This document discusses trends in properties of elements across the periodic table, including: 1. Atomic and ionic radii generally increase down a group and decrease across a period, with exceptions due to electronic configuration effects like the lanthanide contraction. 2. Ionization energies decrease down a group and increase across a period as it becomes easier to remove electrons from larger atoms with more diffuse orbitals. 3. Electronegativity increases across a period as nuclear charge increases, and decreases down a group as atomic size increases.

1. CHEM210 Inorganic
Chemistry
Main Group Descriptive Chemistry
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Recommended Reading: Shriver and Atkins, Inorganic Chemistry, 5th ed., Oxford
University Press, 2010.

2. Descriptive Chemistry
 Group trends: element properties, compounds of the
elements and applications.
 The elements of Group 14.
 Industrial preparation and applications of hydrogen.
 Synthesis and properties of the compounds of hydrogen.
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3. Chemistry
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Inorganic versus Organic Chemistry
Organic chemistry is in general much more
systematic than inorganic chemistry, the latter often
seeming to comprise a string of unrelated facts. This
is in part true since the subject matter of inorganic
chemistry is far more diverse and complicated and
the rules for chemical behavior are often less well
established.

4. Group Trends
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General Properties of the Elements
 Atomic radii
 Ionic radii
 Ionisation energies
 Electronegativity

5. Electronic Configuration
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When n = 1 1s
When n = 2 2s 2p
When n = 3 3s 3p 3d
When n = 4 4s 4p 4d 4f
When n = 5 5s 5p 5d 5f 5g
When n = 6 6s 6p 6d 6f 6g
When n = 7 7s 7p 7d 7f 7g
1s
2s
4s
5s
3p
2p
4p
3d
4d
3s

6. Effective Nuclear Charge
 As electrons are negatively charged they are attracted to nuclei
which are positively charged.
 The force of this attraction depends on the nuclear charge and the
distance between the nucleus and the electron.
 In a many electron atom, each electron is attracted to the nucleus
but also repelled by the other electrons.
 The attraction between the nucleus and the valence electrons is
particularly affected by the core electrons.
 The core electrons partially cancel the attraction force between the
nucleus and the valence electrons.
 The other valence electrons do not have a significant shielding
effect.
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7.  The actual nuclear charge, Z, is equal to the number of protons in
the nucleus.
 Because the core electrons shield the valence electrons from the full
charge, we will only see an effective nuclear charge: Zeff.
Zeff = Z - S
 Where S is the screening constant which is approximately equal to
the number of core electrons.
 In effect we are treating the valence electrons as though they were
moving in the net electric field created by the nucleus and core
electrons.
 As we move across a period the value of Z changes, but the number
of core electrons does not therefore the elements on the right of the
periodic table will experience a higher Zeff than those on the left.
Effective Nuclear Charge
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8. Atomic Radii
 Size is an important property of an atom as it affects many other
aspects of an atom.
 We tend to think of atoms as hard, spherical objects, but this is not
true. Atoms and ions do not have sharply defined edges.
 Atom size can therefore be defined in several ways:
Non-bonding atomic radius
If two non-reacting atoms which are moving freely collide with each
other they will not form a bond. The shortest distance between the
atoms during this collision gives the non-bonding atomic radius.
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9. Atomic Radii
Bonding atomic radius
When two atoms are bonded to each other there is an attractive
interaction between them. The two atoms will therefore approach more
closely than in a non-bonding interaction. The distance between the
two bonded nuclei can be referred to as the bonding atomic radius.
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For Example:
The distance separating two
iodine nuclei in I2 is 266 pm,
therefore the bonding atomic
radius of an iodine atom is 133
pm.

10. Atomic Radii
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Trends
 Increase down a group.
 Decrease from left to
right, within s and p
blocks.
 5d elements have similar
radii to 4d elements
(lanthanide contraction).
 Extra shells of electrons
added.

11. Lanthanide Contraction
 The 4d metals are considerably larger than the 3d metals this is
because an extra shell of electrons is being added.
 One would expect the same increase in radii between the 4d and 5d
transition metals, but this is not the case.
 This is attributed to the lanthanide contraction.
 The so-called lanthanide contraction is attributed to the presence of f
orbitals in the 5d metals which are not present in the 4d metals.
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12. Lanthanide Contraction
 The shielding ability of different electrons vary as follows: s > p > d >
f
 The shielding ability of f-electrons is quite poor due to the shape of
the
f-orbitals.
 Thus with increasing atomic number the effective nuclear charge
experienced by the 4f electrons increases.
 This results in a decrease in the atomic radii of elements that have f
electrons.
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The
shapes of
f orbitals

13. Ionic Radii
 Cations are much smaller than their corresponding atoms (the
whole outer shell is removed and the electron-electron repulsion is
reduced resulting in a considerably smaller radius than the atomic
radius.
 Anions are much larger than their corresponding atoms. When
electrons are added to an atom the electron-electron repulsions
increase, causing the electron cloud to spread. Correspondingly a
larger ionic radius is measured.
 Ionic radii increase down a group and decrease across a period.
(The same as atomic redii)
 How do we measure ionic radii?
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14. Ionic Radii

15. Ionisation Energies
 The ionisation energy of an element is the energy required to
remove an electron from a gas-phase atom.
 Several factors influence ionisation energy:
1. Size of the atom
2. Charge on the nucleus
3. The shielding ability of the inner electron shells
4. The type of electron involved (s, p, d…)
 In a small atom the electrons are tightly held, therefore ionisation
energy is high.
 In a large atom the outer electrons are less tightly held and
ionisation energy is lower.
 Decreases down a group and increases across a period.
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16. Ionisation Energies
 First ionisation energies generally increase with increasing atomic
number.
 As you descend a group the first ionisation energy decreases.
 The first ionisation energies of all transition elements are roughly
equal.
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17. Ionisation Energies
 Removal of successive electrons becomes more difficult: I1 < I2 < I3
 This trend exists because as each electron is removed the next
electron is being removed from a more positive ion, requiring more
energy.
 The removal of inner sphere electrons requires much more energy
than removing valence electrons.
Element I1 I2 I3 I4 I5
Na 495 4562
Mg 738 1451 7733
Al 578 1817 2745 11577
Si 786 1577 3232 4356 16091
P 1012 1907 2914 4964 6274
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18. Electronegativity
 Electronegativity is the power, with which an atom of an element is
able to attract electrons to itself when it is part of a compound.
 Increases across a period.
 Decreases down a group.
 There are exceptions to this trend.
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19. Oxidation States
 Electron configurations explain oxidation states to some extent.
 Full and half-full valence shells are more stable than partially filled
shells.
Example: What would be stable oxidation states for Ag and Mn?
 Heavier elements of the p block also form ions with oxidation
numbers 2 less than the group oxidation number. This due to the
inert pair effect.
 The stability of high oxidation states increases down a group.
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