Download Periodic Law

Periodic Law
History of the Periodic Table
Periodic Trends
Mendeleev - 1869
Organize elements according to
properties
 He noticed that certain properties
appeared at regular intervals
 He left empty spots in his table –
he predicted that there were
elements that scientists did not
yet know about

Mendeleev’s Periodic Table
Dmitri Mendeleev
Chinese Periodic Table
Stowe Periodic Table
Moseley - 1913
Slightly changed Mendeleev’s
table
 Ordered elements according to
Atomic Number
 Periodic Law = physical and
chemical properties of the
elements are periodic functions of
their atomic numbers

Modern Periodic Table
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Arranged according to Atomic Number
Elements with similar properties fall in
the same column (group)
Know these four families:
– Noble Gases
– Halogens
– Alkali Earth Metals
– Earth Metals
Families of the Periodic Table
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Noble Gases = Group 18, Unreactive
Halogens = Group 17, Very reactive
Alkali Metals = Group 1, Very reactive
Alkaline Earth Metals = Group 2, very
reactive
Regions of the Periodic Table
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Metals – to the left of the stair step
line
Nonmetals – to the right of the
stair step line
Metalloids – Elements that border
the stair step line (except for Al)
Transition metals – d block
Main group elements – s and p
block
Periodic Trends – Atomic
Radii
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Atomic Radii – size of the atom, half the
distance between the nuclei of identical
atoms that are bonded together
Across a period – atomic radii decreases due
to an increase in the effective nuclear charge
Down a group – atomic radii increases due to
addition of outermost electrons to higher
energy levels
Trends in Atomic Radius
 Atomic
Radius
K
Atomic Radius (pm)
250
Na
200
Li
150
100
Ar
Ne
50
0
0
5
10
Atomic Number
15
20
Table of
Atomic
Radii
Periodic Trend – Ionization
Energy
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Ionization energy (IE) – the amount of
energy it takes to remove an electron
from a neutral atom
Ion – formed when at atom loses or
gains electrons
Ionization – the process that results in
the formation of an ion

Ionization energy increases across a row.
– It becomes more difficult to lose valence electrons
as you move across a row due to increasing
effective nuclear charge.

Ionization energy decreases down a group.
– It becomes easier to remove valence electrons as
you move down a group because the electrons
are farther from the nucleus.
 Shielding Effect – the inner shell electrons
interfere with the nucleus attraction to the
valence electrons.
Trends in Ionization Energy
1st Ionization Energy (kJ)
 First
Ionization Energy
He
2500
Ne
2000
Ar
1500
1000
500
Li
Na
K
0
0
5
10
Atomic Number
15
20
E. Ionization Energy
 First
Ionization Energy
– Increases UP and to the RIGHT
1
2
3
4
5
6
7
Electron Affinity
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Electron Affinity (EA) - the energy
change that occurs when an
electron is acquired by a neutral
atom
When an atom gains an electron
easily, a large amount of energy is
released (indicated by a high
negative number). These elements
will have a high electron affinity.
Electron Affinity - Trend

In general, electron affinity increases
across a period
– Atoms accept electrons easily.

In general, electron affinity decreases
down a group.
– Larger atoms cannot accept electrons as
easily.
Electron Affinities of the Elements
Why gain or lose electrons?
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The noble gases are stable because they
have 8 electrons in their highest energy
level – a complete octet.
The other families need to gain or lose
electrons so they can also have a complete
octet.
Groups 1, 2, 3 lose valence electrons to
form cations – positive ions (cations).
Groups 5, 6, 7 gain electrons to form
anions – negative ions (anions).
Cations
Cations lose valence electrons and
become positive.
 To name a cation, add the word
ion to the end of the element
name.
 Cations are smaller than the
original atom.
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Anions
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Anions gain valence electrons and
become negative.
To name an anion, change the ending
of the element name to
–ide and add the word ion.
Anions are larger than the original
atom.
Teachers: have students label periodic table
Electronegativity
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Measure of the ability of an atom in a
chemical compound to attract
electrons.
Fluorine is the most electronegative
element and has been assigned an
electronegativity value of 4.
Electronegativity –
Periodic Trends
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Across a period, electronegativity
increases.
– The halogens highly attract electrons.
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Down a group, electronegativity
decreases.
– Larger atoms do not attract electrons as
easily.
Periodic Table of Electronegativities
Summation of Periodic Trends