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INTERMOLECULAR FORCES LIQUIDS AND SOLIDS

INTERMOLECULAR FORCES LIQUIDS AND SOLIDS. Chapter 11. STATES OF MATTER. Weak attractive forces between molecules. Inter molecular forces stronger. Strong inter molecular forces.

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INTERMOLECULAR FORCES LIQUIDS AND SOLIDS

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  1. INTERMOLECULAR FORCESLIQUIDS AND SOLIDS Chapter 11

  2. STATES OF MATTER

  3. Weak attractive forces between molecules Intermolecular forces stronger Strong intermolecular forces The state of a substance depends largely on the balance between the kinetic energies of the particles and the interparticle energies of attraction.

  4. INTERMOLECULAR FORCES intramolecular force - covalent bond between atoms in a molecule (can also be ionic bond) weak 16 kJ/mol Strong 431 kJ/mol intermolecular force - attraction between molecules HCl boils at -85oC at 1 atm When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).  Many properties are governed by the strength of the intermolecular forces e.g. boiling point, melting point, vapor pressure, viscosity, etc.

  5. A liquid boils when: A solid melts when:

  6. 4 types of intermolecular forces: • Ion-Dipole Forces • Dipole-Dipole Forces • London Dispersion Forces • Hydrogen Bonding van der Waals forces electrostatic forces Can you arrange them in order of increasing strength? What is a dipole? - +

  7. Ion-Dipole Forces • Interaction between an ion and a dipole • Strongest of all intermolecular forces. • Ion-dipole interactions make it possible for ionic substances to • dissolve in polar solvents.

  8. Dipole-Dipole Forces • Interaction between two dipoles • Weaker than ion-dipole forces. • Polar molecules need to be close together. • Molecules in liquids are free to move • results in both attractive and repulsive forces.  overall and stronger attractive force

  9. If two molecules have about the same mass and size, then intermolecular attractions (dipole-dipole forces) increase with increasing polarity.

  10. London Dispersion Forces • Interaction between nonpolar atoms or molecules, but can also occur between polar molecules • Occur only when atoms or molecules are close together • Weakest of all intermolecular forces. While the electrons in the 1s orbital of helium would repel each other (and therefore tend to stay far away from each other), it does occasionally happen that they wind up on the same side of the atom. Instantaneous dipoles formed! These dipoles are temporary

  11. Polarisability:the ease with which the distortion of the charge distribution occurs. Greater polarisablility  stronger dispersion In general, larger molecules tend to have greater polarisability as they have more electrons which are further away from the nuclei.

  12. Exercise Explain why n-pentane has a higher boiling point than neopentane? Both have the molecular formula C5H12.

  13. Hydrogen Bonding Special case of dipole-dipole forces Boiling pints polar nonpolar

  14. Hydrogen bonds are dipole-dipole interactions between the H-atom in a polar bond (usually H-F, H-O or H-N) and an unshaired e- pair on a nearby small electronegative ion or atom (usually F, O, N) Hydrogen bond

  15. Why are we specifying F, O and N? Small and electronegative! Hydrogen only has 1e- +ve nucleus is rather exposed. Hydrogen can approach the small electronegative atom closely and interact strongly.

  16. Exercise Why does ice float on water?

  17. O H3C– C–CH3 Exercise Which of the following molecules can hydrogen bond with itself? CH2F2 NH3 CH3OH

  18. SUMMARY

  19. PROPERTIES OF LIQUIDS Viscosity The resistance of a liquid to flow - related to the ease with which molecules can move past each other Viscosity increases with molecular weight and decreases with higher temperature. Which intermolecular forces also increase with molecular weight?

  20. Surface tension A measure of the inward forces that must be overcome in order to expand the surface area of a liquid Surface molecules experience a net inward force. Molecules at the surface can pack more closely together Molecules in the interior are attracted equally in all directions. Surface tension of water at 20oC: 0.0729 J/m2

  21. PHASE CHANGES Every phase change is accompanied by a change in energy of the system. H (change in enthalpy) Hvap Hcond Hsub Hdepos Hfus Hfreez (fusion)

  22. Hvap Hfus • Hvap> Hfus • In the transition from liquid to vapour phase, the molecules must essentially sever all intermolecular interactions. • In melting, many of these interactions remain. Hsub =

  23. Plot of temperature versus heat added is a heating curve e.g. Ice initially at -25oC is heated (constant P = 1 atm) While evaporating, heat added is used to break intermolecular forces Recall: Hvap> Hfus While melting, heat added is used to break intermolecular forces From the graph determine: Hfus Hvap HAB HCD HEF ?

  24. A water 50oC C B 0oC ice -30oC D n = (100 g)/(18.016 g/mol) n = 5.55 mol Exercise Calculate the enthalpy change, ΔH, when 100 g of water at 50 oC is cooled down to ice at -30 oC. Given: Specific heat capacities: Water = 4.18 J/g K, Ice = 2.09 J/g K Δ Hfus = 6.01 kJ/mol HAB HBC = Hfreez H = CmT HCD T = Tf - Ti HAB = (4.18 J/g K)(100 g)(0 - 50 K) = -20.9 kJ HBC = ΔHfreez = -ΔHfus = -(6.01 kJ/mol)(5.55 mol) = -33.4 kJ HCD = (2.09 J/g K)(100 g)(-30 - 0 K) = -6.27 kJ HAD = HAB + HBC + HCD = -60.6 kJ What is the enthalpy change when 100 g of ice at -30 oC is heated up to water at 50 oC?

  25. VAPOUR PRESSURE Closed container • Some of the molecules on the surface of a liquid have enough • energy to escape the attraction of the bulk liquid these • molecules move into the gas phase. • As the number of molecules in the gas phase increases, some of • the gas phase molecules strike the surface and return to the • liquid. • After some time the pressure of the gas will be constant at the • vapour pressure when liquid and vapour reach dynamic • equilibrium.

  26. As the temperature increases, the fraction of molecules that have enough energy to escape increases.

  27. The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure. Normal boiling point  boiling point at 1 atm. Why does water boil at a higher temperature at the coast than here in Jhb?

  28. PHASE DIAGRAMS Phase diagrams display the state of a substance under various pressure and temperature conditions. The solid lines show the conditions P,T conditions under which equilibrium exists between phases.

  29. Line AB: liquid-vapour interface - Each point along this line is the boiling point of the substance at that pressure. Critical point B: above this critical temperature and critical pressure the liquid and vapour are indistinguishable from each other. Line AD: solid -liquid interface - Each point along this line is the melting point of the substance at that pressure.

  30. Line AC: solid-vapour interface - Each point along this line is the sublimation point of the substance at that pressure. - Note: the substance cannot exist in the liquid state below A Triple point A: the temperature and pressure condition at which all three states are in equilibrium.

  31. Phase Diagram of Water Note the high critical temperature and critical pressure:  due to the strong van der Waals forces between water molecules. The slope of the solid–liquid line is negative.  The melting point decreases with increasing pressure. WHY?

  32. Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm. CO2sublimes at normal pressures.

  33. BONDING IN SOLIDS Solids can be: crystalline amorphous or Particles are in highly ordered arrangement No particular order in the arrangement of particles.

  34. The physical properties of crystalline solids (e.g. m.p., hardness) depend on the arrangement of particles and on attractive forces between particles. There are 4 types of crystalline solids: • Molecular • Covalent network • Ionic • Metallic

  35. Explain the m.p.’s and b.p.’s observed below: BenzeneToluenePhenol m.p./oC 5 -95 43 b.p./oC 80 111 182 Molecular Solids Atoms or molecules are held together by intermolecular forces  dipole-dipole forces, London dispersion forces, hydrogen bonds Because of these weak forces they are soft and have relatively low melting points (<200oC) Most substances that are gases or liquids at room temperature form molecular solids at low temperature.

  36. Covalent-Network Solids Atoms are held together in large networks or chains by covalent bonds. Covalent bonds much stronger than intermolecular forces  Harder solids and higher melting points than molecular solids. Diamond Graphite Which one is harder and has the higher m.p.? Explain.

  37. Ionic Solids Ions are held together by ionic bonds  strength of the ionic bond depends on the charges of the ions Ions pack themselves so as to maximize the attractions and minimize repulsions between the ions  Depends on relative size and charge of ions

  38. Metallic Solids Consist of entirely metal atoms. Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces. Bonding due to valence electrons delocalized throughout the solid. In general, the strength of bonding increases as the no. of electrons available for bonding increases The m.p. for sodium is 97.5oC and for chromium is 1890oC. Explain.

  39. SUMMARY

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