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Intermolecular forces

Intermolecular forces. Intermolecular forces: the forces between (among) individual particles (atoms, ions, molecules); they are weak relative to intramolecular forces (i.e. covalent and ionic bonds within a compound) Intermolecular forces are increasingly significant in this order:

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Intermolecular forces

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  1. Intermolecular forces • Intermolecular forces: the forces between (among) individual particles (atoms, ions, molecules); they are weak relative to intramolecular forces (i.e. covalent and ionic bonds within a compound) • Intermolecular forces are increasingly significant in this order: • GAS < LIQUID < SOLID

  2. Strength of Intermolecular forces Ionic compounds • Ion-ion interactions Covalent/molecular compounds • Dipole-dipole interaction • Hydrogen bonding • super strong type of dipole-dipole • occurs in molecules with O-H, N-H or F-H bonds • Dispersion (see text)

  3. Main Types of Intermolecular Attraction • Ion-ion interactions: the force of attraction between 2 oppositely charges ions, ionic bonding/attraction is quite strong, so ionic compounds have high melting points. • containing highly-charged ions have higher melting points than compounds containing univalent (1+ or 1-) ions. • Example: arrange the following ionic compounds in the expected order increasing melting and boiling pts: • NaF, CaO, CaF2 +2 -2 +1 -1 +2 -1 3 Absolute diff. = 2 4 NaF CaF2 CaO Increasing m.p./b.p.

  4. Main Types of Intermolecular Attraction Dipole-dipole interactions: occur between the + end of one polar molecular and the - end of another.

  5. Main Types of Intermolecular Attraction • Hydrogen bonding: occurs in polar molecules that contain hydrogen that is bonded to one of the very electronegative elements O, N, or F. • a + hydrogen atom is attracted to an unshared pair of electrons on an O, N, or F atom on an adjacent molecule:

  6. Liquids and their Properties • surface tension: a measure of the inward forces that must be overcome to expand the surface of a liquid; molecules on the surface are attracted only toward the interior molecules.

  7. example: Hgcohesive forces are stronger than adhesive forces example: H2Oadhesive forces are stronger than cohesive forces Liquids & Their Properties • cohesive forces: the forces that hold a liquid together. • adhesive forces: the forces between a liquid and another surface.

  8. Resonance • A molecule or polyatomic ion for which 2 or more dot formulas with the same arrangement of atoms can be drawn is said to exhibit RESONANCE.

  9. Resonance Example • CO32- • 3 resonance structures can be drawn for CO32- • the relationship among them is indicated by the double arrow. • the true structure is an average of the 3.

  10. Resonance Example • CO32- • 3 resonance structures can be drawn for CO32- • the relationship among them is indicated by the double arrow. • the true structure is an average of the 3.

  11. Resonance Example • CO32- • 3 resonance structures can be drawn for CO32- • the relationship among them is indicated by the double arrow. • the true structure is an average of the 3

  12. Resonance Structures • Another way to represent this is by delocalization of bonding electrons: • (the dashed lines indicate the 4 pairs of bonding electrons are equally distributed among 3 C-O bonds; unshared electron pairs are not shown) • See p. 256

  13. VSEPR valence shell electron pair repulsion

  14. Molecular Shape • Lewis structures (electron dot structures) show the structure of molecules…but only in 2 dimensions (flat). • BUT, molecules are 3 dimensional! • for example, CH4 is:

  15. Molecular Shape • Lewis structures (electron dot structures) show the structure of molecules…but only in 2 dimensions (flat). • BUT, molecules are 3 dimensional! • but in 3D it is: a tetrahedron! = coming out of page = going into page = flat on page

  16. Why do molecules take on 3D shapes instead of being flat? • Valence Shell Electron Pair Repulsion theory • “because electron pairs repel one another, molecules adjust their shapes so that the valence electron pairs are as far apart from another as possible.”

  17. Why do molecules take on 3D shapes instead of being flat? • Valence Shell Electron Pair Repulsion theory • Remember: both shared and unshared electron pairs will repel one another. Non-Bonding Pairs H—N — H — Bonding Pairs H

  18. 5 Basic Molecule Shapes • Linear • Example: CO2

  19. 5 Basic Molecule Shapes • Bent or angular • Example: H2O • Notice electron pair repulsion

  20. 5 Basic Molecule Shapes • tetrahedral • example: CH4

  21. 5 Basic Molecule Shapes • Pyramidal • Example: NH3 • (note: unshared pair of electron repels, but is not considered part of overall shape; no atom there to contribute to the shape)

  22. 5 Basic Molecule Shapes • Trigonal planar or planar triangular • Example: BF3

  23. Geometry and polarity • Three shapes will cancel out polarity. • Shape One: Linear

  24. Geometry and polarity • Three shapes will cancel out polarity. • Planar triangles 120º

  25. Geometry and polarity • Three shapes will cancel out polarity. • Tetrahedral

  26. Geometry and polarity • Others don’t cancel • Bent

  27. Geometry and polarity • Others don’t cancel • Trigonal Pyramidal

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