water molecules

10 Dipole-dipole Forces  Interactions between the dipoles in a polar liquid leads to a net attraction

11 Dipole - Induced dipole force  In a mixture of two liquids were one is polar and the other is not, the dipole of the polar molecule can induce a dipole in the other.  The energy of this force depends on the polarizabilty of the non-polar molecule. Larger molecules are more polarizable

12 London Dispersion Forces  How do we account for the fact that non-polar gases can be liquefied and solidified?  Fritz London proposed instant dipoles. These are dipoles that result from the random movement of the electron cloud.

13 Dispersion forces work over very short distances

14 Polarizability  The ease with which a dipole can be induced depends on the polarizability of the molecule  Large molecules are more polarizable - easier to distort the electron cloud

15 Polarizability and Boiling Points Boiling points (K) of halogens F 2 Cl 2 Br 2 I 2 85.1238.6332.0457.6 Boiling points of Noble gases HeNeArKr Xe 4.6 27.387.5120.9 166.1

16 Dipole or Dispersion?  Dispersion forces operate between all molecules - polar and non-polar  Molecules with comparable molecular weights and shapes have approximately equal dispersion forces. Any difference is due to dipole-dipole attractions  When molecules differ widely in molecular weight, dispersion forces tend to be the decisive ones

17 Water And Ammonia Have Unusual Boiling Points!!

18 The Hydrogen Bomb (err – Bond)  A special dipole-dipole intermolecular attraction H in a polar bond (H-F, H-O or H-N) an unshared electron pair on a nearby electronegative ion or atom (generally F, O or N)  Hydrogen bonds (4 to 25 kJ/mol, or larger) are weaker than covalent bonds but stronger than most dipole-dipole or dispersion forces.

19 Molecules with H bonding  HF (can behave as if it were H 2 F 2 )  NH 3  H 2 O (ice is less dense than liquid at 0ºC)  alcohols (e.g. CH 3 OH)  amines (e.g CH 3 NH 2 )  carboxylic acids (e.g. CH 3 COOH)

20 Summary

21 Liquids - viscosity  Viscosity (resistance to flow)  Molecules that have strong intermolecular forces cannot move very easily - more viscous  Viscosity decreases at higher temperatures. The kinetic energy overcomes the intermolecular forces

22 Liquids - surface tension  Surface tension: the energy that must be expended to increase the surface of a liquid

23 Surface tension Liquids with strong intermolecular bonds have high surface tensions Water: strong H-bonds, so high surface tension, 7.29  10 –2 J/m 2 Mercury: atoms held by metallic bonds, so even higher surface tension, 4.6  10 –1 J/m 2

24 Wetting & capillary action  Water wets clean glass (spreads out) but beads on a waxy surface  Water climbs up a capillary tube  Mercury does not wet glass  Mercury level is depressed in capillary tube

25 To wet or not to wet?  Competition between two tendencies  Cohesive forces; intermolecular forces that bind similar molecules together. Keeps liquid as a bead  Adhesive forces: Intermolecular forces between molecules of a liquid and those of a surface. Makes liquid spread out

26 Are we all wet??  Water on clean glass adhesive forces > cohesive forces explains the upward meniscus of water  Water on polished table cohesive forces win because H 2 O molecules are not attracted to wax  Mercury on glass - cohesive forces win because Hg molecules are not attracted to glass  Explains depression in capillary tube and downward meniscus

27 Phase changes

28  Enthalpy of fusion or heat of fusion for water H fus = 6.01 kJ/mol.  Enthalpy of vaporization or heat of vaporization for water H vap = 40.67 kJ/mol. cooling effect of evaporation refrigeration (Not Freon-12) steam burn generally severe

29 Heating Curve

30 Supercooling, superheating  A liquid cooled below its freezing point is said to be supercooled requires very clean conditions the molecules are moving slowly but have not organized themselves into the solid form  A liquid heated above its boiling point is said to be superheated a danger with heating water in microwave oven

31 Critical T and P  A gas can be liquefied by cooling  A gas can be liquefied by increasing pressure but only if the temperature is below the compound’s critical temperature substancecritical temp K and C ammonia405.6 (133) carbon dioxide304 (31) argon150.9 (-122)  Critical pressure: pressure needed to liquefy gas at critical temperature  Substance at T c and P c - supercritical fluid

32 Vapour Pressure  Every liquid in a closed container gives off vapour until a certain pressure is reached - the liquid’s vapour pressure.  The vapour pressure of a liquid increases with increase in temperature.  We can explain these facts using the kinetic theory


34 Volatile  A liquid in an open container will evaporate  As vapour moves away, the liquid releases more molecules into the vapour phase to try to build up to the correct vapour pressure  Liquids with high vapour pressure evaporate more quickly - they are volatile

35 Boiling point: the temperature at which vp = 760 torr

36 Boiling  A liquid boils when the vp equals the atmospheric pressure  Normal boiling point temp -> when vp is 760 torr. We list normal bp values in textbooks  Actual boiling point -> then the liquid has a vp equal to the external atmospheric pressure  Water boils at temperature lower than 100ºC atop mountains - it never reaches 100ºC  Water boils at higher temp in pressure cooker

37 Phase diagrams  A graphical way to show the equilibria between different phases of a substance  Thing to look for: critical point triple point ( three phases) how bp (and mp) varies with pressure  Triple point is not pressure dependent - the vapour has to be at the critical pressure! useful for thermometer calibration



40 Structure of Solids  Crystalline: atoms, ions, or molecules are well ordered. Have a well-defined melting point. Often the solid has regular shapes.  Amorphous: no order to the particles. Examples are glass and rubber. Have no defined mp; they soften over a range of temperatures (important for glass blowing)

41 Unit cell: crystal lattice  In a brick wall there is a repeating pattern, as there is with most wallpaper  In a crystalline solid there is a repeating pattern - the unit cell. The unit cell repeats to make the crystal lattice

42 The seven unit cells

43 Three Cubic lattices

44 Close Packing

45 Another way of looking at it


47 Diamonds are a …..

48 Some ionic solids- lattice decided by size & charge

49 From sea to shining... array of metal ions in a sea of electrons  A sea of valence electrons  Electrons not tightly held but can move  Explains electrical conduction  Also explains optical properties - most metals “shine”

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