Section 7
Reading assignment: 7.1-7.5
As you read ask yourself:
What is meant by the expression “effective nuclear charge”?
How can you use this concept to explain the trends in atomic
radius in the periodic table?
How can you use effective nuclear charge and electron
configuration to predict what happens to the radius when ions are
formed?
Why is the second ionization energy always greater than the first
ionization energy?
What types of elements have the highest ionization energies and
what types have the lowest ionization energies?
Why can electron affinity values be both positive and negative?
7.1 Development of the Periodic Table
The periodic table is the most significant tool that chemists use for organizing
and recalling chemical facts.
arrangement reflects trends in chemical and physical properties
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Section 7
7.2 Effective Nuclear Charge
Li
Na
7.2 Effective Nuclear Charge
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Section 7
7.3 Sizes of atoms and ions
7.3 Sizes of atoms : trends
higher Zeff pulls the valence
electrons toward the
nucleus
increases
valence
electrons have
higher n values,
higher shells
increases
radius
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3
Section 7
7.3 Periodic trends in Ionic Radii
Ionic radii relative to metallic (or covalent) radii (in Å)
Brown, LeMay, Bursten & Murphy “Chemistry The Central Science” 11th Ed., Pearson 2009, Fig. 7.8, p. 263
7.3 Periodic trends in Ionic Radii
isoelectronic species
ions with same charge:
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Section 7
7.4 Ionization Energy
Note sharp increase in ionization energy when a core electron is removed.
Ionization energy generally increases across a period.
Zeff increases, making it more difficult to remove an electron
Two exceptions are removing the first p electron and removing the fourth p
electron
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Section 7
7.5 Electron Affinities
smallest atoms
Summary:
easy to form
negative ions
easy to form
positive ions
largest atoms
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