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The History of
the Modern
Periodic Table
Modified from www.thecatalyst.com
During the nineteenth century, chemists began
to categorize the elements according to
similarities in their physical and chemical
properties.
The end result of these studies was our
modern periodic table!
Johann Dobereiner 1780 - 1849
In 1829, he classified some elements into groups of
three, which he called triads.
Model of triads
The elements in a triad had
similar chemical and physical
properties.
John Newlands 1838 - 1898
In 1863, he suggested that elements be arranged in
“octaves” because he noticed that certain properties
repeated every 8th element when placed in order of
increasing atomic mass.
Law of Octaves
John Newlands
Newland’s law of octaves failed beyond the element
calcium.
WHY?
Would his law of octaves work today with the first
20 elements?
Elements known at this time
Dmitri Mendeleev 1834 - 1907
In 1869 he published a table of the
elements organized by increasing atomic
mass.
Lothar Meyer 1830 - 1895
At the same time, Lothar
Meyer published his own table
of the elements organized by
increasing atomic mass.
Mendeleev vs. Meyer
• Both Mendeleev and Meyer arranged the elements
in order of increasing atomic mass.
• Both left vacant spaces where unknown elements
should fit.
So…
……why is Mendeleev called the “father of
the modern periodic table” and not Meyer?
And the Winner is…
• Mendeleev stated that if the atomic weight of an
element caused it to be placed in the wrong group,
then the weight must be wrong.
•Mendeleev was so confident in his table that he
used it to predict the physical properties of
three elements that were yet unknown.
•Mendeleev’s predictions for Sc, Ga, and Ge turned
out to be amazingly close to the actual values!!
After the discovery of these unknown elements,
Mendeleev’s table was generally accepted.
Mendeleev’s 15 minutes of Fame?
In spite of Mendeleev’s great achievement, problems
arose when new elements were discovered and more
accurate atomic weights determined.
Can you identify the problematic elements?
Ar and K
Co and Ni
Te and I
Th and Pa
Henry Moseley 1887 - 1915
In 1913, through his work with X-rays, he determined
the actual nuclear charge (atomic number) of the
elements.
He then arranged the
elements in order of
increasing atomic
number (nuclear
charge).
Henry Moseley
“There is in the atom a fundamental quantity
which increases by regular steps as we pass from
each element to the next. This quantity can only
be the charge on the central positive nucleus.”
Moseley’s research was halted when the British government
sent him to serve as a foot soldier in WWI. He was killed
in the fighting , at the age of 28. Because of this loss, the
British government later restricted its scientists to
noncombatant duties during WWII.
Glenn T. Seaborg 1912 - 1999
After co-discovering 10 new elements, he moved 14
elements out of the main body of the periodic table
to their current location below the Lanthanide series.
These became known as the Actinide series.
Glenn T. Seaborg
He is the only person to have an element named after
him while still alive.
"This is the greatest honor ever
bestowed upon me - even better, I
think, than winning the Nobel Prize."
Periodic Table
Geography
Periods
The horizontal rows of the periodic table are called
PERIODS.
Groups
The vertical columns of the periodic table are called
GROUPS, or FAMILIES.
Periodic Law
When elements are arranged in order of increasing
atomic number, there is a periodic pattern in their
physical and chemical properties.
In other words…
The elements in any group of the periodic table
have similar physical and chemical properties!
Alkali Metals
All of the elements in the first Group are
considered part of the Alkali Metals Family.
Alkaline Earth Metals
All of the elements in the second Group are
considered part of the Alkaline Earth Metal
Family.
Transition Metals
The Transition metals are indicated by the
letter B in the Group Number. All of these
elements have a partially filled d-orbital.
Inner-Transition Metals
These elements are also called the rare
earth elements, as they are not commonly
found in nature. In fact, many of them were
“created” in particle accelerators.
Halogens
All of the elements in Group VII A are
considered part of the Halogen Family.
Noble Gases
All of the elements in Group VIII A are considered
part of the Noble Gas Family, as they are rather
inert (i.e. non-reactive) under most conditions.
The periodic table is the most important
tool in the chemist’s toolbox!
Periodic Trends & Properties
We will be looking at the following periodic trends in
this chapter:
Nuclear Charge
Shielding Effect
Atomic size (radius)
Ionic size (radius)
Ionization energy
Electronegativity
Nuclear Charge
•The nuclear charge increases as you move
down a group in the periodic table.
The Reason:
• As Moseley discovered, the number of protons
in the nucleus is what determines the nuclear
charge.
• The more protons in the nucleus, the larger or
stronger the nuclear charge.
Nuclear Charge
• Nuclear Charge increases from left to
right across the periodic table.
The Reason:
• The number of protons in the nucleus increases by one as
you move from left to right across the table.
• Increasing the number of protons in the nucleus results in
an increase in nuclear charge.
Shielding Effect
Within any group of elements, the shielding effect
increases as you move down the column.
The reason:
•The elements located in higher periods have
more electrons, and therefore require more
[higher] energy levels.
•More energy levels translates into more inner
electrons and thus more shielding.
Shielding Effect
Within any period, the shielding effect decreases
from left to right across the table.
The Reason:
• Across the period, the nuclear charge increases
but the number of energy levels remains the same.
• The increasing nuclear force draws the outermost
electrons ever closer to the nucleus…
Atomic Radius
Within any group of elements, the atomic
radius will increase from top to bottom.
The reason:
•The elements located in higher periods
have more electrons, and therefore
require more [higher] energy levels.
•More energy levels translates into a
larger radius.
Atomic Radius
Within a period, the atomic radius decreases
from left to right across the periodic table.
The reason:
•Elements within the same period have the same number of
energy levels and therefore the same shielding effect.
•As you move toward the right, the nuclear charge increases.
•With more nuclear charge, and no increase in shielding, the
outermost electrons are drawn closer to the nucleus, and
hence the atomic radius decreases.
Ionic Radius of Cations
Cations are ALWAYS SMALLER in size than the
neutral atom from which they originate.
The Reason:
• The nuclear charge does not change, but the atom
loses 1 or more electrons.
• Each valence electron experiences more electrostatic
attraction toward the nucleus.
• As a result, the ion is smaller than the original atom.
Lithium
Lithium ion, Li+1
Ionic Radius of Anions
Anions are ALWAYS LARGER in size than the neutral
atom from which they originate.
The Reason:
•The atom gains extra electrons, but the number of
protons attracting those electrons remains the same.
•Each valence electron experiences less electrostatic
attraction toward the nucleus.
•As a result, the ion is larger than the original atom.
Oxygen
Oxygen ion, O-2
Ionization Energies Within Periods
Group 1 Metals
• Have 1 valence electron. It requires a small amount of energy
to remove that single electron. (Low 1st ionization energy)
+ 520 kJ/mol
• In order to remove a second electron, it requires removing an
electron from a stable (noble) configuration. (High 2nd
ionization energy)
+ 7,297 kJ/mol
Ionization Energies Within Periods
Group 2 Metals
•Have 2 valence electrons. It requires a small amount of
energy to remove the first and second electrons.
+ 900 kJ/mol
+ 1,757 kJ/mol
•In order to remove a third electron, it requires removing an
electron from a stable (noble) configuration.
+ 14,840 kJ/mol
Trends in Ionization Energy
Ionization Energy increases as you move
left to right across the periodic table.
The Reason:
•As you move from left to right within a period:
•Atomic radius decreases
•Shielding Effect does
not change
•Nuclear charge increases
•As a result, valence electrons are drawn closer to the
nucleus and held more tightly. This means it requires
more energy to remove an electron.
Trends in Ionization Energy
Ionization Energy decreases as you move top to
bottom on the periodic table.
The Reason:
•As you move from top to bottom in a group:
•Atomic radius increases
•Nuclear charge increases
•Shielding Effect increases
•As a result, valence electrons are farther from
the nucleus and more loosely held. This means
it requires less energy to remove an electron.
Electronegativity of Metals
Metal Atoms
• Compared to non-metals, metals have:
• Larger atomic radii
• Weaker nuclear charge
• As a result, metal atoms tend to lose their valence electrons
when reaching a stable configuration.
•Thus, metals tend to have lower electronegativities.
H
2.1
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Na
0.9
Mg
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
K
0.8
Ca
1.0
Ga
1.6
Ge
1.8
As
2.0
Se
2.4
Br
2.8
Electronegativity of Non-Metals
Non-Metal Atoms
• Compared with metal atoms, non-metals have
•Smaller atomic radii
•Stronger nuclear charge
• As a result non-metals tend to gain valence electrons
when reaching a stable configuration.
•Thus, non-metals tend to have higher electronegativities.
H
2.1
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Na
0.9
Mg
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
K
0.8
Ca
1.0
Ga
1.6
Ge
1.8
As
2.0
Se
2.4
Br
2.8