Electron Configurations

1
Electron Configurations

• Of the three major subatomic particles, the
  electron plays the most significant role in
  determining the physical and chemical properties
  of an element. The arrangement of elements in the
  periodic table depends on these properties.
• Thus, there should be some relationship between
  the electron configurations of the elements and
  their placement in the table.

2
Electron Configurations

• The orbital names s, p, d, and f stand for names
  given to groups of lines in the spectra of the
  alkali metals. These line groups are called
  sharp, principal, diffuse, and fundamental.

3
Electron Configurations
4
Electron Configurations

• The electron configuration of an atom denotes the
  distribution of electrons among available shells.
  
• The standard notation lists the subshell symbols,
  one after another.
• The number of electrons contained in each
  subshell is stated explicitly.
• For example, the electron configuration of
  beryllium, with an atomic (and electron) number
  of 4, is 1s22s2 or He2s2.

5
Electron Configurations

• C 1s2 2s2 2p2 or He 2s2
  2p2
• Ne1s2 2s2 2p6 or He 2s2 2p6
• S 1s2 2s2 2p6 3s2 3p4 or Ne3s2 3p4

6
Aufbau Principle

• Electrons fill orbitals starting at the lowest
  available energy states before filling higher
  states (e.g. 1s before 2s).
• The number of electrons that can occupy each
  orbital is limited by the Pauli Exclusion
  Principle (each orbital can hold two electrons
  with opposite spins).

7
The rules of the Aufbau Principle are

•  The Lazy Tenant Rule
• 1.Electrons are placed in the lowest
  energetically available subshell. 
• 2. An orbital can hold at most 2 electrons. 
• 3.If two or more energetically equivalent
  orbitals are available (e.g., p, d etc.) then
  electrons should be spread out before they are
  paired up (Hund's rule).

8
Hund's Rule

• The Empty Bus Seat Rule
• If multiple orbitals of the same energy are
  available, Hund's Rule says that unoccupied
  orbitals will be filled before occupied orbitals
  are reused (by electrons having different spins).

WRONG
RIGHT
9
Increasing energy
10
The Noble Gases

• These are the elements in which the outermost s
  and p subshells are filled. The noble gases have
  full outer shells notice that these elements
  have filled outermost s and p sublevels
• Helium 1s2 He
• Neon 1s22s22p6 Ne
• Argon 1s22s22p63s23p6 Ar
• Krypton 1s22s22p63s23p63d104s24p6 Kr

11
Representative Elements

• In these elements, the outermost s or p sublevel
  is only partially filled. There are three groups
  of representative elements
• Group 1 alkali metals
• Group 2 alkaline earth metals
• Group 7 halogens

12
Representative Elements

• For any representative element, the group number
  equals the number of electrons in the outermost
  energy level (valence electrons)
• Potassium 1s22s22p63s23p64s1
• Carbon, silicon, and germanium, in Group 4, have
  four electrons in the outermost energy level
• Carbon 1s22s22p2
• Silicon 1s22s22p63s23p2
• Germanium 1s22s22p63s23p63d104s24p2

13
Transition Metals

• These are metallic elements in which the
  outermost s sublevel and nearby d sublevel
  contain electrons. The transition elements are
  characterized by addition of electrons to the d
  orbitals

14
Inner Transition Metals

• These are the metallic elements in which the
  outermost s sublevel and nearby f sublevel
  generally contain electrons

15
Write the electronic configurations for the
following elements

• O
• Na
• Ar
• Fe
• Ca
• Ce

1s22s22p4 or He2s22p4 1s22s22p63s1 or
Ne3s1 1s22s22p63s23p6 or Ne3s23p6 1s22s22p6
3s23p63d64s2 or Ar3d64s2 1s22s22p63s23p64s2
or Ar4s2 1s22s22p63s23p63d104s24p64d105s2
5p64f15d16s2 or Xe4f15d16s2
16
Half-Full and Full Subshells

• full subshell fully-filled shells are lower in
  energy than partially-filled shells (i.e. Noble
  Gases)
• half-filled subshells lower in energy than
  partially-filled subshells
• Cu exception Ar 4s13d10 rather than Ar
  4s23d9
• Cr exception Ar4s13d5 rather than Ar4s23d4