The Periodic Law
History of the Periodic Table
*Antoine Lavoisier
(France, 1789)
• Earned reputation as
“father of chemistry”
• Established a common
naming system of compounds
and elements.
• First to organize elements
• Grouped them into four
categories:Gases, nonmetals, metals and
“earths” (elements that could not be chemically
separated at the time.)
History of the Periodic Table
Dmitri Mendeleev
(Russia, 1869)
1.Placed elements in groups in which
they shared similar properties which resulted in order of
increasing atomic mass with a few exceptions.
2. Ex: If placed solely by atomic mass, iodine was
not in group with chemically similar elements.
3. Left gaps for not-yet-discovered elements and
predicted their properties: gallium, germaniun &
scandium.
History of the Periodic Table
• Henry Mosely
• Moseley (1911) modified the table by
organizing elements in order of increasing
atomic numbers.
• Periodic Law: The phsical and chemical
properties of the elements are the
periodic functions of their atomic
numbers.
Glenn Seaborg
(UC Berkeley, 1944)
• Formed Actinide Series
just like that of the
lanthanides (#58-71)
Basics of the Periodic Table
periodic: a repeating pattern
table: an organized collection of information
period: horizontal row on the P.T.
•Designates e- energy levels
group or family: vertical column on the P.T.
Periodic Table: an arrangement of elements in order
of atomic number; elements with similar properties
appear at regular intervals (are in the same group)
Electron Structures of Atoms
nucleus
e-
1st energy level
e-
e-
1st Period: Hydrogen (#1)
Helium (#2)
e-
2nd energy level
e-
e-
e-
e-
e-
ee-
2nd Period: Lithium (#3)
eNeon (#10)
e-
e-
e-
e-
S block elements: Group 1 & 2
• Chemically reactive metals, group 1 more reactive
than group 2. Group Config: ns 1-2
• Alkali metals:silvery appearance, soft enough to cut
with a knife, not found in nature as free elements. H
shares e-config but not properties.
• Alkaline-earth metals:harder, denser, stronger, and
have a higher melting point than group 1. Too
reactive to be found uncombined in nature. He
shares e- config but not properties.
p block elements: Group 13-18
Includes all the three types of elements: metals, non
metals and metalloids.
Group Config: ns2 np 1-6
Includes Halogens: most reactive of the nonmetals.
React vigorously with most metals to form salts.
P block metals are generally harder & denser than s
block but softer & less dense than d block
metals.Found in nature solely as compounds except
for bismuth.
d block elements: Group 3-12
• Transition Elements: metals with typical properties;
good conductors, high luster.
• Less reactive than s block, many existing in nature as
free elements.
• Electrons added to the d sublevel of the preceding
energy level (n-1).
Group configuration: (n-1)d1-10ns 0-2
• Some deviations from orderly d sublevel filling occur
in group 4-11(s electrons jumping to d sublevel)
f-block elements
• F-block elements are wedged between groups
3 and 4 in the sixth and seventh period,
consisting of lanthanides and actinides
• Most elements are radioactive
• Trans Uranium elements are all synthetic
• Group Config: ns 0-2 (n-1) d 0-1 (n-2)f 1-14
atomic radius:
Covalent Radius for Covalently Bonded Atoms: half the distance between
the nuclei of two covalently bonded atoms

F-F bond length is 144 pm, so F covalent radius is 72 pm.

H-F bond length is 109 pm, so H covalent radius is 37 nm.
Atomic Radius for Elements like the Noble Gases

Ar atomic radius is 131 pm
Metallic Radius for Metals

Al metallic radius is 143 pm.
6.3
Trends in Atomic Size
• The atomic radius is one half of the distance
between the nuclei of two atoms of the same
element when the atoms are joined.
6.3
Trends in Atomic Size
6.3
Ionic Radii
Cation=positive ion, Anion=negative ion
• Forming a cation by losing electron(s) leads to
a decrease in atomic radius, a smaller
electron cloud.
• Forming an anion by electron(s) leads to an
increase in atomic radius, less pull from the
nucleus & there is more repulsion between
the greater number of electrons.
6.3
Cations
6.3
Anions
Across a period atoms become smaller. Down a group
atoms become larger.
Ionization Energy
Amount of energy required to remove an e from
a neutral atom in its gaseous state.
 First Ionization Energy
A(g)
 A+(g) + e Second Ionization Energy
+
A (g)
 A2+(g) + e Third Ionization Energy
2+
A (g)
 A3+(g) + e
Picture
IE and
trend
I.E. increases
across aof
period
decreases
down a group.
6.3
Trends in Ionization Energy
6.3
Trends in Ionization Energy
Electron Affinity




Amount of energy released when an e is
added to a gaseous atom in its neutral state.
First Electron Affinity
A(g) + e-  A-(g)
Second Electron Affinity
A-(g) + e-  A2-(g)
Third Electron Affinity
A2-(g) + e-  A3-(g)
Electronegativity
• A measure of the ability of an atom in a chemical
compound to attract electrons.
• Fluorine, the most electronegative element, is
arbitrarily assigned a value of 4.0. Values for other
elements are calculated in relation to this.
• Tend to increase across a period
• Tend to decrease down a group or remain about
the same.
• If an element does not form a compound, some
noble gases, will not have a value.
Trends in Electron Affinity and
Electronegativity
• Both electron affinity and electronegativity
increase from L to R across a period.
• Both electron affinity and electronegativity
decrease down a group.
Two Factors Used to Explain Trends
The principal energy level

All other factors being equal, increased n for
the orbitals in which electrons are found
means increased size of orbitals, which leads
to decreased attraction for electrons from the
nucleus.
Effective Nuclear Charge





Effective charge is the approximate net nuclear charge felt by
the highest energy electrons.
All other factors being equal, increased effective charge
means increased attraction for electrons, which leads to
decreased size of orbitals.
Effective charge depends upon two factors:
Total nuclear charge: # of protons (greater the total nuclear
charge, higher the attraction felt by electrons)
# of shielding electrons (e present in between the nucleus
and the valence shell electrons, the higher the number of
shielding electrons, the lesser is the effective nuclear charge)
SHIELDING:
• The net nuclear charge felt by an
outer electron is substantially lower
than the actual nuclear charge. the
outer electrons are shielded from the
full charge of the nucleus by the
inner electrons, which is called
shielding effect.
Explanation of Trends (1)
Explanation of Trends
Explanation of Trends
Bibliography
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http://www.homeoint.org/morrell/british/originidea.htm
http://www.chemsoc.org/viselements/pages/history.html
http://www.library.upenn.edu/etext/smith/d/dobereiner.html
http://www.library.upenn.edu/etext/smith/n/newlands.html
http://www.ulb.ac.be/sciences/cudec/ressources/Mendeleev.
gif
• http://intro.chem.okstate.edu/1314F00/Lecture/Chapter7/AT
RADIID.DIR_PICT0003.gif
• http://scidiv.bcc.ctc.edu/wv/4/0004-000-IE.GIF