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The Periodic Table
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Elements are arranged in a way that
shows a repeating, or periodic, pattern.
Dmitri Mendeleev created the first periodic
table of the elements in 1869.
He ordered the ~70 known elements by
their atomic masses and their chemical
properties.
He found that some elements could not be
put into groups with similar properties and
at the same time stay in order.
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Modern Periodic Table
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Later, Henry Moseley carried on the
work.
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Moseley put the elements in order of
increasing atomic NUMBER.
He found that the position of the
element corresponded to its properties.
The modern periodic table shows the
position of the element is related to :
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Atomic number AND
Arrangement of electrons in its energy
levels
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Electron Shells
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Move down P. table: Principal quantum
number (n) increases.
Distribution of electrons in an atom is
represented with a radial electron
density graph.
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Radial electron density is probability of
finding an electron at a particular distance
from the nucleus.
Electron shells are diffuse and overlap a
great deal.
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Examples of Electron
shells
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He: 1s2
Radial plot shows 1 maximum
Ne: 1s2 2s2 2p6
Radial plot shows 2 maxima ( 1 each
for the 1st and 2nd energy levels )
Ar: 1s2 2s2 2p6 3s2 3p6
Radial plot shows 3 maxima ( 1 each
for the 1st,2nd and 3rd energy levels )
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Atomic Sizes: Single atoms
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Colliding argon atoms ricochet apart
because electron clouds cannot penetrate
one another to a significant extent.
The apparent radii are determined by the
closest distances separating the nuclei
during such collisions.
This radius is called the nonbonding
radius.
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Atomic Sizes: Bonded atoms
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The distance between two nuclei is called
the bond distance.
If the two atoms making up the molecule
are the same, then ½ the bond distance is
called the bonding atomic radius of the
atom.
This radius is shorter than the nonbonding
radius.
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Atomic Sizes using Periodic Table
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As we move down a group, atoms
become larger.
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Larger n = more shells = larger radius
As we move across a period, atoms
become smaller.
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More protons = more effective nuclear
charge, Zeff
More positive charge increases the
attraction of nucleus to the electrons in the
outermost shell, so the electrons are pulled
in more “tightly,” resulting in smaller radius
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Ionization energy
Ionization energy of an ion or atom is the
minimum energy required to remove an
electron from the ground state of the isolated
gaseous atom or ion.
 The first ionization energy, I1 is the energy
required to remove one electron from an atom.
Na(g)  Na+(g) + e The 2nd ionization energy, I2, is the energy
required to remove an electron from an ion.
Na+(g)  Na2+(g) + e Larger ionization energy, harder to remove
electron.
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Periodic Trends in Ionization Energy
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Highest = Fluorine
Ionization energy decreases down a
group.
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Easier to remove electrons that are
farther from the nucleus.
Ionization energy increases across a
period.
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Zeff increases, so it’s harder to remove
an electron.
Exceptions: Removing the 1st and 4th p
electrons
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Electron Affinity
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Gain 
Lose
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Electron affinity is the energy change
when a gaseous atom gains an
electron to form a gaseous ion.
Electron affinity: Cl(g) + e-  Cl-(g)
Ionization energy: Cl(g)  Cl+(g) + eAffinity for reaction above is exothermic:
∆E = -349 kJ/mol
If adding the electron makes the species
more stable, it will be exothermic.
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Coulomb’s law
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Which law can best be used to explain
why addition of an electron to the O2–
ion is an endothermic process?
Coulomb’s law: The energy required
for the process is necessary to
overcome the electrostatic repulsion
between the electron and the already
negatively charged O2– ion.
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Ion size
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The oxide ion is isoelectronic (has
exactly the same number and
configuration of electrons) with
neon, and yet O2– is bigger than Ne.
Why?
This is Coulomb's law at work. In
any isoelectronic series the species
with the highest nuclear charge will
have the smallest radius.
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Metals
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Metallic character increases down a group and
from left to right across a period.
Metal properties:
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Lustrous (shiny)
Malleable (can be shaped)
Ductile (can be pulled into wire)
Conduct electricity
Metal oxides form basic ionic solids:
Metal oxide + water  metal hydroxide
 Metal oxides react with acids to form salt and
water
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Metals
Metal oxides form basic ionic solids:
Metal oxide + water  metal hydroxide
MgO(s) + H2O(l)  Mg(OH)2(s)
 Metal oxides react with acids to form
salt and water
MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(l)
 Most neutral metals are oxidized rather
than reduced.
 Metals have low ionization energies.
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Metal reactivity
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Which of the alkali metals would
you expect to react most violently
with water? Li, Na, K, Rb
Of these four, rubidium has the
lowest ionization energy, making it
the most reactive. Rubidium reacts
explosively with water.
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Nonmetals
Lower melting points than metals
 Diatomic molecules are nonmetals.
 Most nonmetal oxides are acidic:
Nonmetal oxide + water  acid
P4O10(s) + 6H2O(l)  4H3PO4(aq)
 Nonmetal oxides react with bases to form
salt and water:
CO2(g) + 2NaOH(aq)  Na2CO3(aq) + H2O(l)
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Nonmetallic oxides
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Which nonmetallic oxide would you
expect to be the strongest acid?
NO2, N2O, N2O4, N2O5
N2O5: Nitrogen has an oxidation
state of +5 in this compound. In
general, the higher the oxidation
state of the nonmetal, the more
acidic the nonmetal oxide.
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