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Electrons in Atoms

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					                         Chapter 5: Electrons in Atoms
5.1 Light and Quantized Energy
     Electromagnetic radiation is a form of energy that exhibits wavelike behavior




          o Wavelength



          o Frequency



          o Energy



      Waves didn’t explain all of light’s behaviors
      The quantum concept
          o Matter can gain or lose energy only small, specific amounts called
             QUANTA
      Wave-particle duality
         o PHOTONS are tiny bundles of energy



           o Radiation is now defined as a stream of photons that move in a wavelike
             pattern

5.2 Quantum Theory and the Atom
     All of an atom’s properties revolves around its electrons

      GROUND STATE is the lowest allowable energy state for an electron
         o If an electron gains energy, it will jump up to an EXCITED STATE
                It has to gain a specific quantum of energy



                     It will be only be excited for a moment
                     It will fall back to the ground state, releasing the same amount of
                      energy it initially gained




           o Many photons have the same wavelength as visible light
                When an electron “falls” the energy released (in the form of a
                   photon) often can be seen as visible light

      An ATOMIC ORBITAL is a 3-D region around the nucleus that describes the
       probable location of an electron
          o PRINCIPLE ENERGY LEVEL is how close an electron is located to the
              nucleus
                   n=1 (close), n=4 (far away)
                   Also defines how much energy the electron has
                   n=1 (little energy), n=4 (lots of energy)
                     Each level has SUBLEVELS; Each sublevel has ORBITALS
                         o 4 different sublevels (s, p, d, f), each with a different
                             number of orbitals.
                         o s = 1 orbital
                         o p = 3 orbitals
                         o d = 5 orbitals
                         o f = 7 orbitals
                     Each ORBITAL can have TWO ELECTRONS at most

                     n=1




                     n=2




                     n=3




                     n=4




5.3 Electron Configurations

      Electron configuration is the arrangement of the electrons in the atom
          o Usually in the ground state, which is the most stable

      Aufbau Principle: Each electron occupies the LOWEST ENERGY ORBITAL
       available
          o Aufbau diagram
   All orbitals related to an energy sublevel are of equal energy
        o i.e., all three 2p orbitals are of equal energy

   The energy sublevels within a principle energy level have different energies
       o i.e., the 2s orbital has less energy than the three 2p orbitals

   The sequence of energy sublevels within a principle energy level (in order of
    increasing energy) is s … p … d … f

   Orbitals in one energy level can overlap the orbitals in another energy level
       o i.e., 4s orbital has less energy than the five 3d orbitals

   Pauli exclusion principle: A maximum of two electrons may occupy a single
    atomic orbital
       o Only if the electrons have opposite spins

   Hund’s rule: Single electrons with the same spin must occupy each equal-energy
    orbital BEFORE additional electrons with opposite spins can occupy the same
    orbital
   Practice Problems p. 139 18 (a-c with boxes; e-f without boxes)




   Helpful Hint for the Representative Groups 1-8
       o Groups 1 and 2 have their last electrons in the s orbital. The principle
          energy level is related to the period number of the element.
               Na = 3s1           K = 4 s1
                           2
               Be = 2s            Sr = 5s2
       o Groups 3-8 have their last electrons in the p orbital. Principle energy level
          is also related to the period number of the element.
               B = 2p1            In = 5p1
                           2
               Si = 3p            Pb = 6p2
                          3
               N = 2p
               S = 3p4
               I = 5p5
               Ar = 3p6
   Noble Gas Configuration
       o This is a shorthand way to show electron configurations
       o Find the first NOBLE GAS that has an atomic number lower than your
          given element
       o Put that symbol in brackets: [He]
       o Complete the rest of the configuration for all of the electrons after the
          noble gas: Magnesium – [Ne] 3s2
       o Phosphorus


       o Practice Problem p. 139 18 (in noble gas configuration), 21, 22




   VALENCE ELECTRONS
      o Remember the electrons determine the chemical properties of an element
      o Valence electrons are the OUTERMOST electrons (highest principle
        energy level)


       o These are the electrons that participate in BONDING
       o OCTET RULE




       o ELECTRON DOT (LEWIS DOT) STRUCTURES


       o Practice Problems p.141 23

				
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