Lecture 1: Introductory Concepts

Read sections 1.3, 1.6, 1.7, 1.8, and 1.9 from your textbook, "Inorganic Chemistry" by C. E. Housecropft and A. G. Sharpe. Information in this section is a review of atomic structure, electron orbitals, and the periodic table from general chemistry. In the web page below, I've summarized some of the key points of your reading. Please go to the online quiz page and complete Survey #1.

Energy

  1. Energy and Mass Relationship
      Remember Einstein? Mass and energy are related through the equation:
      E = mc2

      where E is energy, m is mass, and c is the speed of light. Energy can take the form of heat (q), light (), work (w), or potential energy. One form of potential energy is the energy in chemical bonds.

  2. Units and Conversion


    We typically use joules or kilocalories (1000 calories) to discuss energy differences in chemistry. The conversion factors among different units of energy is shown at the right.

    In this class, always use scientific notation and give numerical answers to 2 significant figures.
    Significant figures are numbers that are NOT used to hold a place value. All of the numbers below have 2 significant figures:
      1500
      27
      4.6
      0.0071
      0.00011

    We can write all of these in scientific notation with a number, a decimal point, a second number, and a power of 10.
      1.5 x 103
      2.7 x 101
      4.6 7.1 x 10-3
      1.1 x 10-4

    Round numbers up (5 and above) or down to go to 2 significant figures. For example:
      36579 = 3.7 x 104

Atomic Structure

  1. nucleus
    • atomic mass: Count the number of neutrons + protons for the atomic mass. We write the atomic mass as a superscript in front of the atomic symbol.The most common isotope of carbon has a mass of 12, 12C.

    • atomic number: Count the number of protons in the nucleus for the atomic number. The number of electrons in a neutral atom is equal to the number of protons in the nucleus, or the atomic number. The atomic number of carbon is 6.

    • isotopes
      These are atoms of the same element that differ in the number of neutrons. Different isotopes have different natural abundance. The isotope of hydrogen with zero neutrons is the most abundant at xxx %. Deuterium and tritium with 1 and 2 neutrons, respectively, are present xxx % and xxx % natural abundance. Below are a couple of examples familiar to you.
        1H, 2H, 3H
        12C, 13C, 14C

    • nuclear spin
      The nucleus of an atom is always postively charged. A spinning charged body will generate a magnetic field. Some nuclei, such as 1H, 13C, or 31P, behave as if they have a spin. In an applied magnetic field, these nuclei generate small magnetic fields opposed to the applied field. This is the basis of 1H NMR and 13C NMR spectroscopy that you covered in organic chemistry. We'll talk about other nuclei with "spin" and multinuclear NMR spectroscopy when we discuss those elements.

    • region of stability
      Some nuclei are unstable and spontaneously break apart (radioactive decay). The energy per nuclear particle that binds these particles together decreases after mass number 120 and very heavy elements typically have short half-lives. We'll return to radioactive decay when we discuss the properties of these heavy elements.

  2. electrons
      Each electron in an atom is described by a unique quantum number.
      1. n, principle quantum number, value= 1, 2, 3...
      2. l, orbital quantum number, value= 0, 1, 2...(n-1)
      3. m, magnetic quantum number, integral values between -l and l
      4. s, spin quantum number, value= -1/2 or 1/2

      For l= 0, there is only one kind of orbital (m= 0), the s orbital. For l= 1, p, there are 3 kinds of orbitals (m= -1, 0, 1) that we know as the px, py, and pz. For l= 2, there are 5 kinds of d orbitals (m= -2, -2, 0, 1, 2). Two electrons at most can occupy each orbital with a spin quantum number of 1/2 or -1/2. The shapes of the simplest s, p, and d orbitals are shown below.



      The orbitals differ in the number of nodes. A nodal plane is a plane where electron density is exactly zero. Look at the sine wave at right. Think of this as a 2 dimensional representation of a p orbital. The part of the line above the base line encloses the mathematically positive part of the electron density function. The curve below the baseline encloses the mathematically negative part of the electron density function. This orbital has a nodal plane, a plane where the electron density is zero, at the nucleus.

      The 1s orbital has 0 radial nodes, 2p orbitals each have 1 node, 3d orbitals have 2 nodes, etc.

      When there are 2 electrons in an orbital, their spins (1/2, -1/2) cancel. Only when there are unpaired electrons in orbitals does the atom or molecule have a net electron spin.

      Oxidation = loss of one or more electrons


      Reduction = gain of one or more electrons



  3. Orbital Energies
    There is a different ordering of orbital energies for neutral atoms and for those same atoms in ions or molecules where there can be a partial charge on the atom.

    1. Neutral, isolated atoms
      1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 5d

    2. Ions and compounds
      1s < 2s < 2p < 3s < 3p < 3d < 4s < 4p < 4d < 5s < 5p < 5d


  4. Absortion Spectroscopy: measures the absorbance of energy by atoms or molecules corresponding to differences in energy levels.

      NMR spectroscopy: nuclear spin states

      EPR spectroscopy: electron spin states

      UV-visible spectroscopy: valence electron energy levels

      X-ray photoelectron spectroscopy: excitation of core electrons

      IR spectroscopy: vibrational energy levels

      microwave spectroscopy: rotational energy levels


  5. Periodic Table
    The periodic table is a device that helps us classify elements according to their electronic configuration. Many of the properties of elements and their reactivity relate to their position in the periodic table.



    1. Each row in the periodic table corresponds to one principle quantum number. The row beginning with H has n=1; the row beginning with Li has n=2; etc.

    2. Counting along a row from the left give the number of valence electrons. Titaniuum is 4 elements from the left and has 4 valence electrons. It is in a row corresponding to n=4 so its electronic configuration as a neutral atom is [Ar]4s24d2.

    3. If you ignore the noble gases (column including He, Ne, Ar...), electronegativity increases as you go from left to right across any row. It also increase as you go from bottom to top in any column. Electrons in orbitals of more electronegative elements are more stabilized. The opposite of electronegative is electropositive.

    4. Because they have the same valence shell electronic configuration, adjacent elements in any column make similar compounds. For example: H2O and H2S, NH3 and PH3.

    5. You should memorize the symbol and atomic number for the elements H through Ne.