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Chem 1211
Class 14
Atomic Structure
Chapter 6
© 2009 Brooks/Cole - Cengage
Arrangement of
Electrons in Atoms
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)
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QUANTUM NUMBERS
There is a hierarchy in quantum numbers:
A few different l can exist for the same n and a
few different ml can exist for the same l
• n determines the energy of the orbital
• for hydrogen atom the energy of the orbital
does not depend on l and ml
• l determines shape of the orbital
• ml determines position of the orbital relative
to others with the same l
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Azimuthal Quantum
Number
• l determines the shape of the orbital
• for given n, l can be any integer from 0 to (n-1)
• every l has its proper (letter) name:
orbital with l = 0 called
s-orbital
orbital with l = 1 called
p-orbital
orbital with l = 2 called
d-orbital
orbital with l = 3 called
f-orbital
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Magnetic Quantum
Number
• ml determines position of the orbital relative
to others with the same l and how many
orbitals with the same l can exist.
• ml is an integer from - l to l, so for given l,
there are (2l +1) orbitals with different ml
l =0
l =1
l =2
l =3
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s-orbital
p-orbital
d-orbital
f-orbital
How many?
1
3
5
7
5
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Types of
Atomic
Orbitals
2px
n, principal
l, azimuthal
ml, magnetic
See Active Figure 6.14
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Arrangement of
Electrons in Atoms:
restriction on number of
residents!
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Each orbital can be occupied by no more than 2
electrons!
Because:
No two electrons in atom can have the same set
of all quantum numbers (Pauli’s Exclusion
Principle)
And:
Wolfgang Pauli
(1900-1958)
There is a 4th quantum number, the electron
spin quantum number, ms.
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Electron
Spin
Quantum
Number,
ms
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• Can be proven experimentally (Stern, Gerlach,
1926) that electron has a magnetic moment (acts
as a small magnet).
• It referred to as “spin.”
• Spin is quantized: it can be only along magnetic
field or against magnetic field.
• Two spin directions are given by
ms = +1/2 and -1/2.
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Electron
Spin
Quantum
Number,
ms
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• If two electrons in atom have the opposite
spins, the total spin is zero (spins are “paired”)
• Electron spin is responsible for magnetic
properties of the materials
• Materials containing unpaired spins are either
paramagnetic or ferromagnetic
• Materials containing only paired spins are
diamagnetic
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Electron Spin and Magnetism
•Diamagnetic: NOT attracted to a magnetic
field (slightly repelled)
•Paramagnetic: substance is attracted to a
magnetic field.
•Ferromagnetic: substance is strongly
attracted to a magnetic field.
•Substances with unpaired electrons are
paramagnetic or ferromagnetic.
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Measuring Paramagnetism
Paramagnetic: substance is attracted to a
magnetic field. Substance has unpaired electrons.
Diamagnetic: NOT attracted to a magnetic field
See Active Figure 6.18
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magnetic field
ferromagnetic
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diamagnetic
paramagnetic
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QUANTUM NUMBERS
Now there are four!
n → shell
1, 2, 3, 4, ...
l → subshell
0, 1, 2, ... n - 1
ml → orbital
- l ... 0 ... + l
ms → electron spin
+1/2 and -1/2
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Magnetic Resonance Imaging
(MRI)
Protons and neutrons are also small magnets,
means, they also have spin;
Proton is the nucleus of hydrogen, energy
transitions of its spin in strong magnetic field can
be measured by Nuclear Magnetic Resonance
(NMR);
Concentration of the protons (i.e. water) in
different parts of the sample can be measured
by NMR - that is called Magnetic Resonance
Imaging (MRI)
It is powerful diagnostic tool, because living
organism mostly consists of water!
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Magnetic Resonance Imaging
(MRI)
X-Ray
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MRI
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Chem 1211
Atomic Electron Configurations
and Periodic Table
Chapter 7
© 2009 Brooks/Cole - Cengage
Arrangement of
Electrons in Atoms
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)
© 2009 Brooks/Cole - Cengage
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Arrangement of
Electrons in Atoms
Each orbital can be assigned no
more than 2 electrons!
This is tied to the existence of a
4th quantum number, the
electron spin quantum
number, ms and…
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Pauli Exclusion Principle
No two electrons in the
same atom can have
the same set of 4
quantum numbers.
That is, each electron has a
unique address.
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When n = 1, then l = 0
this shell has a single orbital (1s) to which
two e- can be assigned.
first shell
Electrons in Atoms
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2s orbital
2e-
three 2p orbitals
6e-
TOTAL =
8e-
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second shell
When n = 2, then l = 0, 1
Electrons in Atoms
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2e6e10e18e-
third shell
When n = 3, then l = 0, 1, 2
3s orbital
three 3p orbitals
five 3d orbitals
TOTAL =
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Electrons in Atoms
When n = 4, then l = 0, 1, 2, 3
4s orbital
three 4p orbitals
five 4d orbitals
seven 4f orbitals
TOTAL =
And many more!
2e6e10e14e32e-
In general, electron capacity of n-th shell is 2n2
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ATOMIC ELECTRON CONFIGURATIONS23
AND PERIODICITY
Length of period corresponds to shell capacity only for the 1st
two periods.
WHY?
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Assigning Electrons to Atoms
• Electrons generally assigned to orbitals of
successively higher energy.
• For H atoms, E = - C(1/n2). E depends only
on n.
• For many-electron atoms, energy depends
on both n and l.
•
See Active Figure 7.1 and Figure 7.2
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Assigning Electrons to Subshells
• In H atom all subshells
of same n have same
energy.
• In many-electron atom:
a) subshells increase in
energy as value of n + l
increases.
b) for subshells of same n
+ l, subshell with lower
n is lower in energy.
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Aufbau scheme: electrons occupy the position with the 26
lowest possible energy
Electron
Filling
Order
See Figure 7.2
Why is this order?
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Effective Nuclear Charge, Z*
• Z* is the nuclear charge experienced by
the outermost electrons. See Figure 7.3
• Explains why E(2s) < E(2p)
• Z* increases across a period owing to
incomplete shielding by inner electrons.
• Estimate Z* = [ Z - (no. inner electrons) ]
• Charge felt by 2s e- in Li
Z* = 3 - 2 = 1
• Be
Z* = 4 - 2 = 2
• B
Z* = 5 - 2 = 3
and so on!
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Effective
Nuclear
Charge
See Figure 7.3
Z* is the nuclear
charge experienced by
the outermost
electrons.
Electron
cloud for 1s
electrons
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Effective
Nuclear
Charge
probability
distribution for
2p electron
Shielding by inner electrons: s < p < d
Effective charge for s > p > d
Correspondingly, energy: s < p < d
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See Figure 7.3
Z* is the nuclear
charge experienced by
the outermost
electrons.
Aufbau scheme: electrons occupy the position with the 30
lowest possible energy
Electron
Filling
Order
See Figure 7.2
Now we
understand
the order!
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Writing Atomic Electron
Configurations
Two ways of
writing configs.
One is called
the spdf
spdf notation
for H, atomic number = 1
notation.
1
1s
value of n
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no. of
electrons
value of l
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Writing Atomic Electron
Configurations
Two ways of
writing
configs. Other
is called the
orbital box
notation.
ORBITAL BOX NOTATION
for He, atomic number = 2
Arrows
2
depict
electron
spin
1s
1s
One electron has n = 1, l = 0, ml = 0, ms = + 1/2
Other electron has n = 1, l = 0, ml = 0, ms = - 1/2
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See “Toolbox” in ChemNow for Electron Configuration tool.
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characteristic
elements
Electron Configurations
and the Periodic Table
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See Active Figure 7.4
Lithium
Group 1A
Atomic number = 3
1s22s1 → 3 total electrons
3p
3s
2p
2s
1s
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Beryllium
3p
3s
2p
2s
1s
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Group 2A
Atomic number = 4
1s22s2 → 4 total
electrons
Boron
3p
3s
2p
2s
1s
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Group 3A
Atomic number = 5
1s2 2s2 2p1 →
5 total electrons
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Carbon
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Group 4A
Atomic number = 6
1s2 2s2 2p2 →
6 total electrons
3p
3s
2p
2s
1s
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Here we see for the first time
HUND’S RULE. When placing
electrons in a set of orbitals having
the same energy, we place them
singly as long as possible
(maximal spin configuration has
the lowest energy).
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Nitrogen
3p
3s
2p
2s
1s
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Group 5A
Atomic number = 7
1s2 2s2 2p3 →
7 total electrons
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Oxygen
3p
3s
2p
2s
1s
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Group 6A
Atomic number = 8
1s2 2s2 2p4 →
8 total electrons
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Fluorine
3p
3s
2p
2s
1s
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Group 7A
Atomic number = 9
1s2 2s2 2p5 →
9 total electrons
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Neon
Group 8A
Atomic number = 10
1s2 2s2 2p6 →
10 total electrons
3p
3s
2p
2s
1s
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Note that we have
reached the end of
the 2nd period, and
the 2nd shell is full!