Chapter 8
Periodic Relationships Among
the Elements
Electron Configurations of the Elements
Valence Electrons
• the outer electrons of an atom, the ones involved in
Bonding
• for the representative elements (main group elements), the
number of valence electrons is equal to the A Group Number
Which of the following electron configurations do represent
similar chemical properties of their atoms?
(i)1s22s22p63s2
(ii) 1s22s22p3
(iii) 1s22s22p63s23p64s23d104p5
(iv) 1s22s1
(v) 1s22s22p6
(vi) 1s22s22p63s23p3
(ii), (vi)
4f
5f
ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Ground State Electron Configurations of the Elements
Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
Atoms gain electrons
so that anion has a
noble-gas outer
electron configuration.
Atoms lose electrons so that
cation has a noble-gas outer
electron configuration.
Remember: electrons are first
removed from orbitals with the
highest principal quantum number.
H 1s1
H- 1s2 or [He]
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
-1
-2
-3
+3
+2
+1
Cations and Anions Of Representative Elements
Na+: [Ne]
Al3+: [Ne]
O2-: 1s22s22p6 or [Ne]
F-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
Isoelectronic: have the same number of electrons
What neutral atom is isoelectronic with H- ?
H-: 1s2
same electron configuration as He
Which of the following pairs are not isoelectronic with Ar?
(a)K+ and Cl(b) Al3+and N3(c)Rh3+ and Ir3+
(d) Ca2+ and S2(e) P3- and Sc3+
(a) K+ [Ar]
Cl- [Ar]
(b) Al3+ [Ne]
N3- [Ne]
(c) Rh3+ [Rn]5s14d5
(d) Ca2+ [Ar]
(e) P3- [Ar]
NOT
Ir3+ [Rn] 5s24d4
S2- [Ar]
Sc3+ [Ar]
NOT
Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal,
electrons are always removed first from the ns orbital and
then from the (n – 1)d orbitals.
Fe:
[Ar]4s23d6
Fe2+: [Ar]4s03d6 or [Ar]3d6
Fe3+: [Ar]4s03d5 or [Ar]3d5
Mn:
[Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
A metal ion with a net +3 charge has five electrons in the
3d subshell. What is this metal?
(a) Cr (b) Mn (c) Fe (d) Co (e) Ni
This species has +3 charges, which indicates that it has
three more protons than the electrons.
According to the question that it has five electrons in the 3d
subshell, and thus the total electrons in valence shells for
its atomic type will be 8.
Note that the transition metals (with d electrons) losing its s
electrons prior to its d electrons and thus the valence
electron configuration will be 4s23d6, which is Fe.
Periodic Variation in Physical Properties
Atomic Radius
Sizes of atoms: Trends
• size increases while going down a Group
• Why?
• Because orbital size increases with increasing n
number, the size of atom increases. The highest energy
electrons can be farther away from the nucleus.
• size decreases going across Period
• Why?
• effective nuclear charge increases. The larger the
effective nuclear charge, the stronger hold of the nucleus
on electrons. The electron clouds shrink.
• Remember n does NOT increases--e- cannot be farther
from nucleus.
Ionic Radius
Cation is always smaller than atom from which it is formed. This is because
the nuclear charge remains the same but the reduced electron repulsion
resulting from removal of electrons make the electron clouds shrink.
Anion is always larger than atom from which it is formed. This is because
the nuclear charge remains the same but electron repulsion resulting from
the additional electron enlarges the electron clouds.
•From top to bottom, both the atomic and ionic radius increase
within group.
•Across a period the anions are usually larger than cations.
• For ions derived from different groups, size comparison is
meaningful only if the ions are isoelectronic.
E.g. Na+ (Z=11) is smaller than F- (Z=9).
The larger effective charge results in a smaller radius.
Radius of tripositive ions < dipositive ions<unipositive ions
Al3+ <Mg2+ <Na+
Radius of uninegative ions < dinegative ions < trinegative ions
N3- > O2->F-
Which of the following is a correct order of atomic radii?
(a)Ar < P < Na < Ca < Cs
(b) Cs < Rb < K < Na < Li
(c) Na < Mg < Al < Si < P
(d) Rb < Mg < C < F < He
(e) Li < H < Al < K < Ar
(a)
Which of the following is a correct order of ionic radii?
(a)Na+ < Mg2+ < Al3+ < O2- < F- < N3(b) Al3+ < Mg2+ < Na+ < F- < O2- < N3(c) F- < O2- < N3- < Na+ < Mg2+ < Al3+
(d) Al3+ < Mg2+ < Na+ < N3- < O2- < F(e) Na+ < Mg2+ < Al3+ < O2- < N3- < F(b)
Ionization energy (I.E.) is the minimum energy (kJ/mol)
required to remove an electron from a gaseous atom in its
ground state.
I1 + X (g)
I2 + X+(g)
I3 +
X2+
(g)
X+(g) + e-
I1 first ionization energy, remove 1st electron
X2+(g) + e- I2 second ionization energy, remove 2nd electron
X3+(g)
+
rd electron
I
third
ionization
energy,
remove
3
3
e
I1 < I2 < I3
General Trend in First Ionization Energies
Increasing First Ionization Energy
Increasing First Ionization Energy
First IE:
Trends down a Group
• remember that Cs is more reactive than Na
• Easier to remove an e- from Cs than from Na
• As size increases, I.E. decreases
Trends across a Period
• remember that Cl gains rather than loses an e• Easier to remove an e- from Al than from Cl
• As size decreases, I.E. increases
Exceptions: Ionization Energy
1.Occur between Group 2A and 3A
Group 3A: ns2np1. one single electron in p orbital.
• removing the first p electron is less than expected
•Why? This p electron is shielded by inner ns2 electrons.
Less energy is needed to remove a single p electron than
to remove a pair of s electron of the same n level.
2. Occur Group 5A and 6A
5A: ns2np3 6A: ns2np4
• removing the fourth p electron is less than expected
•Why? 2nd electron in a p orbital increases the
electron-electron repulsion, which makes it easier to
ionize an atom of Group 6A.
Arrange the following in order of the increasing
first ionization energy: Na, Cl, Al, S, Cs
Hint: Ionization energy increases across a row of
the periodic table and decreases down a column
or group.
Cs < Na < Al < S < Cl
Electron affinity (EA) is the negative of the energy change
that occurs when an electron is accepted by an atom in the
gaseous state to form an anion.
X (g) + e-
X-(g)
F (g) + e-
F-(g)
DH = -328 kJ/mol
EA = +328 kJ/mol
O (g) + e-
O-(g)
DH = -141 kJ/mol
EA = +141 kJ/mol
Variation of Electron Affinity With Atomic Number (H – Ba)
Overall trend:
increases from left to right
across the period.
The values varied little in
the group.
The EA of metals are
generally lower than those
of nonmetals.
Properties of Oxides Across a Period
basic
acidic
Across a period, oxides change from basic to amphoteric to
acidic. Going down a group, the oxides become more basic.