Download Chapter 5 – The Periodic Law

Chapter 5 – The Periodic Law
I.
History of the Periodic Table
A.
Mendeleev and Chemical Periodicity
1.
Dmitri Mendeleev discovered meaningful patterns of
properties among the approximately 63 known elements in
1869
2.
B.
a.
The elements were listed in the order of increasing
atomic mass
b.
Mendeleev noticed that similar chemical and physical
properties recur periodically when the elements were
arranged in this way
Known as the “Father of the Periodic Table”
a.
He insisted that elements with similar characteristics
be listed in the same groups (vertical columns)…for
this reason, he had to leave several blank spaces in
his periodic table
b.
Left empty spaces for elements when he couldn’t “fit”
known elements into the next column
c.
Correctly predicted the physical and chemical
properties of the missing elements based on the
trends he saw in the groups…they were all
discovered by 1886
Moseley and the Periodic Law
1.
In 1913, two years after Rutherford proposed the nuclear
model of the atom, Henry Moseley developed the concept
of atomic numbers
a.
REMEMBER…the atomic number of an element
represents
II.
b.
The periodic chart was modified by Moseley to be
arranged
c.
This explained why in the previous version of the
periodic table the position of some elements had to be
altered in order to place them in the correct group
of elements
The Modern Periodic Table
A.
B.
C.
Groups to MEMORIZE…
1.
Group 1 –
2.
Group 2 –
3.
Group 7 –
4.
Group 8 –
Other groupings to be familiar with…
1.
Main-group (“representative”) elements –
2.
Transition elements –
3.
Lanthanides –
4.
Actinides –
Electron Configurations and the Periodic Table
1.
The periodic table can be divided into 4 “blocks”: s, p, d, & f
2.
a.
s block –
b.
p block –
c.
d block –
d.
f block –
The periodic table can be used as a general guide to write
electron configurations; there are some exceptions in the d
block
3.
For the “s” and “p” blocks (the main-group elements), all
elements in the same group will
III.
Electron Configuration and Periodic Properties
A.
Effective Nuclear Charge, Zeff
1.
The effective nuclear charge increases from left to right in a
period and stays constant from top to bottom in a group
2.
B.
The greater the effective nuclear charge is, the greater the
attractive forces between the nucleus and its electrons will
be
Atomic Radii
1.
Atomic radius (AR) –
2.
The atomic radius decreases from left to right in a period and
increases from top to bottom in a group
D.
Ionization Energy (IE) –
1.
The first ionization energy (related to removing the first
electron from the outer shell of an atom or ion) increases
from left to right in a period and decreases from top to
bottom in a group
2.
As you remove more and more electrons from an atom/ion,
the effective nuclear charge will increase and, therefore, it
will be harder to remove the next electron…second, third,
etc., ionization energies are always larger than the previous
ionization energy
E.
Electron Affinity (EA) –
1.
This energy change is called the electron affininty because it
measures the attraction, or affinity, of the atom for the added
electron
2.
In general, the electron affinity values decrease (become
more negative) from left to right in a period
3.
Electron affinities do not change greatly as we move down a
group
F.
Ionic Radii
1.
Cations are formed when an atom loses one or more
electrons…
2.
Anions are formed when an atom gains one or more
electrons…
3.
Both cations and anions decrease in size from left to right
across a period and increase in size from top to bottom in a
group
G.
Electronegativity (EN)–
1.
Electronegativity values increase from left to right across a
period (although there are exceptions) and either decrease
or remain about the same from top to bottom in a group
2.
The difference in the electronegativity values between two
atoms will determine whether the bond formed will be ionic
or covalent
H.
Summing It Up