Download Chapter Eight Electron Spin

Chapter Eight
Electron
Configurations and
Periodicity
Electron Spin
In Chapter 7, we saw that electron pairs residing
in the same orbital are required to have opposing
spins.
This causes electrons to behave like tiny bar
magnets. (See Figure 8.3)
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An electron configuration of an atom is a particular
distribution of electrons among available subshells.
The notation for a configuration lists the subshell symbols
sequentially with a superscript indicating the number of
electrons occupying that subshell.
For example, lithium (at. # 3) has two electrons in the 1s
subshell and one electron in the 2s sub shell 1s2 2s1.
An orbital diagram is used to show how the
orbitals of a subshell are occupied by electrons.
Each orbital is represented by a circle.
Each group of orbitals is labeled by its subshell
notation.
The Boron atom
1s
2s
2p
Electrons are represented by arrows: up for ms = +1/2 and
3
down for ms = -1/2
The Pauli exclusion principle, which summarizes
experimental observations, states that no two
electrons can have the same four quantum numbers.
- In other words, an orbital can hold at most two
electrons, and then only if the electrons have
opposite spins.
The maximum number of electrons and their orbital
diagrams are:
Maximum
Number of Number of
Sub shell
Orbitals
Electrons
s (l = 0)
1
2
p (l = 1)
3
6
d (l =2)
5
10
f (l =3)
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Aufbau Principle
Every atom has an infinite number of possible electron
configurations.
The configuration associated with the lowest energy
level of the atom is called the ground state.
Other configurations correspond to excited states.
Table 8.1 lists the ground state configurations of
atoms up to krypton.
The Aufbau principle is a scheme used to
reproduce the ground state electron configurations
of atoms by following the building up order.
You need to remember the order of filling the
orbitals (next slide)
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Order for Filling Atomic Subshells
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s,
4f, 5d, 6p, 7s, 5f
1s
2s
3s
4s
5s
6s
2p
3p
4p
5p
6p
3d
4d 4f
5d 5f
6d 6f
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Aufbau Principle
The building up order corresponds for the
most part to increasing energy of the
subshells.
By filling orbitals of the lowest energy first, you
usually get the lowest total energy ( ground
state ) of the atom.
Remember, the number of electrons in the
neutral atom equals the atomic number, Z.
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Table 8.1
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Question
Which ground-state electron configuration is
INCORRECT?
1) Cr: [Ar] 3d6
2) Ca: [Ar] 4s2
3) Na: 1s2 2s2 2p6 3s1
4) Zn: [Ar] 3d10 4s2
5) Kr: [Ar] 3d10 4s2 4p6
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Here are a few examples.
Using the abbreviation [He] for 1s2, the
configurations are:
Z=4 Beryllium 1s22s2 or [He]2s2
Z=3 Lithium
Z=5
Z=6
Z=7
Z=8
Z=9
Z=10
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
1s22s1 or [He]2s1
1s22s22p1
1s22s22p2
1s22s22p3
1s22s22p4
1s22s22p5
1s22s22p6
or
or
or
or
or
or
[He]2s22p1
[He]2s22p2
[He]2s22p3
[He]2s22p4
[He]2s22p5
[He]2s22p6
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With sodium (Z = 11), the 3s subshell begins to fill.
Z=11
Sodium
1s22s22p63s1
or
[Ne]3s1
Z=12
Magnesium
1s22s22p23s2
or
[Ne]3s2
Then the 3p subshell begins to fill.
Z=13
:
:
Z=18
Aluminum
1s22s22p63s23p1
or [Ne]3s23p1
Argon
1s22s22p63s23p6
or [Ne]3s23p6
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PROBLEM Give (1) the expected ground-state electron
configurations (full as well as abbrev.)
(2) the orbital -filling diagrams for;
a. P (Z=15)
b. Zn (Z=30)
c. Ca (Z=20)
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Note that elements within a given family have
similar configurations.
The Group IIA elements are sometimes called
the alkaline earth metals.
Beryllium 1s22s2
Magnesium 1s22s22p63s2
Calcium
1s22s22p63s23p64s2
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Configurations and the Periodic Table
Electrons that reside in the outermost shell of an
atom or in other words, those electrons outside the
noble gas core are called valence electrons.
These electrons are primarily involved in chemical
reactions.
Elements within a given group have the same
valence shell configuration.
This accounts for the similarity of the chemical
properties among groups of elements.
The total number of valence electrons for an atom
equals its group number.
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Another way of writing the electron
configurations is the Orbital Diagram
Consider carbon (Z = 6) with the ground
state configuration 1s22s22p2.
Hund s rule states that the lowest energy
arrangement (the ground state ) of
electrons in a subshell is obtained by
putting electrons into separate orbitals of
the subshell with the same spin before
pairing electrons.
Carbon's orbital
diagram is
1s
2s
2p
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Table 8.2
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Question?
Which of the following orbital filling diagrams
represent :
1) ground state or 2) excited state or 3) forbidden
for Boron (Z= 5)
1s
( )
(
2s
)
( )
( )
( )
(
( )
( )
2p
( ) ( ) ( )
(
)
) ( ) ( )
( ) ( ) ( )
(
) ( ) ( )
Magnetic Properties
Although an electron behaves like a tiny
magnet, two electrons that are opposite in
spin cancel each other. Only atoms with
unpaired electrons exhibit magnetic
susceptibility.
A paramagnetic substance is one that is weakly
attracted by a magnetic field, usually the result
of unpaired electrons.
A diamagnetic substance is not attracted by a
magnetic field generally because it has only
paired electrons.
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Question
The ground-state electron configuration of a Fe (s) is
1s2 2s2 2p6 3s2 3p6 4s23d6. Therefore, Fe is
1) diamagnetic.
2) paramagnetic with one unpaired electron.
3) paramagnetic with two unpaired electrons.
4) paramagnetic with four unpaired electrons.
5) paramagnetic with five unpaired electrons.
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The periodic law states that when the elements
are arranged by atomic number, their physical
and chemical properties vary periodically.
We will look at three periodic properties:
Atomic radius, Ionization energy, Electron affinity
A. Atomic Radius Trends
1) Within each period (horizontal row), the
atomic radius tends to decrease with increasing
atomic number (nuclear charge).
2) Within each group (vertical column), the
atomic radius tends to increase with the period
number.
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Figure 8.10:
Representation
of atomic radii
(covalent radii)
of the maingroup
elements
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Two factors determine the size of an
atom.
1) One factor is the principal quantum
number, n. The larger is n, the larger
the size of the orbital.
2) The other factor is the effective
nuclear charge, which is the positive
charge an electron experiences from the
nucleus minus any shielding effects
from intervening electrons.
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B. Ionization energy
The first ionization energy of an atom is
the minimal energy needed to remove
the highest energy (outermost) electron
from the neutral atom.
For a lithium atom, the first ionization
energy is illustrated by:
Li(1s 2 2s1 )
Li (1s 2 ) e
Ionization energy = 520 kJ/mol
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Ionization energy trends
There is a general trend that ionization energies
increase with atomic number within a given period.
This follows the trend in size, as it is more difficult
to remove an electron that is closer to the nucleus.
For the same reason, we find that ionization
energies, again following the trend in size,
decrease as we descend a column of elements.
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Periodic Properties
Ionization energy
The electrons of an atom can be
removed successively.
The energies required at each step are
known as the first ionization energy, the
second ionization energy, and so forth.
Table 8.3 lists the successive ionization
energies of the first ten elements.
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C. Electron Affinity
The electron affinity is the energy change for the
process of adding an electron to a neutral atom
in the gaseous state to form a negative ion.
For a chlorine atom, the first electron affinity is
illustrated by:
Cl([Ne]3s 2 3p 5 )
e
Cl ([Ne]3s 2 3p 6 )
Electron Affinity = -349 kJ/mol
The more negative the electron affinity, the more stable
the negative ion that is formed.
The general trend goes from lower left to upper right as
electron affinities become more negative.
Table 8.4 gives the electron affinities of the main-group
elements.
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Table 8.4
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The Main-Group Elements
The physical and chemical properties of the
main-group elements clearly display
periodic behavior.
Variations of metallic-nonmetallic
character.
Basic-acidic behavior of the oxides.
Group IA, Alkali Metals
Largest atomic radii
React violently with water to form H2
Readily ionized to 1+
Oxides dissolved in water form basic solutions
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Group IIA, Alkali Earth Metals
Readily ionized to 2+
React with water to form H2
Oxides dissolved in water form basic
solutions
Group III A
Metals (except for boron)
Several oxidation states (commonly 3+)
Group IV A
Form the most covalent compounds
Oxidation numbers vary between 4+ and 4-
Group V A
Form anions generally(1-, 2-, 3-), though positive
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oxidation states are possible
Question
Sodium oxide, Na2O, is added to excess water.
Which one of the following would occur?
1) reaction to yield an aqueous solution of sodium
hydroxide
2) reaction to yield hydrogen and an aqueous
solution of sodium hydroxide
3) reaction to yield oxygen and an aqueous
solution of sodium hydroxide
4) reaction to yield oxygen, hydrogen, and an
aqueous solution of sodium hydroxide
5) no reaction
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Question
An unknown element, Y, reacts with oxygen to
form an oxide with the general formula Y2O3. The
pure element is paramagnetic and the oxide, when
added to water, forms an acidic solution. Which of
the following elements could be described by these
characteristics?
1) Na
2) N
3) Ca
4) S
5) F
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Group VI A
anions generally, though positive oxidation states
are possible
React vigorously with alkali and alkali earth metals,
and nonmetals
Halogens
Form monoanions
High electronegativity (electron affinity)
Diatomic gases
Most reactive nonmetals (F)
Noble Gases
Minimal reactivity
Monatomic gases;
Closed shell
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