Periodic Trends: Periodic Properties
 Certain physical and chemical properties recur at
regular intervals, and/or vary in regular fashion,
when the elements are arranged according to
increasing atomic number.
 Melting point, boiling point, hardness, density,
physical state, and chemical reactivity are periodic
properties.
 We will examine several periodic properties that
are readily explained using electron
configurations.
Periodic Trends
• Atomic Size –
– Group – increase in size down the group
• (addition of energy levels)
– Period – decrease in size across a period
• (increase in the # protons in the nucleus
pulling on the same energy level)
Periodic Properties: Atomic Radius
 Half the distance between the
nuclei of two atoms is the
atomic radius.
Covalent radius: half the
distance between the nuclei
of two identical atoms joined
in a molecule.
Metallic radius: half the
distance between the nuclei
of adjacent atoms in a solid
metal.
Periodic Properties: Atomic Radius
• Atomic radius increases from top to bottom
within a group.
• The value of n increases, moving down the
periodic table.
• The value of n relates to the distance of an
electron from the nucleus.
Periodic Properties: Atomic Radius
• Atomic radius decreases from left to right within a period.
• Why? The effective nuclear charge increases from left to
right, increasing the attraction of the nucleus for the valence
electrons, and making the atom smaller.
Mg has a greater
effective nuclear
charge than Na, and
is smaller than Na.
Multielectron and Shielding
Effective Nuclear Charge (Zeff): The nuclear charge actually felt by an electron.
Zeff = Zactual - Electron shielding
Atomic Radii of the Elements
Example
With reference only to a periodic table, arrange each set of
elements in order of increasing atomic radius:
(a) Na, Al, P
(b) Te, S, Se
(c) Ge, Sb, Sn
Ionic Radii
The ionic radius of
each ion is the
portion of the
distance between
the nuclei occupied
by that ion.
Ionic Radii of Cations
• Cations are smaller than the atoms from which they
are formed; the value of n usually decreases. Also,
there is less electron–electron repulsion.
Ionic Radii of Anions
• Anions are larger than the
atoms from which they are
formed.
• Effective nuclear charge is
unchanged, but additional
electron(s) increase electron–
electron repulsion.
• Isoelectronic species have the
same electron configuration;
size decreases with effective
nuclear charge.
Atomic Radii and the Periodic Table
Example
Use the standard periodic table to arrange the following species in
the expected order of increasing radius:
Sr2+, Ru3+, Rb+, I–, Br–
Isoelectronic atoms= Sr+2, Rb+1, Br-1
Atom
Protons
Electrons
Configuration
Rank
Sr+2
38
36
[Kr]5s2 
[Kr]
1
Ru+3
44
41
[Kr]5s24d6 
[Kr]4d5
4
Rb+1
37
36
[Kr]5s1 
[Kr]
2
I-1
53
54
[Kr]5s24d105p5
 [Xe]
5
Br-1
35
36
[Ar]4s23d104p5
[Kr]
3
Periodic Trend
• Ionization Energy –
– Group – decreases down the group
• (addition of energy levels makes it harder
for the nucleus to hold on and easier to
remove electrons)
– Period – increases across the period
• (addition of protons pulling on the
electrons makes it harder to remove the
electrons)
Ionization Energy
• Ionization energy (I) is the energy required to remove
an electron from a ground-state gaseous atom.
• I is usually expressed in kJ per mole of atoms.
M(g)  M+(g) + e–
ΔH = I1
M+(g)  M2+(g) + e– ΔH = I2
M2+(g)  M3+(g) + e– ΔH = I3
Ionization Energy Trends
• I1 < I2 < I3
– Removing an electron from a positive ion is more difficult than
removing it from a neutral atom.
• A large jump in I occurs after valence electrons are
completely removed (why?).
• I1 decreases from top to bottom on the periodic table.
– n increases; valence electron is farther from nucleus.
• I1 generally increases from left to right, with exceptions.
– Greater effective nuclear charge from left to right holds electrons
more tightly.
Selected Ionization Energies
Compare I2 to I1 for a 2A
element, then for the
corresponding 1A element.
Why is I2 for each 1A element
so much greater than I1?
Why don’t we see
the same trend for
each 2A element? I2
> I1 … but only
about twice as great
…
Example
Why is lithium’s I2 14 times as much as the I1 while beryllium is
only twice as much?
Li
Li+1
Li+2
Be
Be+1
Be+2
[He]2s1
[He] or 1s2
1s1
[He]2s2
[He]2s1
[He]
Selected Ionization Energies
General trend in I1: An increase
from left to right, but …
…I1 drops, moving
from 2A to 3A.
The electron being removed
is now a p electron (higher
energy, easier to remove than
an s).
I1 drops again between
5A and 6A.
Repulsion of the paired
electron in 6A makes
that electron easier to
remove.
Example
Why is boron's I1 less than that of beryllium?
B
Be
Example
Why is oxygen’s I1 less than that of nitrogen?
O
N
First Ionization Energies
Second Ionization Energies
• The amount of energy required to remove the
second electron once the first is gone.
– Always greater than THEIR first IE due to an INCREASED
effective nuclear charge (a greater proton to electron
ratio), making the atom smaller.
– However, 2nd Ionization energies can be compared
BETWEEN different atoms by evaluating the +1 ion of THAT
atom.
• If the +1 ion is a noble gas the 2nd IE will be LARGE. If the +1 ion is
NOT a noble gas the 2nd IE will be SMALLER.
Second Ionization Energies
Example:
1st IE
2nd IE
Na
Small
(wants to lose 1e-)
Big
(Doesn’t want to lose
2e-)
Bigger than Na
(doesn’t want to lose
any)
Smaller than Na
(once 1e- is gone, it is
easy to lose 2e-)
[Ne]3s2
Mg
[Ne]3s2
Periodic Trends: Periodic Properties
 Certain __________________________recur at
regular intervals, and/or vary in regular fashion,
when the elements are arranged according to
______________________.
 Melting point, boiling point, hardness, density,
physical state, and chemical reactivity are periodic
properties.
 We will examine several periodic properties that
are readily explained using electron
configurations.
Periodic Trends
• Atomic Size –
– Group –
– Period –
Periodic Properties: Atomic Radius
 Half the distance between the
nuclei of two atoms is the
_____________________.
_____________________:
half the distance between the
nuclei of two identical atoms
joined in a molecule.
_____________________:
half the distance between the
nuclei of adjacent atoms in a
solid metal.
Periodic Properties: Atomic Radius
• Atomic radius __________ from top to
bottom within a group.
• The value of n increases, moving down the
periodic table.
• The value of n relates to the __________ of
an electron from the nucleus.
Periodic Properties: Atomic Radius
• Atomic radius __________ from left to right within a period.
• Why? The ________________________increases from left
to right, increasing the attraction of the nucleus for the
valence electrons, and making the atom smaller.
Mg has a greater
effective nuclear
charge than Na, and
is smaller than Na.
Multielectron and Shielding
Effective Nuclear Charge (Zeff): The nuclear charge actually felt by an electron.
Zeff = Zactual - Electron shielding
Atomic Radii of the Elements
Example
With reference only to a periodic table, arrange each set of
elements in order of increasing atomic radius:
(a) Na, Al, P
(b) Te, S, Se
(c) Ge, Sb, Sn
Ionic Radii
The ___________
of each ion is the
portion of the
distance between
the nuclei occupied
by that ion.
Ionic Radii of Cations
• __________ are _________ than the atoms from
which they are formed; the value of n usually
decreases. Also, there is less electron–electron
repulsion.
Ionic Radii of Anions
• ________ are _______than
the atoms from which they are
formed.
• Effective nuclear charge is
unchanged, but additional
electron(s) increase electron–
electron repulsion.
• _____________ species have
the same electron
configuration; size decreases
with effective nuclear charge.
Atomic Radii and the Periodic Table
Example
Use the standard periodic table to arrange the following species in
the expected order of increasing radius:
Sr2+, Ru3+, Rb+, I–, Br–
Periodic Trend
• Ionization Energy –
– Group –
– Period –
Ionization Energy
• ____________________(I) is the energy required to
remove an electron from a ground-state gaseous atom.
• I is usually expressed in kJ per mole of atoms.
M(g)  M+(g) + e–
ΔH = I1
M+(g)  M2+(g) + e– ΔH = I2
M2+(g)  M3+(g) + e– ΔH = I3
Ionization Energy Trends
• I1 < I2 < I3
– Removing an electron from a __________ is more difficult than
removing it from a ____________.
• A large jump in I occurs after valence electrons are
completely removed (why?).
• I1 ____________ from top to bottom on the periodic table.
– n increases; valence electron is farther from nucleus.
• I1 generally _________ from left to right, with exceptions.
– Greater effective nuclear charge from left to right holds electrons
more tightly.
Selected Ionization Energies
Compare I2 to I1 for a 2A
element, then for the
corresponding 1A element.
Why is I2 for each 1A element
so much greater than I1?
Why don’t we see
the same trend for
each 2A element? I2
> I1 … but only
about twice as great
…
Example
Why is lithium’s I2 14 times as much as the I1 while beryllium is
only twice as much?
Selected Ionization Energies
General trend in I1: An increase
from left to right, but …
…I1 drops, moving
from 2A to 3A.
The electron being removed
is now a p electron (higher
energy, easier to remove than
an s).
I1 drops again between
5A and 6A.
Repulsion of the paired
electron in 6A makes
that electron easier to
remove.
Example
Why is boron's I1 less than that of beryllium?
Example
Why is oxygen’s I1 less than that of nitrogen?
First Ionization Energies
Second Ionization Energies
• The amount of energy required to remove the
second electron once the first is gone.
– Always greater than _____ first IE due to an ____________
effective nuclear charge (a greater proton to electron
ratio), making the atom smaller.
– However, 2nd Ionization energies can be compared
________ different atoms by evaluating the +1 ion of ____
atom.
• If the +1 ion is a noble gas the 2nd IE will be _____. If the +1 ion is
NOT a noble gas the 2nd IE will be _________.
Second Ionization Energies
Example:
1st IE
2nd IE
Na
Small (wants to lose
1e-)
Big (Doesn’t want to
lose 2e-)
Mg
Bigger than Na
(doesn’t want to lose
any)
Smaller than Na (once
1e- is gone, it is easy
to lose 2e-)