Intermolecular Forces, Liquids, and Solids

1. Intermolecular Forces, Liquids, and Solids By: Ms. Buroker

2. 11.1: A Molecular Comparison of Liquids and Solids There is a fundamental difference between the three states of matter and that is due to the space between the particles.

3. Review of Gases Compressible - gases occupy ALLthe available volume: There is a large space between the molecules. The volume occupied by the gas is much larger than the space occupied by the molecules themselves. B) Intermolecular forces are quite WEAKat high temperature and low pressure. Under these conditions the molecules are far apart and act independently of one another. They are moving fast and the molecular size and _________________ can be ignored.

4. States of Matter Because in the solid and liquid states particles are closer together, we refer to them as condensed phases. The state a substance is in at a particular temperature and pressure depends on two things: 1.) The kinetic energy of the particles and 2.) The strength of attractions between the particles

5. 11.2: Intermolecular Forces The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together. Intramolecular

6. 11.2: Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

7. Intermolecular Forces There are different types of attractions between molecules/ions: Ion-Ion; Ion-Dipole; Dipole-Dipole; London Forces, etc. • Electrostatic attractions are determined by Coulomb’s Law: F = kq1q2/r2 • Attractive forces depend upon three factors: 1.) The distance between the ion and the dipole. *The closer the ion and dipole, the stronger the attraction 2.) The charge on the ion. *The higher the ion charge, the stronger the attraction. 3.) The magnitude of the dipole. *The greater the magnitude of the dipole, the stronger the attraction

8. Effects of intermolecular forces 1.) Affects solvation energy (enthalpy of hydration): 2.) Dissolving substances in one another * In general: Like substances dissolve like substances  polar in polar (ionic); nonpolar in nonpolar, etc.) 3.) Permanent dipole-attractions: DHvap increases with polarity and molar mass

9. “Like Dissolves Like”

11. Ion- Dipole Forces Ion-dipoleinteractions are an important force in solutions of ions. The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

12. Dipole- Dipole Interactions Molecules that have permanent dipoles are attracted to each other. * The positive end of one is attracted to the negative end of the other and vice versa. * These forces are only important when the molecules are close to each other.

13. Dipole- Dipole Forces

14. Dipole- Dipole Interactions In terms of polarity … would the more or less polar molecule have the higher boiling point?

15. London Dispersion Forces 1.) Produce “momentary dipoles” that hold the molecules together 2.) Caused by the temporary shifting of the electron clouds through interactions between Molecules 3.) For an homologous series (CH4, C2H6, C3H8, etc.) polarizability increases with molar mass (greater electron cloud shifting)

16. London Dispersion 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 happen that they occasionally wind up on the same side of the atom.

17. London Dispersion Forces At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

18. London Dispersion Forces Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1. London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

19. London Dispersion Forces *These forces are present in all molecules, whether they are polar or nonpolar. * The tendency of an electron cloud to distort in this way is called polarizability.

20. Factors Affecting London Dispersion Forces The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane). * This is due to the increased surface area in n-pentane.

21. Factors Affecting London Dispersion Forces * The strength of dispersion forces tends to increase with increased molecular weight. * Larger atoms have larger electron clouds, which are easier to polarize.

22. Which Have a Greater Effect: Dipole- Dipole Interactions or Dispersion Forces? 1.) If two molecules are of comparable size and shape, dipole-dipole interactions will likely be the dominating force. 2.) If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

23. * The nonpolarseries (SnH4 to CH4) follow the expected trend. • * The polar series • follows the trend • from H2Te through • H2S, but water is • quite an anomaly.

24. Hydrogen Bonding The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong. * We call these interactions hydrogen bonds.

25. Hydrogen Bonding Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

26. Just to Review …

27. Intermolecular Forces AffectMany Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution.

28. Viscosity Resistance of a liquid to flow is called viscosity. It is related to the ease with which molecules can move past each other. Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

29. Surface Tension Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

30. The "skin" on the surface of a liquid caused by the increase in downward force on surface molecules from by the molecules that exist below the surface. (i.e. there is no counter force from above). By definition, surface tension (γ), is the amount of work required to extend a liquid surface. Units are usually expressed in J/m2. Surface tension decreases with an increase in temperature. Lower Energy Higher Energy

31. Phase Changes

32. Energy Changes Associatedwith Changes of State Heat of Fusion: Energy required to change a solid at its melting point to a liquid. Heat of Vaporization: Energy required to change a liquid at its boiling point to a gas.

33. Energy Changes Associatedwith Changes of State * The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other. * The temperature of the substance does not rise during the phase change.

34. How do we calculate heats of vaporization and fusion? DH = (specific heat) x ( grams of substance) x DT

35. Vapor Pressure At any temperature, some molecules in a liquid have enough energy to escape. • As the temperature rises, the fraction of molecules that have enough energy to escape increases.

36. Vapor Pressure As more molecules escape the liquid, the pressure they exert increases. The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.

37. The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure. The normal boiling point is the temperature at which its vapor pressure is 760 torr.

38. Critical Temperature= the highest temperature at which a substance can exist as a liquid Critical Pressure= the pressure required to bring about the liquefaction