Transcript PowerPoint

Module 3.03
Periodic Trends
Periodic Table
To Review:
• Periods – rows of the periodic table
• Groups – columns of the period table
Periodic Table
Trends
• Trends are useful in making predictions
• Periodic table is arranged based on the
properties of the elements
• Reading across each period and down each
group you will see repeated trends in
properties
Effective Nuclear Charge
Effective Nuclear Charge – the charge (from
the nucleus) felt by the valence electrons
after the number of shielding electrons are
taken into account.
Subtract the number of shielding (core)
electrons from the total nuclear charge
(number of protons) gives the effective
nuclear charge.
Effective Nuclear Charge
Example:
Nitrogen
7 protons (atomic number of 7)
-2 core electrons (1s2 electrons)
-------5 is the effective nuclear charge
Each of the valence electrons in nitrogen feels
an effective nuclear charge of about +5.
Effective Nuclear Charge
Across a Period:
The effective nuclear charge felt by an
atom’s valence electrons increases by one
for each element from left to right in a
period
Down a Group:
The effective nuclear charge felt by an
atom’s valence electrons stays constant for
each element going down a group
Atomic Radius
Atomic Radius – half the distance between
the centers of two atoms that are bonded
together.
Atomic Radius
Across a Period:
There is a gradual decrease in atomic
radii from left to right across a period.
Down a Group:
There is a general increase in atomic radii
going down each group on the periodic
table.
Atomic Radius
Ionization Energy
Ionization energy – energy required to remove
on electron from an element, resulting in a
positive ion.
Elements with a lower effective nuclear charge
felt by their electrons will give up an electron
easier than other elements.
Ionization Energy
Across a Period:
Ionization energy has a general increase
for elements across a period from left to
right.
Down a Group:
Ionization energy of elements decreases
going down a group because the atomic
radius of the atoms increases.
Ionic radius
Ionic radius – ½ the diameter of an ion.
Note: Metals naturally form cations (positive
ions) by losing one or more electrons.
Nonmetals naturally form anions (negative
ions) by gaining one or more electrons.
Atomic Radius vs Ionic Radius
Ionic Radius
Across a Period:
Within each period, the metals at the left form
cations and the nonmetals at the right form
anions. There is a decrease in the ionic radii of the
cations from left to right and a decrease in the
ionic radii of the anions from left to right.
Down a Group:
Ionic radii increase down a group, following the
same trend as atomic radii.
Electronegativity
Electronegativity – measure of the attraction of
an atom for the electrons in a chemical bond.
Electronegativity
Across a Period:
Electronegativity increases from left to right
across a period because of the effective nuclear
charge.
Down a Group:
Electronegativity decreases down a group as a
result of a greater atomic radius.
Module 3.03
Periodic Trends