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Chemical Periodicity
Chapter 6
Chemical Periodicity
• The periodic table is arranged in rows
according to increasing atomic number.
• Physical and chemical properties of the
elements can be explained by electron
configurations or positions in the periodic
table.
• Recall that the periodic table is organized in
blocks according to subshell being filled.
The Periodic Table
• The periodic table is organized into groups
or columns with similar electron
configurations.
– Na and K
– O and S
• For this reason, the elements in a group will
have similar chemical and physical
properties.
– Above examples and also halogens and inert
gases.
The Periodic Table
• Main group elements
– ‘A’ group elements (periodic table)
– Where is the last electron placed?
– Representative elements are the first three period of the
main group elements
• H to Ar
• Noble gases (part of the A group)
– Group VIIIA
– Where is the last electron? Outermost shell?
Demonstration: Noble gas electron configuration
The Periodic Table
• d-transition elements
– B group elements (identify on the table)
– How many series or periods?
– Where does the last electron go?
– The number of the d subshell being filled is one
less than the current shell number. Why?
The Periodic Table
• f-transition elements
– Termed as the inner-transition elements.
– Number of series or periods?
– Where is the last electron placed?
• Why is the f-transition elements termed as inner
transition elements?
Note: Electron buildup for f-transition elements can be complicated
(Appendix B)
Properties of the Elements
• The outermost electrons or the valence
electrons are largely responsible for the
chemical and physical properties of the
elements. The trends in the physical
properties that will be discussed can be
explained in a large degree to the valence
electrons.
Properties of the Elements
• The forces experienced by the valence electrons
can determine trends observed in many atomic
properties. What type of forces can the electrons
experience in an atom?
– Charge from the nucleus. A strong attraction between
the electron and the nucleus lowers the electron energy.
The electron is more tightly bound to the atom.
– Interactions among the electrons. There are repulsions
between electrons in the same orbital. Energy?
Factors that Influence the Energy of
the Electron
1. Shielding from the nucleus – the valence
electrons are shielded by inner electrons.
–
Zeff = Z -  (effective nuclear charge)
•
•
–
Z = nuclear charge,  = amount of shielding
Li example
Electrons with lower l quantum numbers
shield those with higher l quantum numbers.
Why?
– Electrons in the same l do not shield each
other significantly.
Zeff increases from left to right.
Zeff
Trend in Zeff within a Period/Row
7
6
5
4
3
2
1
0
Li
Be
B
C
N
O
F
Ne
IA
IIA IIIA IVA VA VIA VIIA VIIIA
Zeff increases from left to right within a period.
What consequence will this have on the valence electrons?
Energy of the electron?
Factors that Influence the Energy of
the Electron
2. Distance from the nucleus
–
Attraction between
opposite
charges


( q )( q )
F
d2
How will the decreasing d affect the energy of the
electron?
How will the energy of the electron be affected
by shell and subshell numbers?
Combining These Two Affects
2
eff
2
Z
En  
n
Discuss the dependencies:
Decreasing the value of n?
Increasing Zeff?
Factors that Influence the Energy
of the Electron
3. Interactions among electrons
–
Placing electrons in the same orbital creates
repulsions.
What affect will this have on the electron?
The chemical periodicity of the properties can be
discussed based on the preceding factors.
Atomic Radii
The size of an atom is roughly determined by
the ______________.
–
•
Typical sizes range from 0.3 to 3 Å.
The closer the outermost electron to the
nucleus the ______ the energy. 2
n
r
– Qualitative relationship
Z eff
Atomic Radii
• What does the relationship predict going
down a group/column?
• What does the relationship predict going
across a period?
Look at chart on page 241.
STM image of atoms
STM operation
Ionization Energy
• Minimum amount of energy required to
remove the most __________ electron from
a gaseous atom.
– 1st ionization energy (IE1)
• Mg(g) +738 kJ/mol Mg+(g) + e-
– 2nd ionization energy (IE2)
• Mg+(g) + 1451 kJ/mol  Mg2+(g) + e-
Why is IE2>IE1?
Ionization Energy
2
Z eff
En   2
n
How is energy of the electron related to the ionization energy?
Examine Figure 6-1 and Table 6-1. What are the trends? Why?
Ionization Energy
• Ionization energy
down a column?
• Ionization energy
across a period?
– Na through Ar
– Explain the trend
This is a periodicity
pattern. Holds well for
Group A.
Ionization Energies – Electron
Configurations
• Trends for IE’s in the group of the
representative elements can partially be
explained by electron configurations.
– Why?
– Group IA, the alkali metals, has the lowest IE’s.
• 1st and 2nd ionization energies for Na
– Group IIA trends
• 1st, 2nd, and 3rd ionization energies Why is there a
large relative increase in the 3rd ionization energy?
Ionization Energies - Reactivity
• Elements with low or high ionization
energies form _______________.
– Which groups in A?
• Elements with intermediate values of IE
___________ upon reacting.
– Which groups in A?
Electronegativity
• A measure of the relative tendency of an
atom to _________ upon combining
chemically with another atom.
– Elements with high electronegativities?
– Elements with low electronegativities?
– Elements with intermediate electronegativities?
Electronegativity
• Representative
elements
– Increases from left to
right Why?
– Decreases going
down a group Why?
Type of bonding that
will occur in a
compound
Electron Affinity
• The amount of energy absorbed when an
electron is added to an isolated gaseous
atom to form an ion with a 1- charge.
– Measure of an atom’s ability to form a negative
ion.
• Cl(g) + e-  Cl-(g) + 349 kJ (releases energy)
• Mg(g) + e- + 231 kJ  Mg-(g) (takes energy)