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Daniel L. Reger
Scott R. Goode
David W. Ball
www.cengage.com/chemistry/reger
Chapter 8
The Periodic Table:
Structure and Trends
Electron Configurations and the Periodic Table
• The periodic table can be divided into
four blocks of elements: elements
with highest energy electrons in s, p,
d, or f subshells.
• The arrangement of the elements in
the periodic table correlates with the
subshell that holds the highest energy
electron.
Example: Electron Configurations
• Using only the periodic table, determine
the electron configurations of Al, Ti, Br,
and Sr.
Electron Configurations of Anions
• For anions, the additional electrons fill
orbitals following the same rules that
applies to atoms.
Cl: [Ne] 3s2 3p5
As: [Ar] 4s2 3d10 4p3
Cl-: [Ne] 3s2 3p6
As3-: [Ar] 4s2 3d10 4p6
• Many stable anions have the same
electron configuration as a noble gas
atom.
Electron Configurations of Cations
• For the electron configurations of cations,
electrons of highest n value are removed first.
For cases of the same n level, electrons are
first removed from the subshell having highest
l.
As: [Ar] 4s2 3d10 4p3 As3+: [Ar] 4s2 3d10
Mn: [Ar] 4s2 3d5
Mn2+: [Ar] 3d5
• NOTE: For d-block atoms, the ns electrons are
removed before the (n-1)d electrons.
Test Your Skill
• Write the electron configurations of the
following ions: (a) N3- (b) Co3+ (c) K+
Isoelectronic Series
• An isoelectronic series is a group of
atoms and ions that contain the same
number of electrons.
• The species S2-, Cl-, Ar, K+, and Ca2+
are isoelectronic – they all have 18
electrons.
Atomic Radii
• An atomic radius
is one half the
distance between
adjacent atoms of
the same element
in a molecule.
198/2 = 99
228/2 = 114
Sum = 215
Size Trends for an Isoelectronic Series
Sizes of the Atoms and Their Cations
• Atoms are always larger than their
cations.
Sizes of the Atoms and Their Cations
• If an atom makes more than one cation,
the higher-charged ion has a smaller size.
Atomic and Ionic Radii
• Anions are always larger than their
atoms.
Test Your Skill
• Identify the larger species of each
pair: (a) Mg or Mg2+ (b) Se or Se2-
Atomic Radii of Main Group Elements
Sizes of Atoms
• The sizes of atoms are impacted by the
effective nuclear charge felt by the
outermost electrons.
Effective Nuclear Charge & Size
• The sizes of atoms increase going
down a group.
Sizes of Atoms
• The increase in effective nuclear charge
causes a size decrease across the
period.
Test Your Skill
• Identify the larger species of each
pair: (a) Mg or Na (b) Si or C
Ionization Energy
• The ionization energy is the energy
required to remove an electron from a
gaseous atom or ion in its electronic
ground state.
Ionization Energies
• An atom has as many ionization
energies as it has electrons.
• Example:
Mg(g) → Mg+(g) + eI1 = first ionization energy
Mg+(g) → Mg2+(g) + eI2 = second ionization energy
Trends in 1st Ionization Energies
• The increase in the
effective nuclear
charge across a
period causes an
increase in the
ionization energy
as you go across
that period.
Trends in 1st Ionization Energies
Trends in 1st Ionization Energies
• The slight dip in
ionization energy
for O is because
the fourth p
electron now pairs
with another
electron, slightly
repelling each
other.
Trends in First Ionization Energies
Ionization Energy Trends in Isoelectronic
Series
• Isoelectronic species with the greatest
charge in the nucleus will have the
largest ionization energy.
• For the isoelectronic series S2-, Cl-, and
Ar, Ar has the largest ionization energy
because it has the most protons in its
nucleus.
Ionization Energy
• Predict which species in each pair
has the higher ionization energy.
(a) Ca or As (b) K+ or Ca2+
(c) N or As
Successive Ionization Energies
• Successive ionization energies always
increase because of the increasing hold
the nucleus has on remaining electrons.
I1
Mg 738
Al 578
(all values in kJ/mol)
I2
1450
1817
I3
7734
2745
I4
10550
11600
• A much larger increase is seen when an
electron comes from a lower-energy
subshell.
Test Your Skill
• Which element, magnesium or sodium,
has the greater second ionization
energy?
Electron Affinity
• The electron affinity of an element is
the energy change that accompanies
the addition of an electron to a
gaseous atom to form an anion.
A(g) + e- → A-(g)
• Electron affinities are generally
favorable (exothermic) for elements
on the right side of the periodic table.
Electron Affinities
Alkali Metals – Group 1A (1)
• The reactivity of the Group 1A metals
increases down the group. Their
chemistry is dominated by the formation
of M+ ions.
2M(s) + H2O(l) → 2MOH(aq) + H2(g)
2M(s) + H2(g) → 2MH(s)
2M(s) + X2(g) → 2MX(s) X = F, Cl, Br, I
Alkali Metal Reactions with O2
• Only lithium reacts with O2 to give the
expected product, lithium oxide.
4Li(s) + O2(g) → 2Li2O(s)
• Sodium reacts mainly to yield sodium
peroxide.
2Na(s) + O2(g) → Na2O2(s)
• Potassium reacts to yield mixtures of the
oxide, peroxide, and superoxide.
K(s) + O2(g) → KO2(s)
Flame Colors of the 1A Elements
The Alkaline Earth Metals – Group 2A (2)
• The Group 2A metals are not as
reactive as the Group 1A metals.
Reactivity increases down the group,
and they all form M2+ ions.
• Magnesium alloys are useful in
aeronautical applications, where low
density and high strength are
important.
Flame Colors of 2A Elements
Calcium
Strontium
Barium
The Halogens – Group 7A (17)
• The halogens all exist as diatomic
molecules, but they are very reactive.
• The reactivity decreases as you go
down the group. Their chemistry is
dominated by the formation of X- ions.
• The interhalogens are compounds
formed from different halogens, like IF3
and BrCl.