Characteristics of Electrons
Extremely small mass
Located outside the nucleus
Moving at extremely high speeds in a sphere
Have specific energy levels
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Energy of Electrons
When atoms are heated, bright lines appear
called line spectra
Electrons in atoms arranged in discrete levels.
An electron absorbs energy to “jump” to a
higher energy level.
When an electron falls to a lower energy level,
energy is emitted.
Bohr Model
• First model of the electron structure
• Gives levels where an electron is most likely
to be found
• Incorrect today, but a key in understanding
the atom
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Bohr Model
• e- exist only in orbits with specific amounts of energy called energy levels
• Therefore…
– e- can only gain or lose certain amounts of energy
– only certain photons are produced
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Bohr’s Model
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23
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Energy of photon depends on the difference in energy levels
Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom
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Other Elements
Each element has a unique bright-line emission spectrum.
“Atomic Fingerprint”
Helium
Bohr’s calculations only worked for hydrogen!
The Quantum Mechanical Model
• Describes the arrangement and space
occupied by electrons in atoms
• This model is based on Quantum Theory,
which says that matter also has properties
associated with waves.
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Quantum Theory
• State that it is impossible to know an electron’s exact position and momentum at the same time (Uncertainty of Principle).
• So the scientist had to develop the concept of orbitals (sometimes called electron clouds), volumes of space in which an e- is likely present (certainty was replaced with probability).
Quantum Mechanical Model
Describes the arrangement of electrons in
atoms in terms of:
Main or principal energy levels (n)
Energy subshells
Orbitals (space occupied within the atom)
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Quantum Numbers
The principal quantum number, n
• describes the average distance of the orbital from the nucleus and the energy of the electron in an atom.
• (n=1) is the energy level nearest to the nucleus and the succeeding energy level is located farther from the nucleus.
• Electrons in energy levels at increasing distances from the nucleus have increasingly higher energies.
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Maximum No. of Electrons / Energy LevelPrincipal Energy
Level, n
Allowed per Energy Level = 2n2
1 2 x (1)2 = 2
2 2 x (2)2 = 8
3 2 x (3)2 = 18
4 2 x (4)2 = 32
5 2 x (5)2 = 50
6 2 x (6)2 = 72
Quantum Numbers
• The azimuthal quantum number, l sometimes called the Orbital Angular Momentum
Quantum Number
designates the overall shape of the orbital within a shell
affects orbital energies (bigger l = higher energy)
all electrons in an atom with the same value of l are said to
belong to the same subshell
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Subshells - Orbitals that have the same value of
n but different values of l
n = 4
n = 3
n = 2
n = 1
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
Quantum Numbers
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• The azimuthal quantum number, l
Letter Designation of the Subshells
Value of l Letter Subshell Name
# of electrons
that can exist
0 s sharp 2
1 p principal 6
2 d diffuse 10
3 f fundamental 14
Quantum Numbers
• The magnetic quantum number, ml
– determines the orientation of orbitals within a
subshell
– does not affect orbital energy (except in magnetic
fields!)
– only integer values between -l and +l are allowed
– the number of m values within a subshell is the
number of orbitals within a subshell
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Quantum Numbers
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• The magnetic quantum number, ml
The number of possible ml values determines the number of orbitals in a subshell.
l possible values of ml Number of orbital in this subshell
0 0 1
1 -1, 0, +1 3
2 -2, -1, 0, +1, +2 5
3 -3, -2, -1, 0, +1, +2, +3 7
Quantum Numbers
• The magnetic quantum number, ml
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Orbitals combine to form a spherical shape.
2s
2pz2py
2px
Quantum Numbers
• The spin quantum number, ms
several experimental observations can be
explained by treating the electron as though it were
spinning
spin makes the electron behave like a tiny magnet
spin can be clockwise or counterclockwise
spin quantum number can have values of
+1/2 or -1/2
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Electronic Configuration
• The statement of how many electrons an atom
has in each of its subshells.
• a list showing how many electrons are in each
orbital or subshell in an atom or ion
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Aufbau Principle
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– Electrons fill the
lowest energy orbital
first.
– “Lazy Tenant Rule”
Methods of Illustrating Electronic Configuration
• s p d f Notation
the distribution of electrons using this method
indicates the energy level and sublevels that
are filled.
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Methods of Illustrating Electronic Configuration
Energy sequence
1s 2s 3s 4s 5s 6s 7s
2p 3p 4p 5p 6p 7p
3d 4d 5d 6d 7d
4f 5f 6f 7f
Note:
s ≤ 2
p ≤ 6
d ≤ 10
f ≤ 14
A way to remember the filling order of subshells, write subshells
designations and follow the diagonal arrows starting from left.
Methods of Illustrating Electronic Configuration
• s p d f Notation
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1s2
main shell
subshell
number of electrons
Methods of Illustrating Electronic Configuration
• s p d f Notation
Example:
1.) 9F 1s2 2s2 2p5
2.) 13Al 1s2 2s2 2p6 3s2 3p1
3.) 2He 1s2
4.) 12Mg 1s2 2s2 2p6 3s2
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Pauli Exclusion Principle
• Limits the number of electrons in any orbital to
not more than two
• Each orbital can hold TWO electrons with
opposite spins.
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Hund’s Rule
• Priciple of Maximum Multiplicity
• Within a sublevel, place one e- per orbital before pairing them.
• “Empty Bus Seat Rule”
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RIGHTWRONG
Methods of Illustrating Electronic Configuration
• Orbital diagram
using arrows to represent electrons and boxes
for orbitals. (Pauli exclusion principle and
Hund’s Rule should be followed)
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Methods of Illustrating Electronic Configuration
• Orbital diagram
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↑↓ ↑↓ ↑↓ ↑ ↑
1s 2s 2p
Methods of Illustrating Electronic Configuration
Orbital diagram
1. write the electron configuration in subshell notation
2. draw a box for each orbital.
• Remember that s, p, d, and f subshells contain 1, 3, 5, and 7 degenerate orbitals, respectively.
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Methods of Illustrating Electronic Configuration
Orbital diagram
• Remember that an orbital can hold 0, 1, or 2 electrons only, and if there are two electrons in the orbital, they must have opposite (paired) spins (Pauli principle )
3. within a subshell (depicted as a group of boxes), spread the electrons out and line up their spins as much as possible (Hund's rule )
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Methods of Illustrating Electronic Configuration
• Orbital diagram
Example:
1.) 9F
2.) 13Al
3.) 2He
4.) 12Mg
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↑↓ ↑↓↑↓ ↑↓ ↑↓ ↑↓ ↑
↑↓ ↑↓↑↓ ↑↓ ↑↓ ↑↓
↑↓ ↑↓↑↓ ↑↓ ↑
↑↓
Learning Check
Using the periodic table, write the complete
electronic configuration using electron notation
for each:
A. Cl
B. Sr
C. I
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Solution
Using the periodic table, write the complete electronic configuration using electron notation for each:
A. Cl
1s2 2s2 2p6 3s2 3p5
B. Sr
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
C. I
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5
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Paramagnetic and Diamagnetic
Paramagnetic
- When there is unpaired electron
- configurations with unpaired electrons are attracted to magnetic fields (paramagnetism )
Diamagnetic
- When all electrons are paired
- configurations with only paired electrons are weakly repelled by magnetic fields (diamagnetism )
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Some other important terms
Differentiating electron
- the last entering electron in the electron configuration
Valence Shell
- the outermost shell
Valence Electron
-all the electrons in the outermost shell or valence shell
Core and Valence Electron
• chemistry involves mostly the shell with the highest
value of principal quantum number , n, called the
valence shell
• the noble gas core under the valence shell is
chemically inert
• simplify the notation for electron configurations by
replacing the core with a noble gas symbol in square
brackets:
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Example of Electron Configuration written with the
Core and Valence Electron
Atom Full configuration Core valence configuration
Full configuration using core/valence
notation
O 1s2 2s2 2p4 He 2s2 2p4 [He] 2s2 2p4
Cl 1s2 2s2 2p6 3s2 3p5 Ne 3s2 3p5 [Ne] 3s2 3p5
Al 1s2 2s2 2p6 3s2 3p1 Ne 3s2 3p1 [Ne] 3s2 3p1
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Psuedocore Electrons
• electrons in d and f subshells outside the noble
gas core are called pseudocore electrons
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Examples of electron configurations containing
pseudocore electrons
Atom Core Psuedocore Valence Full Configuration
Fe Ar 3d6 4s2 [Ar] 3d6 4s2
Sn Kr 4d10 5s2 5p2 [Kr] 4d10 5s2 5p2
Hg Xe 4f14 5d10 6s2 [Xe] 4f14 5d10 6s2
Pu Rn 5f6 7s2 [Rn] 5f6 7s2
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