Chapter 7 periodic trends
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Development of the Periodic
Table
• 1869 Mendeleev &
Meyer publish nearly
identical
classifications of
elements.
Mendeleev
Meyer
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•Insisted that elements with similar
characteristics be groups into
families.
•He left blanks spaces for unknown
elements and predicted their
physical properties.
•In 1913 Mosley developed the
concept of atomic numbers that we
use today to classify elements.
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Periodicity
• The valence electron structure of atoms
can be used to explain various properties
of atoms.
• In general properties correlate down a
group and across a period.
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Periodicity Vocabulary
• Valence Electrons: Outermost electrons.
Requires less energy to remove due to
increased distance from the nucleus and
positive protons.
• Core Electrons : An inner electron in an atom.
Harder to remove due to strong bond between
positive nucleus
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7.2 Effective Nuclear Charge
(Zeff)
• Tells us how strongly an electron is attached to
the nucleus.
• The force of attraction of an electron (negatively
charged) INCREASES as nuclear charge
(positive due to protons in the nucleus)
INCREASES
• The force of attraction of an electron
DECREASES as the electron moves away from
the positively charged nucleus.
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Zeff =
(number of p+ in the nucleus) – (average number of e-)
Zeff = (Z) – (S)
• The effective nuclear charge experienced by
outer electrons is determined by the difference
between the charge on the nucleus and the
charge of the core electrons. Pg 322
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Periodic Trends
H
Li Mg
Na Mg
atom size
1st ionization energy
electron affinity
electronegativity
He
B C N O
Al Si P S
F Ne
Cl Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I
Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac Rf Db Sg Bh Hs Mt Ds
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
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Ionization Energy
• Energy required to
remove one electron
from an atom.
• The greater the
propensity for an
atom to hold onto an
electron the higher
the ionization energy
required to remove
that electron.
Note the size in electron
clouds
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Ionization Energy cont.
First Ionization Energy (I1): Energy required
to remove the first electron from an atom
in the ground state (no charge)
Second Ionization Energy(I2):Energy
required to remove the second electron
from an atom (X+)
Table 8.2 page 331
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NOTE:
You will see a large SPIKE in
energy when you begin to remove
core electrons.
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Trends in Ionization
• One can perform multiple ionizations
Al (g)  Al+ +e-
Al+ (g)  Al2+ +eAl2+ (g)  Al3+ +e-
I1 = 580 kJ/mol
I2 = 1815 kJ/mol
I3 = 2740 kJ/mol
I1 < I 2 < I 3
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Periodic Table - Trends
ionization energy
-
Ionization energy increases across a period
And decreases down a group
+
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Order the indicated three elements according to the
ease with which each is likely to lose its third
electron.
A vertical transition is greater than a horizontal
transition
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Removing Valence and Core
Electrons
1. Na (g)

[Ne]3s1
2. Na+ (g)
[Ne]

Na+ (g) + e-
I1 = 495 kJ/mol
[Ne] (removing valence e- )
Na2+ (g) + e-
I2 = 4560 kJ/mol
1s22s22p5 (removing core electrons)
* It takes significantly more energy to remove core electrons
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Electron Affinity
• The energy change (ΔE) associated with the
addition of an electron.
(affinity for chocolate)
X(g) + e  X(g)
Example:
Cl (g) + e-  Cl- (g)
∆E = -349 kJ/mole
(negative E thus energy is emitted to add an electron)18
Electron Affinity
A negative ΔE indicates a strong attraction
between atom and the added electron.
The stronger the attraction the more
energy will be released.
Cl + e-  Cl- ΔE = -349kJ/mol
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Periodic Table - Trends
Electron Affinity
+
More negative
ΔE
More positive ΔE
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Which of the indicated three elements has the
least favorable Eea, and which has the most
favorable Eea?
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Atomic Radii
• Allows us to determine the
bond length between two
covalently bonded atoms.
• Ex:
the Br-Br bond distance of Br2
is 228 ppm therefore the
atomic radius of Br is
228/2 = 114 ppm
Pg 324 fig. 8.5
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Periodic Table - Trends
Atomic Radii
-
+
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Shielding
• moving down the periodic table the
number of Ve stays the same but the
number of core electrons increases.
Atomic radius increase down a group
because the increase in core electrons
SHIELDS the valence electrons
from the nucleus allowing them to pull
farther away.
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The Why
• There is a correlation between atomic radii
and the principle quantum number n.
• As n increases atomic radii increases due
to the e- moving farther and farther away
from the nucleus, pulling on the e- less
and less and allowing them to spread out
and be less dense.
n=2
n=5
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Isoelectric Ions
Please read this section it is confusing and reading
will help!
• These are ions that have the same number of
electrons.
– Example: O22- , F-, Na+, Mg2+, Al3+ all have 10 electrons
so we can’t use our repulsive and attractive forces the
figure out their size so we need to look at their effective
nuclear charge (Zeff).
• In isoelectric series size (radii) of the atom
DECREASES with an INCREASE in nuclear
charge as the electrons are attracted more
strongly to the nucleus.
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Radii of Ions
• Size of ions is based on
the distance between the
ions in the ionic
compound
• Would you expect the
cations of these
elements to be larger or
smaller than the ground
state atom?
• Pg 327 fig. 8.9
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Radii of Ions
An anion occurs when we add an electron to an atom making it
more negative
Anions of an atom are larger than its parent.
A cation occurs when we remove an electron from an atom
making it more positive
A cation of an atom is smaller than its parent.
For atoms with the same charge (all –2 atoms or +2 atoms)
size of the atom INCREASES as you move down a group
and INCREASES across a group from left to right. (Figure
7.6)
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Electronegativity
• How strongly the nucleus attracts the
electrons of OTHER atoms in a bond.
EN increases
Decreases
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Homework
• Chang pg 348
Ionization/Affinity: 21,24,49,52,55,59,61,64
Radii: 27, 32-37,89
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