Atomic Models _In chronological order of discovery__

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Atomic Models _In chronological order of discovery__ Powered By Docstoc
					Scott Anderson, James Baek
Chapter 13
Atomic Models (In chronological order of discovery):

Nothing – By John Dalton
    Dalton created Atomic Theory
    Saw atoms as solid invisible masses

Plum Pudding – By J.J Thomson
    Thomson discovered electrons
    Model included:
         o Negatively Charged Electrons in a lump of Positively Charged Material
         o Nothing about:
                 # of Protons and Electrons
                 Arrangement of Protons and Electrons

Nuclear Atom – By Earnest Rutherford
    Rutherford discovered Nucleus
    Model includes:
          o Dense, Positively Charged Nucleus
          o Electrons surround the Dense Nucleus
          o Empty Space as the rest of the atom

Planetary Model – By Niels Bohr
    Theorized electrons in Orbit around nucleus, like planets, to prevent from falling
       into nucleus
    Electrons have Fixed Energy Levels – An energy level is a region around the
       nucleus where the electron is likely to be
    Quantum = amount of energy needed to move energy levels
Scott Anderson, James Baek
Chapter 13
Quantum Mechanical Model – Erwin Schrödinger
    Quantum mechanical model based on studies of light
    A Mathematical Equation describing the Location and Energy of electrons in
    Estimates Probability of finding an electron in a position – Represented by a
      fuzzy cloud
    Designates energy levels with Principle Quantum Numbers (n=1, 2, 3…)
         o Higher principle quantum numbers = larger distance from nucleus
         o Within each energy level, there are Sublevels –
                  n = # of Sublevels (Maximum number of electrons that can occupy
                   an energy level = 2n2)
                  Sublevels called Atomic Orbitals
                        s – Spherical – up to 2 electrons in “s” orbitals
                        p – Dumb-bell shaped – up to 6 electrons in “p” orbitals
                        d – Clover shaped – up to 10 electrons in each “d” orbital
                        f – Complex and hard to describe – up to 14 electrons in
                           each “f” orbital
Scott Anderson, James Baek
Chapter 13
Light and Spectra:
    Light consists of electromagnetic waves
    Light Terms:
           o Photons – Light quanta. Each corresponds to a particular energy change
                    Red/Yellow < Blue/Purple – See Page 373 in book for a chart
           o Photoelectric Effect – Electrons eject from metals when light is shone on
              the metal. Needs a particular frequency.
           o Frequency – Number of wave cycles to pass a point in a unit of time
                    Measured in hertz (Hz)
                    Inversely proportional to wavelength, as frequency increases,
                     wavelength decreases
                    Frequency X Wavelength = The speed of light
           o Atomic Emission Spectrum – Light emitted by an element through a prism
           o Planck’s Constant – Amount of radiant energy absorbed/emitted by
              something is proportional to the frequency of the radiation. (You can’t
              jump energy levels unless you have enough energy available)
    Bohr’s Model explained atomic spectra emissions of Hydrogen only. It didn’t
       explain how atoms bonded with each other to form molecules. Instead, quantum
       mechanics was developed to help explain atoms. (379 in book)
           o Quantum Mechanics
                    Describes motions of subatomic particles like electrons and
                     protons as wavelike. They gain and lose energy in packages.

Electron Configurations:
       - The way electrons are arranged around the nuclei of atoms
       - Can be expressed as Aufbau Diagrams or written in a short-hand form

Making Them:
               1. Aufbau Principle: Electrons enter orbitals of lowest energy first
                        Things to Remember:
                               D is of a higher energy than the S after
   Aufbau                             o Ex. 1s22s22p63s23p64s23d104p2
   Diagram                     F is of a higher energy level than the S 2 levels before
                                      o Ex.
               2. Pauli Exclusion Principle: An atomic orbital can contain up to two
                  electrons of opposite spins
               3. Hund’s Rule: When filling atomic orbitals, an entire sublevel must be
                  filled with single electrons with parallel spins before being paired with
                  a second of opposite spin.
                        “Like filling college dorms.”

          Remember:
Scott Anderson, James Baek
Chapter 13

             o   Filled electron sublevels are more stable than partially filled sublevels
                 Ex. Cu 1s22s22p63s23p64s13d10
             o   Half Filled sublevels are less stable than filled, but are more stable
                 than others.
                  Ex. Cr 1s22s22p63s23p64s13d5

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