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Electron Configuration
Background:
In 1911, Rutherford proposed our present view of the nuclear atom. His model described the atom as
having a very small, dense nucleus containing most of the atom’s mass and all of the atom’s positive charge. Bohr
proposed that negatively charged electrons were distributed around the nucleus at great distances giving an atom a
volume which is only empty space.
Schrodinger developed a theory describing the most likely location of electrons around the atom. Using the
results of this theory, a method was developed for writing electron configurations. Whole numbers, 1,2,3, and so
on, are used to represent energy levels, and s, p, d, and g denote energy sublevels. Subscripts above the sublevel
letter indicate the number of electrons in that sublevel. For example, 2p4 indicates 4 electrons in the p-sublevel of
the second energy level.
Steps for writing the electron configuration:
1. Start with the hydrogen atom on the periodic table and continue across and down until you get to the
atom of choice.
2. Write down the energy level (row number, for the s & p sublevels, row number – 1, for the d-sublevel,
and row number – 2, for the f-sublevel)
3. Write the correct sublevel (s, p, d, or f).
4. Use a superscript to write the correct number of electrons in each sublevel.
The maximum number of electrons in each sublevel is indicated below:
s-sublevel
2 electrons
p-sublevel
6 electrons
d-sublevel
10 electrons
f-sublevel
14 electrons
Example: Oxygen = 1s22s22p4 (which means – 2 electrons in the s-sublevel of the first energy level, 2
electrons in the s-sublevel of the second energy level, and 4 electrons in the p-sublevel of the second energy level).
Procedure: Write the electron configuration of the following atoms:
1.
Boron =
2.
Magnesium =
3.
Nickel =
4.
Sulfur =
5.
Fluorine =
6.
Silver =
7.
Mercury =
8.
Iron =
9.
Sodium =
10. Lead =
Steps for writing orbital filling diagrams:
1. Once you have completed writing the electron configuration for an atom, you need only to write the
orbital filling diagram for those electrons in the highest energy level (valence electrons) and any in
incomplete sublevels.
2. Dashes are used to represent orbitals of sublevels.
The table below shows how many orbitals (dashes) are used for each sublevel:
S – sublevel
1 orbital
P – sublevel
3 orbitals
D – sublevel
5 orbitals
F – sublevel
7 orbitals
Note: The maximum number of electrons in each orbital is two. If you multiply the number of
orbitals by 2 electrons/orbital, you will get the maximum number of electrons in each sublevel.
3.
4.
When adding electrons to orbitals, remember to follow HUND’S RULE and place electrons in empty
orbitals before beginning to pair them up.
Electrons are represented with arrows either face up (clockwise spin) or face down (counter-clockwise
spin).
Example: For oxygen, the electron configuration is 1s2
____
2s2
____
2p4
____ ____ ____
Procedure: Draw orbital filling diagrams for each of the atoms for which you have already written the electron
configuration.
1.
Boron
2.
Magnesium
3.
Nickel
4.
Sulfur
5.
Fluorine
6.
Silver
7.
Mercury
8.
Iron
9.
Sodium
10. Lead
Steps for writing drawing planetary models:
1. Draw the nucleus, including the number of protons and approximate number of neutrons.
2. Referring to the electron configuration, draw the same number of rings around the nucleus as the
number of occupied energy levels.
3. Using the superscripts, determine the number of electrons on each ring (each energy level) and place
this number below the ring.
4. As a check step, add up the number of electrons at the bottom of each ring. This number should equal
the atomic number of the atom.
Example: For oxygen, the planetary model is shown below:
1s22s22p4 -------Procedure: Draw the planetary model for each of the atoms for which you have already written electron
configurations.
Planetary models
Electron-Dot
1.
Boron
2.
Magnesium
3.
Nickel
4.
Sulfur
5.
Fluorine
6.
Silver
7.
Mercury
8.
Iron
9.
Sodium
10. Lead
Steps for writing Electron-Dot Diagrams:
1. Write the symbol of the atom.
2. Determine the number of valence electrons (electrons in the highest energy levels) in the atom.
Note: Valence electrons can only come from s or p sublevels!
3. Use X’s or DOT’s to represent the electrons and place the electrons around the atom in the sequence
shown below.
Example: Oxygen – has 6 valence electrons (2s22p4)
Procedure. Draw Electron-Dot diagrams for each of the atoms above.