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unit 3 atomic structure electron configurations

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					ELECTRON CONFIGURATIONS
Atomic Models



  Dalton’s Model
                    Thomson’s Model      Rutherford’s Model
                    (Plum-pudding)

 Rutherford’s atomic model lacked detail on how
 electrons occupied the space around nucleus
   Why don’t electrons crash into nucleus?
 Chemical Behavior of elements
• Light

• At high temperatures or voltages, elements in the
  gaseous state emit light of different colors.

 • When the light is passed through
 a prism or diffraction grating a
 line spectrum results.
 Bohrs Atomic model




 Bohr model : After Rutherford's discovery, Bohr proposed
 that electrons travel in definite orbits around the nucleus.
 (planetary model)



 •Explained hydrogen’s spectrum but not
 for other elements.
Each element has its own
unique set of spectral emission
lines that distinguish it from
other elements.




  Line spectrum of hydrogen. Each line corresponds
  to the wavelength of the energy emitted when the
  electron of a hydrogen atom, which has absorbed
  energy falls back to a lower principal energy level.
                   Modern View
 Movement of electrons is not
  completely understood
 The atom is mostly empty space
 Two regions
   Nucleus
     protons and neutrons

   Electron cloud
     region where you might

     find an electron
                               Quantum mechanical model
                               Modern atomic theory described the
                               electronic structure of the atom as the
                               probability of finding electrons within
                               certain regions of space.
• Instead of being located in orbits, the
  electrons are located in orbitals.
• An orbital is a region around the nucleus
  where there is a high probability of
  finding an electron, can hold a
  maximum of 2 electrons.
Quantum Numbers
 Four Quantum Numbers:
   Specify the “address” of each electron in an atom

Principal Quantum Number ( n )

Angular Momentum Quantum number ( l )

Magnetic Quantum Number ( ml )

Spin Quantum Number ( ms )
1. Principal Quantum Number ( n )

   Indicates the number of the energy level

   As n increase, size of electron cloud
    increases.                                                                                    1s


   Energy increases as n increases. (electrons
    closer to nucleus have less energy)
                                                                                                  2s

   2n2 = maximum # of electrons possible in
    the energy level

       Ex. if n=1 (energy level 1) it can only have 2
        electrons
                                                                                              n    3s
                    Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
2. Angular Momentum Quantum Number ( l )
   Describes the sublevel within each energy level
   # of sublevels = value of principal quantum
    number of that level
       Ex. n=1, has 1 sublevel
            n=2, has 2 sublevels


   lowest sublevel :       s
   second sublevel : p
   third sublevel :      d
   fourth sublevel : f
 There is just one s sublevel , thus it has one orbital that
  can hold only 2 electrons.


 There are three p sublevels and thus it has        s
  three orbitals; each can hold 2 electrons.
 There are five d sublevels and thus it has
 five orbitals; each can hold 2 electrons.




 There are seven f sublevels and thus it has
 seven orbitals; each can hold 2 electrons.
    Too complicated to show with drawings
Classwork
 P 118 # 6 and p122 # 7,8
3. Magnetic Quantum Number ( ml )
   Specifies the exact orbital within each sublevel


4. Spin Quantum Number ( ms )
   An orbital can hold 2 electrons that spin in opposite
    directions.
   Indicated by arrows (in opposite direction):
 General Rules For Writing
 Electron Configurations
1. Pauli Exclusion Principle
   Each orbital can hold TWO electrons with opposite
    spins.




  In the following diagrams boxes represent orbitals.
 • Electrons are indicated by arrows: ↑ or ↓.
2. Aufbau Principle
   Electrons fill the
   lowest energy
   orbitals first.
   The number
   represents n, the
   principal quantum
   number
3. Hund’s Rule
   Within a sublevel, place one e- per orbital before pairing
    them.




        WRONG
                                             RIGHT
Notation
 Orbital Diagram



  O
8e-       1s                    2s   2p




Electron Configuration

               1s 2 2s2 2p4

               C. Johannesson
 Classwork p 128 #14
Electron Dot Diagrams
 The electrons in the outer energy level (greatest value
 of n ) called valence electrons are the most
 important electrons for chemical reactions.

 Lewis electron dot diagrams are used to represent
  these outer electrons around the symbol of an
  element.
 Examples
  Lithium      Electron configuration: 1s22s1
Select electrons that are in the outer energy level (the
  ones with the largest principal quantum number):
                    1s22s1          Largest principal
                                         quantum number is 2
                                         and there is 1 electron in
                                         this level
         Li             Valence electron
1. Symbol of element represents nucleus and all electrons except
those in outer level
2. Write the electron configuration of element to determine
valence electrons.
 3. Each side of symbol represents an orbital, draw dots to
 represent electrons in that orbital.
Oxygen: 1s22s2 2p4

Oxygen: 1s2 2s2 2p4

Oxygen: has 6 valence electrons (2 +4)




              O
Krypton: 1s22s2 2p6 3s23p6 4s2 3d10 4p6

 Krypton: 1s22s2 2p6 3s23p6 4s2 3d10 4p6


krypton: has 8 valence electrons (2 +6)




               Kr
   Classwork p 130 # 15 (Z= atomic number)

				
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