Download History of the Periodic Table Chapter 6.1 Who developed the first

24-Oct-11
Who developed the first periodic
table?
History of the Periodic Table
Chapter 6.1
• Dobereiner (Germany, 1817) noticed that
atomic mass of Sr was closely related to the
atomic mass of Ca and Ba.
• He called this grouping a triad.
• Several other triads were created based on
his system but eventually there were not
enough triads to make the system useful.
Who developed the modern periodic
table?
Why was Mendeleev credited with
the periodic table?
• Dmetri Mendeleev (Russia)
and Lothar Meyer (Germany)
both published almost
identical systems of
classifying elements.
• Arranged periodic table based
on their physical and
chemical properties and
atomic mass.
• He insisted that elements be grouped
together by properties and thus by family.
• The gaps in his periodic table were actually
elements yet to be discovered!!
• He was able to predict the properties of
these undiscovered elements.
Prediction of Germanium’s
Properties
Property
Mendeleev’s Prediction
Observed Properties
Atomic Weight
72
72.59
Density
5.5
5.35
Specific Heat
0.305
0.309
Melting point
High
947C
Color
Dark gray
Grayish white
Formula of oxide
XO2
GeO2
Density of oxide
4.7
4.70
Formula of Chloride
XCl4
GeCl4
Boiling point of chloride A little under 100
Periodic Law
• Mendeleev stated that the physical and
chemical properties of elements vary
periodically with increasing atomic mass.
84C
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How were the table and atomic numbers proven?
• Two years after Rutherford created his atomic
model, Robert Moseley developed the concept of
atomic numbers.
• He bombarded elements with high energy
electrons and measured the frequency of x-rays
given off.
• Then he arranged the elements in order according
to the frequency given off. It was then concluded
that as the atomic mass of the elements increased
the frequency also increased.
• Moseley then assigned a unique whole number to
each element and called it the atomic number.
Glenn Seaborg and Actinoid Series
• After discovering U, Th, Pa, Seaborg was
advised to revise the periodic table.
• After examining the properties of the
elements, Seaborg was convinced that the
elements were part of the inner transition
elements because they behaved similarly to
other transition metals.
• Instead of creating a new system, he
integrated the elements into a new section
called the actinoid series, which would be
inserted into the transition elements.
•
How did this change the
Mendeleev’s Periodic Table?
• This allowed elements
such as Argon, which
has a greater atomic
mass, to be placed
before Potassium.
• It also allowed others
to determine fully the
number of “holes” in
the periodic table.
Modern Periodic Law
• Elements are placed on the table according
to increasing atomic number.
• Trends and properties are a result of the
similar atomic structure in a family.
Metals and Non Metals
Metals
And
Non-metals
• For the following, write ‘M’ if it applies to metals, and
‘NM’ if it applies to nonmetals:
– ____ have luster or shine
– ____ can be gases at room temperature
– ____ good conductors of heat and electricity
– ____ high melting points
– ____ ductile and malleable
– ____ are found on the right side of the periodic table
• What about metalloids?
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Classification of Elements
Classification of Elements
Chapter 6.2
Alkali Metals
•
•
•
•
____ column
___-block
# of valence electrons ____
How many valence electrons would it like to have
______
• What it will do to follow the Octet Rule ______
• Charge of ion formed______
Transition and
Inner Transition Metals
• ___, ___-blocks
• # of valence electrons ____
• With all the d and f electrons, can these metals really
ever be happy?____
• What can they do?
– ___________________
– ___________________
Alkaline Earth Metals
•
•
•
•
____ column
___-block
# of valence electrons ____
How many valence electrons would it like to have
______
• What it will do to follow the Octet Rule ______
• Charge of ion formed______
Halogens
•
•
•
•
7th column (or 17th)
___-block
# of valence electrons ____
How many valence electrons would it like to have
______
• What it will do to follow the Octet Rule ______
• Charge of ion formed______
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Noble Gases
•
•
•
•
8th column (or 18th)
___-block
# of valence electrons ____
How many valence electrons would it like to have
______
• What it will do to follow the Octet Rule ______
• Charge of ion formed______
Electron Shielding
• Electron
Shielding
reduces the
force of the
nucleus’
positive charge
and its outermost
electron due to
cancellation of
charge by inner
electrons
Effective Nuclear Charge
• In many atoms, there is repulsion between
electrons.
• Each electron is also attracted to the
positive charge on the nucleus.
• “Effective” or actual nuclear charge (“pull”)
experienced by an valence electron is
lessened as a result of repulsion (“push”)
from core electrons.
Periodic Trends
Chapter 6.3
Effective Nuclear Charge
• Effective nuclear Charge-changes because
of shielding
Effective Nuclear Charge
• Zeff = Z – S where Z is the nuclear charge
and S is the shielding constant.
• Electrons in the same type of orbital are
not effective in shielding one another.
• Core electrons closer to the nucleus
provide effective shielding (“push”) for
outer electrons.
• Effective nuclear charge increases from
left to right in a periodic table.
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Atomic radii
Periodic Properties
Atomic Radii:
Atomic radii
Trend:atomic radii gets ________as ones goes down the column
Reason: as one moves down the column of the PT you are
adding entire _____________ ___________ of electrons,
increasing the # of ______ ________________ and therefore
increasing the _______________ _______________
(________), and decreasing the _____________ ____________
______________.
Trend: atomic radii gets ________ as one moves left to right
Reason: as one moves left to right, you are keeping the same
number of __________ electrons (________), and increasing
the number of ________________ ________________ (pull),
therefore increasing the ___________ _____________
___________.
Atomic Radius = Snowman
Ions and Ionic Radii
Chapter 06
02
30
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Ions and Ionic Radii
Periodic Trends – Ionization Energy
Ionization energy is defined as the energy
required to remove the electron from an atom.
As you move down a column the ionization energy
_________.
R: Atoms are getting ______; ____ to lose electrons
due to ____ effective nuclear charge (pull).
As you move left-to-right across a period, the
ionization energy __________.
R: Atoms are getting ______; ____ to lose electrons
due to ____ effective nuclear charge (pull).
Periodic Trends – Ionization Energy
Ionization energy is defined as the energy
required to remove the electron from an atom.
Ionization Energy
• The smallest elements are really good at
holding onto electrons, thus it is opposite of
Atomic Radius
As you move down a column the ionization energy
_________.
R: Atoms are getting ______; ____ to lose electrons
due to ____ effective nuclear charge (pull).
As you move left-to-right across a period, the
ionization energy __________.
R: Atoms are getting ______; ____ to lose electrons
due to ____ effective nuclear charge (pull).
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Ionization Energy
• The smallest elements are really good at holding onto
electrons, thus the trend is opposite of Atomic Radius
Periodic Trends-Electronegativity
2nd vs 1st Ionization Energy
• The same trend as the 1st, but shifted
Electronegativity
Electronegativity is defined as the relative ability
of an atom to attract electrons in a chemical
bond.
T: As you move down a column the electronegativity
_______.
R: Atoms are getting ______; ____ able to attract
electrons due to ____ exposure to outside world.
As you move left-to-right across a row, the
electronegativity _________.
R: Atoms are getting ______; ____ able to attract
electrons due to ____ exposure to outside world.
Electronegativity
Periodic Trends - Reactivity
Reactivity is defined as the ability for an
atom to react
With reactivity, we must look at the 2 separate groups
metals and non-metals.
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Periodic Trends
Atomic Radius
Reactivity
WHY???
Ionization Energy
Electronegativity
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