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Electron Structure 2: Energy Level Diagram of hydrogen and polyelectronic atom and
electronic configuration
IMPORTANT!!!!!
Definition: an ENERGY LEVEL is a specific amount of energy which an electron in an atom
can possess
The energy levels of hydrogen have the pattern shown below (“n” is the number of the
energy level.) The actual number of energy level is very large; only a few are shown
The pattern of lines in the spectrum reflects the energy level pattern. The observed
spectrum represents energy level differences occurring when an electron in higher
energy levels gives off energy and drops down to a lower level. The energy difference
between two particular energy levels is called the QUANTUM of energy associated with
the transition between the two levels.
Neil Bohr thought the electron orbited the nucleus along a specific path. This concept was
replaced by Quantum Mechanics (which you will learn in Grade 12 chemistry or First Year
University Chemistry). In Quantum Mechanics, different electrons, depending on their
energies simply occupy particular regions of space called “orbitals”.
Important
Definition: an ORBITAL is the actual region of space occupied by an electron in a
particular energy level. Each Orbital can hold 2 electrons.
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The Energy Level Diagram for Hydrogen
Experiment show that the lowest set of energy levels for hydrogen is arranged as follows.
Each dash represents he energy possessed by a particular orbital in the atom. The letter
s, p, d, and f refer to four different types of orbital’s which are discussed in class.
Definition:
A SHELL is the set of all orbital’s having the same n-value
For example, the 3rd shell consists of 3s, 3p, 3d, 3 f orbitals
A SUBSHELL is the set of orbital’s of the same type
For example, the set of five 3d- orbital’s in the 3rd shell is a subshell
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The energy level diagram for polyelectronic atoms
As the repulsion between electrons changes the energy of different types of orbitals,
the energy level diagram for hydrogen must be modified to describe any other atom.
Fortunately, the modified diagram can be used for ALL polyelectronic atoms (atoms
having more than one electron)
As can be seen on the above energy level diagram, all the orbital’s for a hydrogen atom
with a given value of n have the same energy (this is not true for atoms with more than
one electron).
Tue rules governing which types of orbital can occur for a given energy level and how many
orbital of a given type can exist are given below:
1) For a given value of “n”, n different types of orbital’s are possible
 For n=1: only the s-type is possible
 For n=2: only the s and p-type is possible
 For n=3: only the s, p, and d-type is possible
 For n=4: the s, p, d, and f-type is possible
2) An s-type subshell consists of ONE s- orbital
A p-type subshell consists of THREE p- orbital
A d-type subshell consists of FIVE d- orbital
An f-type subshell consists of SEVEN f- orbital
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The Energy Level Diagram for Polyelectronic (more than 1 electron) atoms
The energy level diagram for hydrogen must be modified to describe any other atom.
Fortunately, the modified disgram below can be used for all polyelectronic atoms (atoms
having more than one electrons)
The energy level order from low to high: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f,
5d, 6p, 7s, 5f, 6d, 7p
An easy way to remember it:
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ELECTRON CONFIGURATIONS
The addition of electrons to atomic orbitals follows three simple rules.
(a) The Aufbau ("Building Up") Principle: As the atomic number increases, electrons are added to the
available orbitals. To ensure the LOWEST POSSIBLE ENERGY for the atom, electrons are added to
the lowest–energy available orbitals.
(b) The Pauli Exclusion Principal: Each orbital can have a maximum of 2 electrons.
(c) Hund's Rule: Electrons placed in degenerate orbitals tend to remain unpaired.
Writing Electron Configurations for Neutral Atoms
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An ELECTRON CONFIGURATION lists the orbitals that contain electrons in an atom and how many
electrons are present in each orbital. In order to show where the electrons exist, the above three rules are
used for placing electrons in an atom. Refer to the energy level diagram on p. 9 to see how the “electron–
filling” proceeds.
Hydrogen has 1 electron, which goes in the lowest (1s) energy level [Aufbau Principle]. Therefore, H has
the ELECTRON CONFIGURATION
H (1s1) — pronounced “Hydrogen, one s one”.
Helium has 2 electrons. Since each orbital can contain up to 2 electrons [Pauli Exclusion Principle], both
of helium’s electrons go into the 1s orbital, and helium has the electron configuration
He (1s2) — pronounced “Helium, one s two”.
Two electrons fill the 1st electron shell completely. The 1st electron shell, having n = 1, consists of only
one orbital, the 1s.
Lithium has 3 electrons. The first 2 electrons fill the 1s orbital and the 3rd electron goes into the orbital
with the next–higher energy [according to the Aufbau Principle]: the 2s. The electron configuration of Li is
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Li (1s22s1).
Similarly, the addition of a 4th electron gives
Be (1s22s2).
A problem arises in filling up the 2p orbitals. The three 2p orbitals are designated as 2p x , 2py and 2pz
and can be filled in any order. Since all three p–orbitals have the same energy (they are DEGENERATE),
sequential addition of electrons DOES NOT completely fill one orbital, say 2px , before going on to 2py and
2pz. Hund's Rule for degenerate energy levels requires that one electron be placed in each of the 2p
orbitals before pairing up the electrons in the 2p orbitals.
B (1s2 2s2 2p1x )
or simply
B (1s22s22p1)

C (1s2 2s2 2p1x 2p1y )
or simply
C (1s22s22p2)

N(1s2 2s2 2p1x 2p1y 2p1z )
or simply
N (1s22s22p3)

O (1s2 2s2 2p2x 2p1y 2p1z )
or simply
O (1s22s22p4)

F (1s2 2s2 2p2x 2p2y 2p1z )
or simply
F (1s22s22p5)

Ne (1s2 2s2 2p2x 2p2y 2p2z )
or simply
Ne (1s22s22p6)
The constant
 reference to 2px , 2py and 2pz is quite cumbersome, so that the simpler method (shown
above, to the right) is used in which the TOTAL number of electrons in a given subshell is shown, rather
than specifying which individual orbitals in the subshell actually contain electrons.
At this point the second shell is completely filled (all orbitals having n = 2 are completely filled) and the
second shell is said to be CLOSED (as in “closed to further filling”).
The diagram below shows the manner in which the electron energy levels dictate the structure of the
periodic table.
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EXAMPLE: Arsenic, As, has 33 electrons and is situated halfway along the “4p” block of orbitals. Getting
to As on the table requires passing through all the intervening blocks of orbitals and filling
them from left to right in a given horizontal line before going down to the next line and
continuing.
As (1s22s22p63s23p64s23d104p3)
Technetium, Tc, has 43 electrons and is situated in the “4d” block of orbitals. Again, all the
intervening orbitals are filled in order to get to the 4d orbitals.
Tc (1s22s22p63s23p64s23d104p65s24d5)
EXERCISE:
14. Predict the electron configuration of the following.
(a) P
(e) Sr
(i) Ca
(m) Ga
(b) Ti
(f) Nb
(j) Kr
(n) Y
(c) Co
(g) Ge
(k) Cs
(o) Xe
(d) Br
(h) Cd
(l) Ar
(p) Rh
Core Notation
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An atom’s electrons are divided into two subsets: CORE electrons and OUTER electrons.
Definition: The CORE of an atom is the set of electrons having the configuration of the nearest inert gas
with an atomic number LESS than that of the atom being considering.
The OUTER electrons consist of all electrons outside the core.
Since core electrons normally do not take part in chemical reactions, core electrons are frequently not
explicitly included when writing the electron configuration of an atom.
CORE NOTATION is a way of writing the electron configuration in terms of the core and outer electrons.
The rules for writing an electron configuration using core notation are straight–forward.
EXAMPLE:
Write the core notation for aluminum atom.
• Locate the atom and note the noble gas at the end of the row above the element
Ne is the noble gas at the end of the row above Al.
• Start to write the electron configuration as usual, but replace the part of the electron
configuration corresponding to the configuration of the noble gas with the symbol for the
noble gas in square brackets: [...]. Follow the core symbol with the configuration of the
remaining electrons in the row containing the element.
Al (1s22s22p63s23p1) becomes
EXAMPLES:
Al ( [Ne] 3s23p1)
Full notation
Core notation
S (1s22s22p63s23p4)
S ( [Ne] 3s23p4)
Rb (1s22s22p63s23p64s23d104p65s1)
Rb ( [Kr] 5s1)
Kr (1s22s22p63s23p64s23d104p6)
Kr ( [Ar] 4s23d104p6)
Note: If you are given the electron configuration of an element in the 3rd or greater row, and are asked to
re–write the configuration in core notation, look backward from the end of the given electron
configuration until you find a "p6": this marks off the end of the core electrons.
EXERCISE:
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15. Re–write the electron configurations in exercise 14 using core notation.
Electron Configuration Exceptions
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There are two exceptions to the expected configurations of the elements up to Kr. The configurations
expected for Cr and Cu are
Cr ( [Ar] 4s23d4) ; d4 is one electron short of a half–filled subshell
Cu ( [Ar] 4s23d9) ; d9 is one electron short of a filled subshell
but experiments show that the configurations are
Cr ( [Ar] 4s13d5) ; 4s1 and 3d5 are two half–filled subshells
Cu ( [Ar] 4s13d10) ; 4s1 is a half–filled subshell, and 3d10 is a filled subshell
The following generalization is based on the behavior of Cr, Cu and a few other atoms and ions.
A filled or exactly half–filled subshell is especially stable and is favoured by a lowered energy.
From an energy point of view, the order of preference for the filling of orbitals is
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Writing Electron Configurations for Ions
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(a) NEGATIVE IONS To write the electron configuration of a NEGATIVE ION, add electrons to the last
unfilled subshell starting where the neutral atom left off.
Examples: O ( [He] 2s22p4) + 2 e–
Na ( [Ne] 3s1) + e–
(b) POSITIVE IONS
O2– ( [He] 2s22p6)
Na– ( [Ne] 3s2)
To write the electron configuration of a POSITIVE ION, two rules are used.
• Electrons in the outermost shell (largest n–value) are removed first.
• If there are electrons in both the s and p orbitals of the outermost shell, electrons in the p–
orbitals are removed first.
The idea is that the outermost electrons are most readily available and therefore are removed
preferentially. In addition, electrons in the highest–energy outermost orbital require the least
amount of energy to be completely removed from the atom so that p–electrons are removed
before s–electrons.
Examples: Sn ( [Kr] 5s24d105p2)
2 e– + Sn2+ ( [Kr] 5s24d10)
– 5p is removed before 5s
Sn ( [Kr] 5s24d105p2)
4 e– + Sn4+ ( [Kr] 4d10)
– 5p and 5s are removed before 4d
V ( [Ar] 4s23d3)
2 e– + V2+ ( [Ar] 3d3)
– 4s is removed before 3d (4s is OUTSIDE 3d)
EXERCISE:
16. Write the electron configuration of the following ions, using core notation.
(a) H–
(c) Br–
(e) Ti2+
(g) Mn2+
(i) Au3+
(k) Ru3+
(m) Cr2+
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(b) Sr2+
(d) N3+
(f) N2–
(h) Cu+
(j) Ge2+
(l) Sb3+
(n) Zn2+
Determining the Number of Valence Electrons
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Definition: VALENCE ELECTRONS are electrons that can take part in chemical reactions.
or
VALENCE ELECTRONS are electrons that are outside the core electrons.
Example: Al ( [Ne] 3s23p1) has 3 valence electrons: 3s23p1
The following electrons are NOT COUNTED AS VALENCE ELECTRONS.
• core electrons
• electrons in filled d and f subshells
EXAMPLE:
Ga ( [Ar] 4s23d104p1) has 3 valence electrons (omit 3d10)
Pb ( [Xe] 6s24f145d106p2) has 4 valence electrons (omit 4f14 and 5d10)
Xe ( [Kr] 5s24d105p6) has ZERO valence electrons (inert gas configuration)
EXERCISE:
17. How many valence electrons do each of the following contain?
(a) O
(c) Ag+
(e) Zn2+
(g) I5+
(i) Tc4+
(k) Cr
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(b) Cl–
(d) Nb3+
(f) Ge4+
(h) Xe2+
(j) Sb3+
(l) Cu2+