1. Periodic
Relationships
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2. PERIODIC TRENDS IN PROPERTIES
OF ELEMENTS
 There are many observable patterns in the physical and chemical
properties of elements as we descend in a group or move across a
period in the Periodic Table.
 We will rationalize observed trends in
 Sizes of atoms and ions.
 Ionization energy.
 Electron affinity.
The bonding atomic
radius is defined as
one-half of the distance
between covalently
bonded nuclei.
The Size of an Atom

3. Trend in Atomic Radius
 Different methods for measuring the radius of an atom, and
they give slightly different trends
 Van der Waals radius = Nonbonding
 Covalent radius = Bonding radius
 Atomic radius is an average radius
of an atom based on measuring large
numbers of elements and compounds.

4.  Valence shell farther from nucleus
 Effective nuclear charge fairly close
 Atomic Radius Decreases across period (left to right)
The number of energy levels increases as you move down a group as
the number of electrons increases. Each subsequent energy level is
further from the nucleus than the last. Therefore, the atomic radius
increases as the group and energy levels increase.
 Atomic Radius Increases down group
Atomic Radius
adding electrons to same valence shell
effective nuclear charge increases
valence shell held closer
- As you go across a period, electrons are added to the same energy
level. At the same time, protons are being added to the nucleus. The
concentration of more protons in the nucleus creates a "higher
effective nuclear charge." In other words, there is a stronger force
of attraction pulling the electrons closer to the nucleus resulting in a
smaller atomic radius.

5. Atomic Radius

6. Sizes of Ions
• Ionic size depends upon
– The nuclear charge.
– The number of electrons.
– The orbital in which
electrons reside.

7. Sizes of Ions
• Cat ions are smaller than their parent atoms:
– The outermost electron is removed and repulsions between
electrons are reduced.
• Anions are larger than their parent atoms
– Electrons are added and repulsions between electrons are
increased.
• In an Isoelectronic series, ions have the same number
of electrons.
• Ionic size decreases with an increasing nuclear charge.

8. Ionization Energy
The ionization energy is the amount of energy
required to remove an electron from the ground state of
a gaseous atom or ion.
– The first ionization energy is that energy required to
remove the first electron.
– The second ionization energy is that energy
required to remove the second electron, etc.
• It requires more energy to remove each successive electron.
• When all valence electrons have been removed, the ionization
energy takes a quantum leap.

9. Ionization Energy

10. First Ionization Energies
 Larger the effective nuclear charge on the electron,
the more energy it takes to remove it
 The farther the most probable distance the electron is
from the nucleus, the less energy it takes to remove it
 1st IE decreases down the group
 valence electron farther from nucleus
 1st IE generally increases across the period
 effective nuclear charge increases

11. First Ionization Energies

12. First Ionization Energies

13. Irregularities in the Trend
 Ionization Energy generally increases from left to
right across a Period
 except from 2A to 3A, 5A to 6A
Be
1s 2s 2p
B
1s 2s 2p
N
1s 2s 2p
O
1s 2s 2p
Which is easier to remove an
electron from B or Be? Why?
Which is easier to remove an
electron from N or O? Why?

14. Irregularities in the
First Ionization Energy Trends
Be
1s 2s 2p
B
1s 2s 2p
Be+
1s 2s 2p
To ionize Be you must break up a full sublevel, cost extra energy
B+
1s 2s 2p
When you ionize B you get a full sublevel, costs less energy

15. Irregularities in the
First Ionization Energy Trends
To ionize N you must break up a half-full sublevel, cost extra energy
N+
1s 2s 2p
O
1s 2s 2p
N
1s 2s 2p
O+
1s 2s 2p
When you ionize O you get a half-full sublevel, costs less energy

16. Trends in Successive
Ionization Energies
 Removal of each successive electron
costs more energy
– shrinkage in size due to having more
protons than electrons
– outer electrons closer to the nucleus,
therefore harder to remove
 Regular increase in energy for each
successive valence electron
 Rarge increase in energy when start
removing core electrons

17. Ionization Energies (kJ/mol)

18. Electron Affinity
Electron affinity is the energy change accompanying
the addition of an electron to a gaseous atom:
Cl + e− Cl−
1) As you move down a group, electron affinity decreases.
2) As you move across a period, electron affinity increases.

19. Electron Affinity
 The first occurs between
Groups IA and IIA.
– The added electron must go in a
p orbital, not an s orbital.
– The electron is farther from the
nucleus and feels repulsion from
the s electrons.
 The second discontinuity
occurs between Groups IVA
and VA.
– Group VA has no empty
orbitals.
– The extra electron must go into
an already occupied orbital,
creating repulsion.

20. Oxidation States
 A way of keeping track of the electrons.
 Not necessarily true of what is in nature, but it works.
 need the rules for assigning .
The oxidation state of elements in their standard states is zero.
Oxidation state for monatomic ions are the same as their charge.
Oxygen is assigned an oxidation state of -2 in its covalent
compounds except as a peroxide.
In compounds with nonmetals hydrogen is assigned the oxidation
state +1.
In its compounds fluorine is always –1.
The sum of the oxidation states must be zero in compounds or
equal the charge of the ion.

21. Oxidation States
1. The oxidation state of any element such as Fe, H2, O2, P4,
S8 is zero (0).
2. The oxidation state of oxygen in its compounds is -2, except
for peroxides like H2O2, and Na2O2, in which the oxidation
state for O is -1.
3. The oxidation state of hydrogen is +1 in its compounds,
except for metal hydrides, such as NaH, LiH, etc., in which
the oxidation state for H is -1.
4. The oxidation states of other elements are then assigned to
make the algebraic sum of the oxidation states equal to the
net charge on the molecule or ion.
5. The following elements usually have the same oxidation
states in their compounds:+1 for alkali metals - Li, Na, K, Rb,
Cs;
6. +2 for alkaline earth metals - Be, Mg, Ca, Sr, Ba;
7. -1 for halogens except when they form compounds with
oxygen or one another;

22. Element
Oxidation
state
Compound
or ion
Fe +2 Fe2+ Fe = Fe2+ + 2 e-
+3 Fe3+ Fe2+ = Fe3++ e-
Zn 0 Zn Zn is reducing agent
+2 Zn2+
O -1 H2O2 H2O2 = O2 + H2O
0 O2
-2 H2O
Oxidation States

23. The End
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