Chapter 8
Periodic Properties & Electron Configurations
Bushra Javed
1
Contents
•
•
•
•
Electron Spin
Electron Configurations of elements
The development of Periodic Table
Periodic Trends
2
Electron Configurations
• Quantum-mechanical theory describes the
behavior of electrons in atoms
• The electrons in atoms exist in orbitals
• A description of the orbitals occupied by
electrons is called an electron configuration
principal energy level of
orbital occupied by the
electron
1s1
number of electrons in the
orbital
sublevel of orbital occupied
by the electron
3
How Electrons Occupy Orbitals
• Calculations with Schrödinger’s equation show
hydrogen’s one electron occupies the lowest
energy orbital in the atom
• Schrödinger’s equation calculations for multielectron atoms cannot be exactly solved due to
electron-electron interactions in multi-electron
atoms
4
How Electrons Occupy Orbitals
• The interactions that occur in multi-electron
atoms are due to :
• electron spin
• & energy splitting of sublevels
5
The Property of Electron Spin
• Spin is a fundamental property of all electrons
• All electrons have the same amount of spin
• The orientation of the electron spin is quantized, it
can only be in one direction or its opposite
– spin up or spin down
• The electron’s spin adds a fourth quantum number
to the description of electrons in an atom, called
the Spin Quantum Number, ms
– not in the Schrödinger equation
6
The Property of Electron Spin
• Experiments by Stern and Gerlach showed a
beam of silver atoms is split in two by a magnetic
field.
• The experiment reveals that the electrons spin on
their axis
• spinning charged particles generate a magnetic
field
Electron Spin
• If there is an even number of electrons, about
half the atoms will have a net magnetic field
pointing “north” and the other half will have a
net magnetic field pointing “south”
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Electron Spin
9
The two
possible spin
orientations
of an electron
and the
conventions
for ms are
illustrated
here.
ms, and Orbital Diagrams
• ms can have values of +½ or −½
• Orbital Diagrams use a square to represent each
orbital and a half-arrow to represent each
electron in the orbital
• By convention, a half-arrow pointing up is used
to represent an electron in an orbital with spin
up
• Spins must cancel in an orbital
– paired
11
Orbital Diagrams
• We often represent an orbital as a square or a
circle and the electrons in that orbital as arrows
– the direction of the arrow represents the spin of the
electron
unoccupied
orbital
orbital with
one electron
orbital with
two electrons
Pauli Exclusion Principle
• No two electrons in an atom may have the same set
of four quantum numbers
• Therefore no orbital may have more than two
electrons, and they must have the opposite spins
• Knowing the number orbitals in a sublevel allows us
to determine the maximum number of electrons in
the sublevel
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Pauli Exclusion Principle
s sublevel has 1 orbital, therefore it can hold
2 electrons
p sublevel has 3 orbitals, therefore it can
hold 6 electrons
d sublevel has 5 orbitals, therefore it can
hold 10 electrons
f sublevel has 7 orbitals, therefore it can hold
14 electrons
14
Sublevel Splitting in Multi-electron Atoms
• The sublevels in each principal energy shell of
Hydrogen all have the same energy
• We call orbitals with the same energy degenerate
• For multi-electron atoms, the energies of the sublevels
are split
– caused by charge interaction, shielding and penetration
• The lower the value of the l quantum number, the less
energy the sublevel has
– s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3)
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Shielding & Effective Nuclear Charge
• Each electron in a multi-electron atom experiences
both the attraction to the nucleus and repulsion by
other electrons in the atom
• These repulsions cause the electron to have a net
reduced attraction to the nucleus – it is shielded
from the nucleus
• The total amount of attraction that an electron
feels for the nucleus is called the effective nuclear
charge of the electron
Shielding & Penetration
17
penetration & Effective Nuclear Charge
• The closer an electron is to the nucleus, the
more attraction it experiences
• The better an outer electron is at penetrating
through the electron cloud of inner electrons,
the more attraction it will have for the nucleus
• Penetration causes the energies of sublevels in
the same principal level to not be degenerate
Effect of Penetration and Shielding
• In the fourth and fifth principal levels, the effects of
penetration become so important that the s orbital lies
lower in energy than the d orbitals of the previous
principal level
• The energy separations between one set of orbitals and
the next become smaller beyond the 4s
– the ordering can therefore vary among elements
– causing variations in the electron configurations of
the transition metals and their ions
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Filling the Orbitals with Electrons
• Energy levels and sublevels fill from lowest energy
to high
s→p→d→f
Aufbau Principle
• Orbitals that are in the same sublevel have the
•
same energy
No more than two electrons per orbital
Pauli Exclusion Principle
• When filling orbitals that have the same energy,
place one electron in each before completing pairs
Hund’s Rule
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Filling the Orbitals with Electrons
The lowest-energy configuration of an atom is
called its ground state.
Any other allowed configuration represents an
excited state.
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Electron Configuration of Atoms
• The electron configuration is a listing of the sublevels
in order of filling with the number of electrons in
that sublevel written as a superscript.
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
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‘Building up’ order of Filling in sublevels
(Ground State Electron Configurations)
Start by drawing a diagram
putting each energy shell on
a row and listing the sublevels,
(s, p, d, f), for that shell in
order of energy (left-to-right)
Next, draw arrows through
the diagonals, looping back
to the next diagonal
each time
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
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Electron Configuration & the Periodic
Table
• The Group number corresponds to the number
of valence electrons
• The length of each “block” is the maximum
number of electrons the sublevel can hold
• The Period number corresponds to the principal
energy level of the valence electrons
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s1
1
2
3
4
5
6
7
s2
p1 p2 p3 p4 p5 s2
p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 f14d1
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Blocks of Elements
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Blocks of Elements
For main-group (representative) elements, an s
or a p subshell is being filled.
For d-block transition elements, a d subshell is
being filled.
For f-block transition elements, an f subshell is
being filled.
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Electron Configurations
Example 1
Write the ground state electron configuration of
the chlorine atom, Cl,
For chlorine, Cl, Z = 17.
1s2 2s2 2p6 3s2 3p5
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Electron Configurations
Example 2
Write the ground state electron configuration of
the manganese atom, Mn
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Electron Configurations
Example 3: What is the ground-state electron
configuration of tantalum (Ta)?
a. 1s22s22p63s23p64s23d104p65s24d105p6 6s25d104f7
b. 1s22s22p63s23p64s23d104p65s24d105p6 6s24f145d3
c. 1s22s22p63s23p64s23d104p65s24d105p6 6s2 5d3
d. 1s22s22p63s23p64s23d104p65s24d105p6 6s24f14
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Electron Configurations
Example 4
Which of the following electron configurations
corresponds to the ground state electron configuration
of an atom of a transition element?
a.
b.
c.
d.
1s22s22p2
1s22s22p63s23p5
1s22s22p63s23p64s2
1s22s22p63s23p63d54s2
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Exceptions to the ‘Building up’ order of Filling
• About 21 of the predicted configurations are
inconsistent with the actual configurations
observed.
One Possible Reason:
• Half filled and fully filled subshells are highly
Stable
• Cr, Cu ,Ag and U are some of the exceptions.
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Exceptions to the ‘Building up’ order of
Filling
Chromium (Z=24) and copper (Z=29) have been
found by experiment to have the following groundstate electron configurations:
Cr:
Cu:
1s2 2s2 2p6 3s2 3p6 3d5 4s1
1s2 2s2 2p6 3s2 3p6 3d10 4s1
In each case, the difference is in the 3d and 4s
subshells.
8 | 33
Exceptions to the ‘Building up’ order of Filling
•
•
•
•
•
•
Expected
Cr = [Ar]4s23d4
Cu = [Ar]4s23d9
Mo = [Kr]5s24d4
Ru = [Kr]5s24d6
Pd = [Kr]5s24d8
•
•
•
•
•
•
Found Experimentally
Cr = [Ar]4s13d5
Cu = [Ar]4s13d10
Mo = [Kr]5s14d5
Ru = [Kr]5s14d7
Pd = [Kr]5s04d10
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Electron Configurations
Example 5
Which of the following electron configurations
represents an allowed excited state of the indicated
atom?
a.
b.
c.
d.
He: 1s2
Ne: 1s2 2s2 2p6
Na: 1s2 2s2 2p6 3s2 3p2 4s1
P: 1s2 2s2 2p6 3s2 3p2 4s1
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The Noble Gas core
Electron Configuration
• The noble gases have eight valence
electrons.
– except for He, which has only two
electrons
• We know the noble gases are
especially non-reactive
– He and Ne are practically inert
• The reason the noble gases are so
non-reactive is that the electron
configuration of the noble gases is
especially stable
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Noble Gas Core Electron Configuration
• A short-hand way of writing an electron
configuration is to use the symbol of the
previous noble gas in [] to represent all the
inner electrons, then just write the last set.
Rb = 37 electrons =
1s22s22p63s23p64s23d104p65s1 = [Kr]5s1
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Noble Gas Core Electron Configuration
Example 6
Which ground-state electron configuration is
incorrect?
a.
b.
c.
d.
Fe: [Ar] 3d5
Ca: [Ar] 4s2
Mg: [Ne] 3s2
Zn: [Ar] 3d10 4s2
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Writing Electron Configurations
The pseudo-noble-gas core includes the noblegas subshells and the filled inner, (n – 1), d
subshell.
For bromine, the pseudo-noble-gas core is
[Ar]3d10
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Hund’s rule
Hund’s rule states:
that the lowest-energy arrangement of
electrons in a subshell is obtained by putting
electrons into separate orbitals of the subshell
with the same spin before pairing electrons
40
Electron Configurations
41
Electron Configurations
Example 7
Draw an orbital diagram for nitrogen.
1s
2s
2p
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Electron Configurations
Example 8
Which of the following electron configurations
or orbital diagrams are allowed and which are
not allowed by the Pauli exclusion principle? If
they are not allowed, explain why?
a. 1s22s12p3
b. 1s22s12p8
c. 1s22s22p63s23p63d8
d. 1s22s22p63s23p63d11
e.
1s
2s
43
Electron Configurations
Example 9
write the complete ground state orbital diagram
and electron configuration of potassium
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Valence Electrons
• The electrons in all the sublevels with the
highest principal energy shell are called
the valence electrons
• Electrons in lower energy shells are called
core electrons
• Chemists have observed that one of the
most important factors in the way an atom
behaves, both chemically and physically, is
the number of valence electrons
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46
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Valence- shell configuration
For main-group elements, the valence configuration
is in the form
nsAnpB
The sum of A and B is equal to the group number.
So, for an element in Group VA of the third period,
the valence configuration is
3s23p3
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Valence-shell configuration
Example 10
What are the valence-shell configuration for
arsenic and zinc?
Arsenic is in period 4, Group VA.
Its valence configuration is 4s24p3.
Zinc, Z = 30, is a transition metal in
the first transition series.
Its noble-gas core is Ar, Z = 18.
Its valence configuration is 4s23d10.
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Valence-shell configuration
Example: 11
What is the valence -shell configuration of Tc, Z = 43
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Mendeleev’s Periodic Table
• Order elements by atomic mass
• Saw a repeating pattern of properties
• Periodic Law – when the elements are
arranged in order of increasing atomic mass,
certain sets of properties recur periodically
• Put elements with similar properties in the
same column
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Mendeleev’s Periodic Table
• Used pattern to predict properties of
undiscovered elements
• Mendeleev’s Periodic Law allows us to predict
what the properties of an element will be
based on its position on the table
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Mendeleev's Predictions
53
Everyone Wants to Be Like a Noble Gas!
The Alkali Metals
• The alkali metals have one more
electron than the previous noble gas
• In their reactions, the alkali metals
tend to lose one electron, resulting
in the same electron configuration as
a noble gas
– forming a cation with a 1+ charge
54
Periodic Trends
Periodic law states that when the elements are
arranged by atomic number, their physical and
chemical properties vary periodically.
We will look in more detail at three periodic
properties:
1. atomic radius,
2. ionization energy, and
3. electron affinity.
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Atomic Radius
While an atom does not have a definite size, we
can define it in terms of covalent radii (the
radius in covalent compounds).
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Trend in Atomic Radius – Main Group
• There are several methods for
•
•
measuring the radius of an atom,
and they give slightly different
numbers
van der Waals radius = nonbonding
covalent radius = bonding radius
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Trend in Atomic Radius – Main Group
• Atomic Radius Increases down group
valence shell farther from nucleus
effective nuclear charge fairly close
• Atomic Radius Decreases across period (left to
right)
adding electrons to same valence shell
effective nuclear charge increases
valence shell held closer
58
Trend in Atomic Radius – Main Group
Trends
Within each group (vertical column), the atomic
radius increases with the period number.
This trend is explained by the fact that each
successive shell is larger than the previous shell.
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Trend in Atomic Radius – Main Group
Within each period (horizontal row), the atomic
radius tends to decrease with increasing atomic
number (nuclear charge).
60
A representation of atomic radii is shown below.
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Trend in Atomic Radius – Main Group
Example 12:
Refer to a periodic table and arrange the
following elements in order of increasing atomic
radius: Br, Se, Te.
34
35
Se
Br
Te is larger than Se.
Se is larger than Br.
Br < Se < Te
52
Te
62
Trend in Atomic Radius – Main Group
Example 13
An atom of which of the following elements has
the largest atomic radius?
a) Rb
b)Cl
c) Mg
d) P
63
Trend in Atomic Radius – Main Group
Example 14
An atom of which of the following elements has
the smallest atomic radius?
a) Mg
b) Cl
c) As
d) Rb
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Trend in Ionization Energy
• Minimum energy needed to remove an
electron from an atom or ion in the gaseous
state
• endothermic process
• valence electron easiest to remove, lowest IE
•
•
M(g) + IE1  M1+(g) + 1 eM+1(g) + IE2  M2+(g) + 1 efirst ionization energy = energy to remove electron from neutral atom;
2nd IE = energy to remove from 1+ ion; etc.
65
Trend in Ionization Energy
• Going down a group, first ionization energy
decreases.
• This trend is explained by understanding that
the smaller an atom, the harder it is to
remove an electron, so the larger the
ionization energy.
• In general : E1< IE1< IE3
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Trend in Ionization Energy
Generally, ionization energy increases with atomic
number.
Ionization energy is proportional to the effective
nuclear charge divided by the average distance
between the electron and the nucleus.
Because the distance between the electron and the
nucleus is inversely proportional to the effective
nuclear charge, ionization energy is inversely
proportional to the square of the effective nuclear
charge.
67
Trend in Ionization Energy
Example 15
Refer to a periodic table and arrange the
following elements in order of increasing
ionization energy: As, Br, Sb.
33
35
As
Br
51
Sb
Sb is larger than As.
As is larger than Br.
Ionization energies:
Sb < As < Br
8 | 68
Trend in Ionization Energy
Small deviations occur between Groups IIA and IIIA and
between Groups VA and VIA.
Examining the valence configurations for these groups
helps us to understand these deviations:
It takes less energy to remove the
IIA
ns2
np1 electron than the ns2 electron.
IIIA ns2np1
VA
VIA
ns2np3
ns2np4
It takes less energy to remove the
np4 electron than the np3 electron.
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Trend in Ionization Energy
Example16
An atom of which of the following elements has
the largest second ionization energy?
a) Li
b) C
c) F
d) Be
70
Trend in Ionization Energy
Example 17
Which of the following elements has larger first
IE and WHY ?
a)
b)
c)
d)
Al
N
O
K
71
Electron affinity (E.A.)
The energy change for the process of adding an
electron to a neutral atom in the gaseous state to
form a negative ion.
A negative energy change (exothermic) indicates a
stable anion is formed. The larger the negative
number, the more stable the anion. Small negative
energies indicate a less stable anion.
A positive energy change (endothermic) indicates
the anion is unstable.
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Trends in Electron affinity (E.A.)
Electron affinities in the main-group elements show
a periodic variation when plotted against atomic
number, although this variation is somewhat more
complicated than that displayed by ionization
energies.
In a given period, the electron affinity rises from the
Group IA element to the Group VIIA element but
with sharp drops in the Group IIA and Group VA
elements.
73
8 | 74
Trends in Electron affinity (E.A.)
Broadly speaking, the trend is toward more
negative electron affinities going from left to
right in a period.
Let’s explore the periodic table by group.
75
76
Trend in Electron Affinity (EA)
Example 18
An atom of which of the following elements has the
most negative electron affinity?
a) K
b) Cl
c) Br
d) Se
e) N
77
Groups IIA and VIIIA do not form stable anions;
their electron affinities are positive.
Group Valence
IA
ns1
IIIA ns2np1
IVA ns2np2
VA ns2np3
VIA ns2np4
VIIA ns2np5
Anion Valence
ns2
stable
ns2np2
stable
ns2np3
stable
ns2np4
not so stable
ns2np5
very stable
ns2np6
very stable
Except for the members of Group VA, these values
become increasingly negative with group number.
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Metallic Character
Elements with low ionization energies tend to
be metals. Those with high ionization energies
tend to be nonmetals. This can vary within a
group as well as within a period.
79
Metallic Character
• Metallic character is how closely an element’s
properties match the ideal properties of a metal
– more malleable and ductile, better conductors, and
easier to ionize
• Metallic character decreases left-to-right across
a period
– metals are found at the left of the period and
nonmetals are to the right
• Metallic character increases down the column
– nonmetals are found at the top of the middle Main
Group elements and metals are found at the bottom
80
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Oxides of Main Group Elements
Oxides
A basic oxide reacts with acids. Most metal oxides
are basic. If soluble, their water solutions are basic.
An acidic oxide reacts with bases. Most nonmetal
oxides are acidic. If soluble, their water solutions
are acidic.
An amphoteric oxide reacts with both acids and
bases.
82
Trends in Metallic Character
Group IA, Alkali Metals (ns1)
These elements are metals; their reactivity
increases down the group.
The oxides have the formula M2O.
Hydrogen is a special case. It usually behaves as a
nonmetal, but at very high pressures it can exhibit
metallic properties.
83
Reactions of Alkali Metals with Water
• Alkali metals are oxidized to the 1+ ion
• H2O is split into H2(g) and OH− ion
• The Li, Na, and K are less dense than the water
– so float on top
• The ions then attach together by ionic bonds
• The reaction is exothermic, and often the
heat released
ignites the H2(g)
84
Trends in Metallic Character
Group IIA, Alkaline Earth Metals (ns2)
These elements are metals; their reactivity
increases down the group.
The oxides have the formula MO.
85
Trends in Metallic Character
Group IIIA (ns2np1)
Boron is a metalloid; all other members of Group
IIIA are metals.
The oxide formula is R2O3.
B2O3 is acidic.
Al2O3 and Ga2O3 are amphoteric
The Group IIIA oxides are basic.
86
Trends in Metallic Character
Group IVA (ns2np2)
Carbon is a nonmetal; silicon and germanium are
metalloids; tin and lead are metals.
The oxide formula is RO2 and, for carbon and lead,
RO.
CO2, SiO2, and GeO2 are acidic (decreasingly so).
SnO2 and PbO2 are amphoteric.
87
Some oxides of Group IVA
PbO
(yellow)
PbO2
(dark brown)
SnO2 (white)
SiO2
(crystalline solid quartz)
88
Group VA (ns2np3)
Nitrogen and phosphorus are nonmetals; arsenic
and antimony are metalloids; bismuth is a metal.
The oxide formulas are R2O3 and R2O5, with some
molecular formulas being double these.
Nitrogen, phosphorus, and arsenic oxides are acidic.
Antimony oxides are amphoteric.
Bismuth oxide is basic.
89
Group VIA, Chalcogens (ns2np4)
Oxygen, sulfur, and selenium are nonmetals;
tellurium is a metalloid; polonium is a metal.
The oxide formulas are RO2 and RO3.
Sulfur, selenium, and tellurium oxides are acidic
except for TeO2, which is amphoteric.
PoO2 is also amphoteric.
90
Group VIIA, Halogens (ns2np5)
These elements are reactive nonmetals, with the
general molecular formula being X2. All isotopes
of astatine are radioactive with short half-lives.
This element might be expected to be a
metalloid.
Each halogen forms several acidic oxides that
are generally unstable.
91
Group VIIIA, Noble Gases (ns2np6)
These elements are generally unreactive, with
only the heavier elements forming unstable
compounds. They exist as gaseous atoms.
92