CHAPTER 7 Periodic Properties of the
Elements
 Development of the Periodic Table
 Electron Shells and the Size of the Atoms
 Ionization Energy
 Electron Affinity
 Others (read on your own)
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Development of the Periodic Table
Mendeleev developed periodic table to group
elements in terms of chemical properties.
 Atomic radii increase down a column and
decrease across a row.
 Ionization energy, electron affinity varies in
uniform manner.
 Alkali metals develop +1 charge, alkaline earth
metals + 2
 Nonmetals usually develop negative charge
(1 for halides, 2 for group 6A, etc.)
 Blank spots where elements should be were
observed.
 Discovery of elements with correct properties
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Electron Shells and the Sizes of
Atoms
Periodic trends in atomic and ionic radii (and
other things) also agree with the quantum
mechanical configuration.
 Atomic radii actually decrease across a row in
the periodic table. Due to an increase in the
effective nuclear charge.
 If positively charged the radius decreases
while if the charge is negatively the radius
increases (relative to the atom).
 When substances have the same number of
electrons (isoelectronic) the radius will depend
upon which has the largest number of protons.
E.g. Predict which of the following substances
has the largest radius: P3, S2, Cl, Ar, K+,
Ca2+.
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IONIZATION ENERGY
Ionization energy, Ei: minimum energy
required to remove an electron from the
ground state of atom (molecule) in the gas
phase. M(g) + h  M+ + e.
 Ei related to electron configuration. Higher
energies = stable ground states.
 Sign of the ionization energy is always
positive, i.e. it requires energy for ionization to
occur.
 The ionization energy is inversely proportional
to the radius and directly related to Zeff.
 Exceptions to trend:
a. B, Al, Ga, etc.: their ionization energies are
slightly less than the ionization energy of
the element preceding them in their period.
i. Before ionization ns2np1.
ii. After ionization is ns2. Higher energy 
smaller radius.
b. Group 6A elements.
i. Before ionization ns2np4.
ii. After ionization ns2np3 where each p
electron in different orbital (Hund’s rule).
Electron-electron repulsion by two electrons in
same orbital increases the energy (lowers EI).
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HIGHER IONIZATION ENERGIES
 The energies for the subsequent loss of more
electrons are increasingly higher. For the
second ionization reaction written as
M+(g) + h  M2+ + e Ei2.
 Large increases in the ionization energies
vary in a zig-zag way across the periodic
table.
n
0
Ei,n
NA
Na
Final.
Config.
Ei,n
1s22s22p63s1
NA
Mg
Final
Config.
1s22s22p63s2
Ei,n
NA
Fi
C
1s
1
496
1s22s22p6
738
1s22s22p63s1
578
1s
2
4562
1s22s22p5
1451
1s22s22p6
1817
1s
3
9543
1s22s22p4
7733
1s22s22p5
2745
1s
4
13353
1s22s22p3
10540
1s22s22p4
11575
1s
 States with higher ionization energies have:
1s22s22p6 (stable).
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ELECTRON AFFINITY
Electron Affinity, Eea, is the energy change
that occurs when an isolated atom in the
gas phase gains an electron.
E.g. Cl + e  Cl Eea = 348.6 kJ/mol
 Energy is often released during the process.
 Magnitude of released energy indicates the
tendency of the atom to gain an electron.
a. From the data in the table the halogens
clearly have a strong tendency to become
negatively charged
b. Inert gases and group I & II elements have
a very small EA.
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