1 / 29

Intermolecular Forces

Intermolecular Forces. Mr. Nelson AP Chem. Intermolecular Forces. Forces that exist between molecules Also called Van der Waals forces Compare to intramolecular forces, which are the forces that hold atoms together

yaphet
Download Presentation

Intermolecular Forces

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  1. Intermolecular Forces Mr. Nelson AP Chem

  2. Intermolecular Forces • Forces that exist between molecules • Also called Van der Waals forces • Compare to intramolecular forces, which are the forces that hold atoms together • Holds individual molecules together, but does not determine macro properties (aka boiling point)

  3. Intermolecular Forces • Largely determines physical properties of solids and liquids • Intermolecular forces are weaker than intramolecular forces (chemical bonding) • Boiling points and melting points are an indication of the strength of intermolecular forces

  4. Ion-Dipole forces • Interaction between an ion and a partial charge in a polar molecule • Not technically a van der Waals force • Positive ions are attracted to negative end of a dipole • Force depends on dipole moment of polar molecule and charge of the ion • Example: NaCl in Water

  5. Dipole-dipole force • Attractive forces between polar molecules • Weaker than ion-dipole forces • Positive end of one molecule is near negative end of another • In liquids, molecules of equal size and mass have increasing intermolecular attractions with increasing polarity

  6. London Dispersion Forces • London proposed that the motion of electrons in an atom or molecule can create an instantaneous dipole moment • Example: If both electrons on an He atom are on the same side of a nucleus, you have an instantaneous dipole moment

  7. London Dispersion Forces • A temporary dipole on one atom can induce a similar dipole on an adjacent atom (called an induced dipole)

  8. London Dispersion Forces • Because of induced dipoles, molecules attract one another • Dispersion forces exist in all types of molecules (polar or non-polar, charged or not) • In many cases, dispersion forces can be stronger than dipole-dipole • Example: CH3F (B.P. -78.9 °C) vs. CCl4 (B.P. 76.5 °C)

  9. London Dispersion Forces • Dispersion forces increase with molar mass of the molecule • More electrons allows for greater chances of dispersion • When comparing molecules of similar weights and shapes, dipole-dipole forces tend to be the decisive factor • When comparing differing weights, dispersion forces is the decisive factor

  10. Intermolecular Forces Recap • Example • Of Br2, Ne, HCl, HBr, and N2, which has: • The largest intermolecular dispersion forces? • The largest dipole-dipole attractive forces?

  11. Hydrogen Bonding • A special type of dipole-dipole attraction that exists between hydrogen atoms in a polar bond and small electronegative ions • Can only form between hydrogen and N, O, or F

  12. Hydrogen Bonding • Water has a lower molar mass and a higher boiling point • Dispersion forces do not account for this

  13. Hydrogen Bonding • Explains density of ice • Solid is less dense than liquid (not common)

  14. Review of IM forces

  15. Phase Changes • Changes in state of matter • In general, each state of matter can change into either of the two other states • Phase changes require energy • When becoming a more disordered state, requires energy to overcome intermolecular forces that hold them together

  16. Phase Changes

  17. Phase Changes • Melting, vaporization, and sublimation are all endothermic processes • Freezing, condensation, and deposition are all exothermic processes

  18. Phase Changes • Fusion • Melting of a solid • Molar heat of fusion or enthalpy of fusion (∆Hfus) • ∆Hfusion= -∆Hsolidification • Vaporization • Vaporizing of a liquid • Molar heat of vaporization or enthalpy of vaporization (∆Hvap) • ∆Hvaporization= -∆Hcondensation

  19. Heating Curves • Graph of temperature of the system versus heat added to the system • Some segments of the graph are heating as single phase others are converting one phase to another • Heat change when temp. inc. is given by q=mC∆T. During phase change, it is q=n∆H(n is moles of substance)

  20. Heating Curve of Water

  21. Phase Change Problem • Freon-11 (CCl3F) has a normal boiling point of 23.8 °C. The specific heat of CCl3F (l) is 0.87 J/g K and CCl3F (g) is 0.59 J/g K. The heat of vaporization is 24.75 kJ/mole. Calculate the heat required to convert 10.0 g of Freon-11 from liquid at -50.0 °C to gas at 50.0 °C.

  22. Vapor Pressure • In a closed system, liquid will initially evaporate and then condense back into its liquid form • The pressure exerted by this liquid/gas equilibrium is called vaporpressure • When evaporation occurs in an open system, no equilibrium can be established

  23. Vapor Pressure • Substances with high vapor pressure evaporate more easily than those with a low vapor pressure • Volatile liquids (aka liquids that evaporate easily) have high vapor pressures • Water can vaporize at room temp. because molecules in liquid move at different speeds and some can overcome IM forces (think space shuttles/escape velocity)

  24. Vapor Pressure • As the number of gas molecules increases in a closed system, the probability that gas molecules will strike the liquid increases. • This allows the molecules to become “trapped” by the IM forces of the liquid • http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/vaporv3.swf

  25. Temperature & Vapor Pressure • Vapor pressure increases nonlinearly with temperature • More molecules have enough KE to escape into gas phase

  26. Vapor Pressure & Boiling Point • Boiling point is when the vapor pressure equals the external pressure acting on the surface • A “normal” boiling point is measured at 1 atm • A more volatile liquid has a lower boiling point and a higher vapor pressure (and vice versa)

  27. Phase Diagrams • A graph that displays the conditions under which an equilibrium exists between different states of matter • Allows prediction of a substance’s state of matter at a given temperature and pressure

  28. Phase Diagrams • Critical Temperature (Tc) is the highest temperature a substance can exist as a liquid • Meaning above Tc molecular motion breaks IM attraction (Tc measures strength of IM forces) • Critical Pressure (Pc) is pressure required to bring about liquefaction at the critical temperature

  29. Phase Diagrams • Triple Point is the place at which all three phases exist at equilibrium

More Related