Liquids and Solids

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AP Chemistry Chapter 10

Intramolecular force: forces within a molecule
Intermolecular force: forces between molecules

Intermolecular forces area very similar for liquids and solids.

  1. Properties of Molecular Substances
    1. Trends in melting point, boiling point
      1. Usually increase with molar mass.

        F2 <  Cl2 <  Br2 <  I2

        CH4 <  C2H6 <  C3H8 <  C4H10

      2. Slightly higher for polar molecules than for nonpolar molecules of comparable molar mass.
      3. Unusually
    2. Types of intermolecular forces
      1. Dipole forces: electrical attractive forces between + end of polar molecule and - end of adjacent polar molecule.

        dipole force

      2. Hydrogen Bond - unusually strong dipole forces. Hydrogen atom is very small and differs greatly in electronegativity from Fluorine, Oxygen, and Nitrogen. (easily remembered as "NOF")

        Note that water has many unusual properties in addition to high boiling point. Open structure of ice, as a result of hydrogen bonding, accounts for ice's low density.

        Click here to see a rendering of ice's structure.

      3. London Dispersion Forces - temporary dipoles formed in nonpolar molecules.
        • Strength depends upon how readily electrons are "polarized"
        • Increases with molecular size or molar mass.
  2. Types of Solids
    1. Ionic (NaCl, KNO3, etc.)
      High melting point (strong attractive forces between + and - ions). Nonconducting as solids, but conduct when melted. Often soluble in polar solvents such as water.
    2. Molecular (Cl2, CH4, S8, etc.)
      Low melting point (weak attractive forces between molecules). Nonconducting (no charged particles). Seldom water soluble.
    3. Atomic
      • Network covalent (C, SiC, SiO2, etc.)
        Conduct electricity and heat through movement of electrons. Cover wide range of melting points. insoluble in water.
      • Metallic
        metal atoms, conduct with delocalized electrons throughout the substance.
        • substitutional (brass = copper and zinc): atoms of comparable size in an alloy.
        • interstitial (steel = iron and carbon): atoms of different size in an alloy.
  3. Unit Cells in Solids
    - Unit Cell is the simplest unit which, repeated over and over again in three dimensions, generates the solid lattice.
    - Diffraction: the scattering of light from a regular array of points or lines, where the spacings between the points or lines are related to the wavelength of the light (X-rays for lattice structures).

    Bragg Equation: n lambda = 2 d sin theta
    n lambda = xy + yz (if the waves are in phase (n, integer))

    n = the order of diffraction
    lambda = wavelength of the incident energy
    d = the spacing between planes of atoms
    sin theta = the sine of the angle of reflection of the beam

    1 angstrom (Å) = 1 x 10-8 cm = 10 nm

    1. simple cubic - unit cell consists of eight atoms at the corners of a cube.

      2 r = s

      r = atomic radius
      s = length of cube edge

    2. face-centered cubic - Atoms at corners of cube and in center of each face. Atoms at corners of cube do not touch each other; instead, atoms touch along face diagonal.

      4 r = s square root of 2

      Typical of many metals and LiCl (Cl- ions form face-centered cubic lattice with Li+ ions fitting between them.)

    3. body centered cubic - Atoms at corners of cube and one atom in the center of the cube. Contact is along a body diagonal, which has a length of s square root of 3.

      4 r = s square root of 3

      In CsCl, Cs+ ion is at center of cube, touching 8 Cl- ions at corners.

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