Transcript The Periodic Table - Miami East Schools

Organizing the Elements
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How do you organize a collection of different
things?
Elements are organized into groups based
upon their chemical properties
Elements that have similar chemical behavior
are grouped together
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Mendeleev is the father
of the modern periodic
table (1869)
Organized elements by
◦ increasing atomic mass
◦ Repeating pattern of
properties
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Was able to predict
unknown elements
using his table
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Reorganized periodic
table based upon
increasing atomic
number
Resolved
inconsistencies of
Mendeleev’s table
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When elements are arranged by increasing
atomic number….
Their physical and chemical properties repeat
in a periodic fashion (repeating pattern)
Properties of elements within a group are very
similar. Why?
Properties of elements vary in a regular way
across a period
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Most elements are metals (~ 80%)
Properties of metals
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Solid at RT (except for Hg)
Good conductors of heat
Electrical conductivity
Luster
Ductile
Malleable
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Most nonmetals are gases at RT
Some solids, one liquid (Br)
Properties of Nonmetals
◦ Poor conductors of heat
◦ Poor electrical conductivity (except C)
◦ Brittle
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Metalloids sometimes called “semi-metals”
What does this tell you about metalloids?
Seven Metalloids
Properties of Metalloids
◦ Intermediate between metals and nonmetals
◦ Often depends upon conditions
◦ Example: Si is a poor electrical conductor, but is a
good conductor when mixed with small amounts of
boron.
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Groups
◦ vertical columns
◦ 1-18
◦ IA – VIII A are
“representative elements”
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Periods
◦ horizontal rows
◦ 1-7
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Squares
◦ element symbol and other
information
◦ ~115
Inner Transition Metals
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Group 1:
Group 2:
Groups 3-12:
Group 17:
Group 18:
___________________?
___________________?
___________________?
___________________?
___________________?
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Elements are within a
group have the same
ending electron
configurations
This is why elements
within a group have
similar chemical
properties
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Atomic Radius: half the distance between nuclei of
two like atoms joined together
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Increases from top to bottom down a group
Decreases from left to right across a period
Arrow points
toward increase
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Down a group, added energy levels increase
radius
Because atomic number increases across a
period, there is greater nuclear (+) charge
No energy levels are added across a period
This tends to draw electrons closer to
nucleus, decreasing atomic radius
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Ions are atoms which have gained or lost
electrons
Cations are positively charged because they
have lost electrons (negative charges)
Anions are negatively charged because they
have gained electrons (negative charges).
Na0 → Na+ + eCl0 + e- → Cl-
6.3
Positive and negative ions form when
electrons are transferred between atoms.
6.3
Positive and negative ions form when
electrons are transferred between atoms.
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Cations and Anions
Cations are smaller
than their parent
atoms
Anions are larger
than their parent
atoms
Relative Sizes
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Ionic radius follows the same pattern as
atomic radius but remember…
Cations are smaller than their parent atoms
Anions are larger than their parent atoms
This is because of a change in the relative
strength of the nucleus, i.e….
◦ The ratio of protons to electrons changes
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In the case of cations, an energy level is lost
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Energy required to remove an e- from an
atom
Decreases from top to bottom down a group
Increases from left for right across a period
A0 + E i → A+ + e Na0 + Ei → Na+ + e-
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Energy required to remove 1st, 2nd, & 3rd eAl0 → Al+ + eAl+ → Al2+ + eAl2+ → Al 3+ e-
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Ei1= 578 kJ/mol
Ei2 = 1816 kJ/mol
Ei3 = 2744 kJ/mol
Why would it require more energy to remove
the 2nd electron?
The 3rd electron?
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Why are these important?
◦ These explain trends in atomic size, ionic size, and
ionization energy
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Nuclear charge
◦ tends to draw e-s toward nucleus (reducing atomic
radius)
◦ The bigger the atomic number, the more positive it is
◦ Increases across a period
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Electron Shielding:
◦ Inner (core) electrons “screen” outer electrons from
attractive force of the nucleus
◦ Allows outer e-s to move further from nucleus
(increasing atomic radius)
◦ Decreases down a group
◦ Does not change across a period
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Tendency of an atom to attract electrons to itself
when bonded to another atom
A very important property!
Electronegativity
4.5
Increases left to right across a period
Decreases down a group
4.0
3.5
Electronegativity
3.0
2.5
2.0
1.5
1.0
0.5
0.0
0
10
20
30
40
50
Atomic Number
60
70
80
90
100