Topic 8
Periodic Properties of
the Elements
1
Electron Configuration
An “electron configuration” of an atom is a
particular distribution of electrons among
available subshells.
The notation for a configuration lists the subshell
symbols (s, p, d, f) sequentially with a superscript
indicating the number of electrons occupying that
subshell.
For example, lithium (atomic number 3) has two
electrons in the “1s” sub shell and one electron in
the “2s” sub shell 1s2 2s1.
2
Electron Configuration
Ground state configuration refers to the most stable
(lowest energy) way of assigning the electrons.
All other ways are higher energy or excited state
configurations that are unstable and caused by absorbing
light or collisions with other particles.
Excited atoms eventually revert to their ground state
configuration either by giving off light or as a result of
collisions with other particles.
If ground or excited state is not specified, we assume the
term electron configuration refers to the ground state.
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The Pauli Exclusion Principle
The maximum number of electrons allowed per
subshell based on Pauli’s Exclusion Principle:
Sub shell
Number of
Orbitals
Maximum
Number of
Electrons
s (l = 0)
1
2, s2
p (l = 1)
3
6, p6
d (l =2)
5
10, d10
f (l =3)
7
14,
f14
4
Aufbau Principle
The Aufbau principle is a scheme used to reproduce
the ground state electron configurations of atoms by
following the “building up” order.
Listed below is the order in which all the possible subshells fill with electrons.
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f
lowest energy
highest energy
This order is based on increasing energy levels using the following
factors:
• Orbitals with higher (n + l values) have higher energies
i.e. 3d (3+2=5) is higher energy than 4s (4+0=4)
because the n+l value is higher (note: d, l = 2 and s, l=0).
• For orbitals with the same (n + l values), the orbital with the higher n
have higher energies
i.e. 3s (3+0=3) has the same n+l value as 2p (2+1=3) but 3s is
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higher energy because of the higher n of 3.
Configurations and the Periodic Table
We can use the periodic table as a memory aid for the
typical order of filling orbitals.
p block – 3 orbital with max
2e- each = 6 electrons ; p6 max
s block – 1 orbital with max
2e- = 2 electrons; s2 max
d block – 5 orbital with max 2eeach = 10 electrons; d10 max
f block – 7 orbital
with max 2e- each
= 14 electrons;
f14 max
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To write the ground state electron configuration of an atom, you first need
to locate the atom on the periodic table. For a neutral atom, the number
of electrons equals the number of protons (atomic number). You will write
the subshells with the number of electrons in the subshell written as a
superscript (nl #e-) from left to right and top to bottom until you reach the
particular atom in question on the periodic table. Basically, each block
represents an electron that you are adding to the configuration.
i.e.
21
𝑆𝑐
(21 electrons to account for in electron config.)
1s2 2s2 2p6 3s2 3p6 4s23d1
Ground state electron config. for Sc
21
Sc
7
i.e. 𝟓𝑩 (5 electrons to account for in ground state electron config.)
1s2 2s2 2p1
i.e.
𝟏𝟎𝑵𝒆
(10 electrons to account for in ground state electron config.)
1s2 2s2 2p6
i.e.
𝟏𝟏𝑵𝒂
(11 electrons to account for in ground state electron config.)
1s2 2s2 2p6 3s1
8
i.e.
𝟖𝟐𝑷𝒃
(82 electrons to account for in ground state electron config.)
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d105p6 6s24f14 5d106p2
There is a short notation for writing electron configurations which
involves using the noble gas core. We specify the noble gas in the
row above the atom in question in brackets representing the electron
configuration to that point and then write the configuration of the
remaining electrons. For the above Pb example, we would write:
[Xe] 6s24f14 5d10 6p2
Xe
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Configurations and the Periodic Table
There are exceptions to the order of filling the
electron configuration: notable exceptions are
Cr, Cu, Nb, Ru, Rh, Pd, Ag, Gd, and Pt.
– The commonly used justification for these
exceptions is that the electrons are “shifted” in
order to have half-filled (d5) or fully-filled (d10)
– There seems to be some stability attained with
half-filled or fully-filled subshells.
– But this justification is tenuous since the
exceptions are not consistent.
i.e. Cu: [Ar] 3d10 4s1 is the actual electron configuration instead
of [Ar] 3d9 4s2 as predicted; one electron has shifted from the “s”
orbital to the “d” level for the stability gained from having the inner
10
shell completely filled.
Valence Electrons
Valence Electrons are electrons that reside in the
outermost shell of an atom - or in other words, those
electrons outside the “noble gas core”.
– These electrons are primarily involved in chemical reactions.
– Elements within a given group (vertical column of periodic table)
have the same “valence shell configuration.”
– This accounts for the similarity of the chemical properties
among groups of elements in the same family (vertical group).
P: [Ne] 3s2 3p3
As: [Ar] 4s2 3d10 4p3
note: 3d electrons are inner electrons and not valence electrons.
Sb: [Kr] 5s2 4d10 5p3
These three atoms reside in the same VA group and have the same
valence electrons: 2 in the s block and 3 in the p block.
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Valence Electrons
Valence electrons for the groups:
Group IA:
ns1
Group IIA:
ns2
Group IIIA: ns2 np1
Group IVA: ns2 np2
Group VA:
ns2 np3
Group VIA: ns2 np4
Group VIIA: ns2 np5
Group VIIIA: ns2 np6
Transition metals: ns2 (n-1)dx where x = 1 to 10
i.e. Sc: [Ar] 4s2 3d1
12
Configurations and the Periodic Table
So far we have done ground state electron
configurations for neutral atoms. What happens if you
have an ion (charged atom)?
Since we are counting electrons, any charge will either
add or subtract the number of electrons we are placing
in the electron configuration meaning the number of
electrons does not equal the number of protons.
For cations, we will subtract electrons from the total.
For anions, we will add electrons to the total.
– Ca2+ subtract 2e- because you have lost 2e– Cl- add 1e- because you have gained 1e-
13
Configurations and the Periodic Table
Here are a few examples:
Ca: 1s2 2s2 2p6 3s2 3p6 4s2
Ca2+: 1s2 2s2 2p6 3s2 3p6 lost two outer shell electrons
Cl:
Cl-:
1s2 2s2 2p6 3s2 3p5
1s2 2s2 2p6 3s2 3p6 gained one electron
Note that Ar (1s2 2s2 2p6 3s2 3p6), Ca2+, and Cl- all have the exact same
electron configurations. We say that these species are isoelectronic with
each other. They have the same arrangement of electrons but are not the
same species; same number of electrons but different number of protons.
Electrons do not come off the same way they go on; the outer most valence
electrons come off before any inner electrons;
V: [Ar] 4s2 3d3
V3+: [Ar] 3d2 lost three electrons; the 2 outermost 4s and
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1 inner 3d. Note we did not lose all 3 in d level.
Orbital diagrams
An orbital diagram gives more detailed
information than an electron configuration. In
an orbital diagram,
• electrons are represented by arrows
• boxes, circles, or blanks are used to represent orbitals
• upward-pointing arrows represent electrons with +1/2 spin and
downward-pointing arrows represent electrons with -1/2 spin
• Pauli’s principle must be followed meaning no more than 2
electrons per orbital with opposite spin.
• We fill the orbitals in the highest occupied subshell singly with
electrons of the same spin before putting a second electron (of
opposite spin) in any orbital. This is known as Hund’s Rule of
Maximum Multiplicity. The justification for Hund’s rule is spin
correlation; electrons with same spin repel each other less than
electrons with opposite spin.
15
Orbital diagrams
An orbital diagram is used to show how the
orbitals of a subshell are occupied by electrons.
– Each orbital is represented by a box, circle or a line.
– Each group of orbitals is labeled by its subshell
notation.
– Electrons are represented by arrows:
up for ms = +1/2
and down for ms = -1/2
– We fill the orbitals in the highest occupied subshell
singly with electrons of the same spin before putting
a second electron (of opposite spin) in any orbital.
Neon:
10e-
1s
2s
2p
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Orbital Diagrams
Here’s another way orbital diagrams are displayed by increasing
energy levels. This example is for Vanadium with 23 electrons:
3d
Energy
4s
3p
3s
2p
2s
1s
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Orbital Diagrams
The maximum number of electrons and their
orbital diagrams are:
Sub shell
Number of
Orbitals
Maximum
Number of
Electrons
s (l = 0)
1
2, s2
p (l = 1)
3
6, p6
d (l =2)
5
10, d10
f (l =3)
7
14,
f14
18
Paramagnetic and Diamagnetic
Substances
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 properties.
Paramagnetic substance – is one that is weakly attracted by a
magnetic field as the result of one or more unpaired electrons.
Whenever the electron configuration of an atom has partially
filled subshells, the atom is paramagnetic. Vanadium on slide 17
had unpaired electrons in the 3d level; therefore, is
paramagnetic.
Diamagnetic substance – is not attracted by a magnetic field
because it has only paired electrons. Neon on slide 16 had all
electrons paired; therefore, is diamagnetic.
HW 61-62
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code: electron
Periodic Properties
Recall that periodic law states that when the
elements are arranged by atomic number,
their physical and chemical properties vary
periodically (hence the name periodic table).
We will look at three periodic properties:
– Atomic radius
– Ionization energy
– Electron affinity
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Periodic Properties
Atomic radius (size)
Covalent radius – of atom is ½ distance between the
nuclei of two like atoms joined in a molecule.
ionic radius same as above.
– Within each period (horizontal row), the atomic
radius tends to increase with decreasing atomic
number (right to left).
– Within each group (vertical column), the atomic
radius tends to increase with the period number (top
to bottom).
increases
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Two factors determine the size of an atom:
– Why increase down?
As we go down the periodic table, the valence shell
gets larger. We are adding more shells as we go
down the table making the outer electrons farther
from the nucleus with less attraction for the nucleus.
– Why decrease across?
The effective nuclear charge, which is the positive
charge an electron experiences from the nucleus
minus any “shielding effects” from intervening
electrons, gets larger as we go across the periodic
table. As the effective nuclear charge increases,
valence electrons will be pulled toward the nucleus
more and held more tightly making the atom smaller.
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Figure:
Representation
of atomic radii
(covalent radii)
of the maingroup elements.
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Ex. Arrange the following in order of increasing atomic radius:
Ca, N, As, F, Ba
increasing means smallest
atomic radius to largest radius
N
Locating the atoms
on the periodic table
and comparing them,
we find
Ca
Ba
F
As
Atomic size periodic trend increases
across from right to left and top to
bottom on the table
(smallest size) F < N < As < Ca < Ba (largest size)
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Ionic Radii
Cations are smaller than neutral atom: Na+ < Na
Cation lost 3s shell
when lost e-; therefore,
cation smaller
Cation has more protons
than electrons; therefore
electrons are pulled more
towards the nucleus
making it smaller
Anion has more electrons than
protons; therefore, electrons are
pulling more outwards from
nucleus making it larger. There is
more electron-electron repulsion
and a smaller effective nuclear
charge causing the anion to get
larger.
Anions are larger than neutral atom: Cl- > Cl
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Within an isoelectronic group of ions, the one with the
greatest nuclear charge (largest excess of protons) will be
the smallest.
Ex. Ar, Cl-, Ca2+, S2-, K+
Z= 18, 17, 20, 16, 19
all have 18e-; same electron configuration but
different number of protons
For isoelectronic groups, the larger the positive charge (higher
effective nuclear charge), the more pull the valence electrons will
have by the nucleus, the smaller the species.
smallest
Ca2+ < K+ < Ar < Cl- < S2- largest
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Periodic Properties
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 in
gaseous state.
For a lithium atom, the first ionization energy is
illustrated by:

Li(1s 2s )  Li (1s )  e
2
1
2

Ionization energy = 520 kJ/mol
Metal atoms tend to give up valence electrons while nonmetal
tend to gain electrons. Energy must be absorbed (endothermic
process) to remove an electron; therefore, we have +ionization 27
energies.
Periodic Properties
Ionization energy (IE)
increases
– There is a general trend that ionization
energies increase across a period (horizontal
row); this trend is the opposite of atomic radius
because it is more difficult to remove an
electron that is closer to the nucleus in smaller
atoms.
– The trend for ionization energies increases up a group
(vertical column).
– Summary: ionization energy increases towards the top
right corner of the periodic table with He having the highest
ionization energy (larger the radius, smaller the IE)
– Note there are some exceptions to the trend such as
between groups IIA and IIIA and VA and VIA.
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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.
• 2nd IE tends to be larger than 1st IE because you are
dealing with ions (charged species) which make it harder
to pull an e- from a charged ion than a neutral species;
therefore, more energy will be required.

2
Li  Li  e

2nd Ionization energy = 7298 kJ/mol
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1st Ionization energy = 520 kJ/mol
Ex. Arrange the following in order of increasing 1st
ionization energy:
increasing means
smallest IE to largest IE
As, Ca, O, N
N O
Locating the atoms
on the periodic table
and comparing them,
we find
Ca
As
IE periodic trend increases
across and up the table
(smallest IE) Ca < As < N < O (largest IE)
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Periodic Properties
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.
– Electron affinity process releases energy
(exothermic process); therefore, tend to be negative
values.


Cl([Ne]3s 3p )  e  Cl ([Ne]3s 3p )
2
5
2
6
Electron Affinity = -349 kJ/mol
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Periodic Properties
Electron Affinity
– The more negative the electron affinity, the more
stable the negative ion that is formed.
– Broadly speaking, the general trend goes from
lower left to upper right as electron affinities
become more negative.
increases in negative value
HW 63
code: trend
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