Periodic Properties
of Elements
Tadas Rimkus
Krishna Trehan
Rachel Won
Soo Jeon
What is the SWBAT????
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Objective
 Know a brief history of the periodic table
 Know the effective nuclear charge
 Examine periodic tends in the atomic size,
ionization
energy, and electron affinity
 Examine the sizes of ions and their electron
configurations
 Explore some of the difference in the physical and
chemical properties of metals, nonmetals, and
metalloids.
 Discuss some periodic trends.
I <3 The Periodic Table
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Mendeleev, Meyer, Moseley strived to investigate the possibilities of
classifying elements in useful ways.
Mendeleev and Meyer - developed the periodic table on the basis of
the similarity in chemical and physical properties. He predicted Ga,
Ge, and Sc.
Moseley- established that each element has a unique atom number.
- found out that each element produces X rays of a
unique frequency; the frequency increases as the atomic mass
increases
Rutherford- proposed the nuclear model of the atom.
Element in the same column of the periodic table have the same
number of electrons
Sizes of Atoms and Ions
Ionic size - the size of an ion plays an
important role in determining the structure
and stability of ionic solids!
 Periodic Trend - Metals tend to lose
electrons. The size of the atom becomes
smaller with the loss of each!
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Continue…
Consecutive electron. The nuclear charge
can have a greater pull on the electrons
with the loss of each electron.
 Iso-electronic ions - ions containing the
same number of electrons, thus having the
same electron configuration!
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Continue… AGAIN!
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Periodic Trend - Nonmetals tend to gain
electrons. The size of the atom becomes
larger with the addition of each
consecutive electron. The nuclear charge
has less of a pull on the electrons with the
gain of each electron due to greater
electron repulsion!
7.5 Electron Affinities
Ionization Energy: energy required to
remove an valence electron
 Cl(g)  Cl+(g) + e- ΔE = 1251 kJ/mol
 See how E is positive?
 This is the energy which you must add, in
order to remove an electron from Chlorine
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Electron Affinity
Did you know that atoms can gain
electrons to gain a negative charge?!
 This energy change that occurs is called
ELECTRON AFFINITY because it
measures the ATTRACTION between the
newly added electron and the atom
 It shows how much the atom WANTS the
electron!
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Electron Affinity
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So……. FLUORINE
wants an electron
MORE than
LITHIIUM, meaning it
would have a
LARGER electron
affinity!
http://www.chemguide.co.uk/inorganic/group7/eadiag.gif
Ionization Energy
the amount of energy need to remove an
electron from a specific atom or ion in its
ground state
 1st Ionization Energy
 energy needed to remove the 1st electron
from an atom
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Ionization Energy… CONTINUE
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ACROSS- increase
As you go across the periodic table, the
electrons are closer to the nucleus increasing
the energy necessary to remove an electron
DOWN- decrease
As you go down the periodic table, the electrons
are farther from the nucleus
Continue!
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Nuclear Charge
the change in the nucleus or the number of protons
ACROSS- increase
DOWN- increase
Atomic Radius
one half the distance from center to center of like atom downincreases
As you go down the periodic table, a new energy level is added
increasing the size of the atoms
Ex) of the following, which has the highest and lowest first ionization
energy?
Answer: Highest- C
Lowest- AL
Periodic Trends Of Electron Affinity
Electron Affinity tends to decrease as we
go from left to right!
 BUT WHY?!?!?!?!?!?!?!?!?!?!?!
 This is because the elements on LEFT
side of the Periodic Table of LESS valence
electrons.
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Electron Affinity
the energy change when an electron is
added to a neutral atom to from a negative
ion
 ACROSS- increase
 DOWN- decrease
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Continue?
The elements on the right, are one
electron away from GETTING A FULL
OCTET!
 They will have a high affinity, so they can
gain the last electron.
 (The more negative the electron affinity,
the greater the attraction for an electron)
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METALS!!! Woohooo
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Metals are the shiny things we have learned to
love and treasure
They are ductile, conduct heat and electricity,
and malleable.
They are all solid (except Mercury!) at room
temperature.
They can have high melting points, for example
Chromium has a melting point at 1900 C.
Metals, CONTINUED!
Metals have low ionization energies
(energy required to remove an valence
electron)
 WHICH MEANS, they tend to form positive
ions EASILY!!!!!!!
 Because of this… when metals go through
CHEMICAL REACTIONS they are
OXIDIZED
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But what does oxidized mean?!?!
An metal has been OXIDIZED when it
loses electrons.
 Metals are oxidized by many different
substances INCLUDING OXYGEN GAS
(O2) and ACIDS!
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Alkali and Alkaline Earth Metals
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The charge of ALL alkali metals is ALWAYS 1+
The charge of ALL alkaline earth metals is
always 2+
Transition Metals!
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Many transition
metals have a +2
charge, BUT +2 AND
+3 is also
encountered!
A characteristic of the
Transition elements is
to be able to form
ions!
NONMETALS!
Nonmetals are overall worse than metals
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 They are not lustrous and are generally
poor conductors of heat and electricity.
 They have a lower melting point than
those of metals, BUTT the diamond form
of carbon melts at 3570 C
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Nonmetals, CONTINUED!
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UNDER ORDINARY CONDITION:
 Seven
elements of the nonmetals exist in
diatomic elements ****BrNClHOF****
 Because of their ELECTRON AFFINITY,
nonmetals tend to GAIN ELECTRONS when
they react with metals.
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Compounds composed entirely of
nonmetals are molecular substances
 Ex:
Oxides, Halides, and Hydrides
METALLOIDS
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Properties of Metalloids include propertis
from both Metals and Nonmetals
 Ex:
Silicon is shiny like a metal, but is brittle
like a nonmetal
Group Trends of Active Metals!
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Group 1A Alkali Metals
 Soft,
metallic solids
 Shiny, and have a high thermal electrical
conductivity.
 Low density, and low melting points.
 As you move down the group, the atomic
radius increases, and ionization energy
decreases
Alakali metals are extremely reactive
toward WATER and OXYGEN
 Alkali Metal Ions are colorless and
produce a characteristic a color when
combusted
 EX: Burning sodium produces a yellow
flame.
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Group 2A: Alkaline Earth Metals
Ionization energies of these metals are
low, but not as low as the alkali metals
 They are less reactive then the 1A Metals
 The “heavier” alkaline earth metals are
more reactive than the lighter ones
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Group Trends for Non Metals
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Hydrogen:
 Even
though Hydrogen is in the Alkali metals,
it does not belong to any specific group
 However, Hydrogen can have metallic
properties under extreme pressure
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Ex: the center of Saturn is surrounded by a thick
shell of pressurized Hydrogen
Group 6A: The Oxygen Group
In the 6A group, there is a change from
non-metallic to metallic character
 Oxygen through Selenium are non-metals
and Tellurium has some metallic properties
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Group 7A: Halogens
The Halogens are tpically nonmetals
 Their melting and boiling points increase
as they go down the coloumn.
 Halogens, have ridiculously high electron
afifnities, this is because they really want
to get ONE MORE ELECTRON, to mimic
the Noble Gasses.
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Noble Gasses
All gasses at room temperature.
 Because of their stable and filled outer
orbitals, they are extremely un-reactive
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