Download unit iv – the periodic table

UNIT IV – THE PERIODIC TABLE /
PERIODICITY
Periodicity
Johann Dobereiner, 1817
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1. Triad - group of three elements with similar properties.
John Newlands, 1863
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1. Arranged elements in order of their atomic masses.
2. Noticed that their properties repeated every 8th element.
3. Law of Octaves - the same properties repeated every 8th
element
Dmitri Mendeleev, 1869
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Father of the Periodic Table
Believed that similar properties occurred after periods that could
vary in length:
1. First two rows: 7 elements
2. Second two rows: 17 elements
Organized the periodic table according to the properties of the
elements. Properties of the elements repeat in an orderly way.
Such a repeating pattern is "periodic"
Periodic Law - Properties of the elements are a periodic function of
their atomic masses.
Henry Moseley, 1911
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Worked with Ernest Rutherford and showed that the nucleus had
a positive charge. He recognized the pattern of chemical
properties when the elements were organized according to their
nuclear charge (i.e., atomic number), Thus, the periodic law was
revised:
Periodic Law - Properties of the elements are a periodic function of
their atomic numbers.
The periodic table is the arrangement of the elements in order of their
atomic numbers so that elements with similar properties fall in the
same column or group.
Noble Gases were added to the periodic table later - early 1900's
Lanthanides - f-Block was added in the early 1900's after the identity
and properties of these elements (at. no. 58-71) were sorted out.
Actinides - f-Block elements 90-103 were then added as they were
identified
Periodicity - can be observed by observing the properties of the
elements either along any row (Period) or column (Group/Family).
Electron Configuration and The Periodic Table
Electron Configuration determines a chemical's reactivity.
s-Block Elements (Groups 1 and 2; or Groups I A and II A)
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Outermost electrons are added to an s-orbital
Group 1: s1 - Alkali Metals
Group 2: s2 - Alkaline Earth Metals
p-Block Elements (Groups 13-18; Groups III A through VIII A)
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Outermost electrons are added to a p-orbital
Group 13: p1
Group 14: p2
Group 15: p3
Group 16: p4 - Chalcogens
Group 17: p5 - Halogens
Group 18: p6 - Noble Gases (Inert Gases)
d-Block Elements (Groups 3-12; Groups I B through VIII B)
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Outermost electrons are added to a d-orbital
Known as the transition metals
All metals
Reading the electron configuration directly off the Periodic
Table.
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Main Group Elements - s-Block and p-Block
Transition Elements - d-Block
Lanthanides and Actinides - f-Block
Elemental Properties:
Metals
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hard and shiny; conduct electricity
on left side of the periodic table
1-3 electrons in outer level
Nonmetals
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gases or brittle solids; good insulators
on right side of the periodic table
>5 electrons in outer level
Metalloids
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properties of both metals and nonmetals
Rule of Thumb
have 1-3 electrons in
Metals
the outer shell
have >5 electrons in
Nonmetals
the outer shell
The Octet Rule - Eight electrons in the outer shell render an atom
essentially unreactive.
Rule of Thumb: An atom having a filled or half-filled sublevel is
slightly more stable (less reactive) than an atom without a filed or
half-filled sublevel
Relative Atomic Stability
(Decreasing order of stability)
Full outer shell
Full sublevel (s, p, d or f)
half-filled sublevel
no special arrangement of
electrons
If the last electron for an atom is in a full or half-full sublevel, then the
atom is inherently more stable
PERIODIC PROPERTIES
We will look at the following periodic properties of the elements:
1 Atomic Radii
2 Ionic Radii
3 First Ionization Energy
Oxidation Numbers / Valence
4
Electrons
5 Electronegativity
6 Electron Affinity
1. Atomic Radius
As the principal quantum number (n) increases, the size of the
electron cloud increases. That is, the atomic size increases as you go
down the table.
The reason for this is that you are adding energy levels as you go
down the table (1, 2, 3,...).
The positive charge of the nucleus increase as you go from left to right
across the table. This increase in nuclear charge increase the pull on
the electron cloud by the nucleus pulling the the electron cloud in
tighter to the nucleus. Thus, the atom decreases in size.
2. Ionic Radius
Metallic Ions
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Formed by the loss of electrons
Smaller than the atoms from which they were formed
Nonmetallic Ions
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Formed by the gain of electrons
Larger than the atoms from which they were formed
3. First Ionization Energy
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The energy required to remove the most loosely-held electron
from a neutral atom.
The first ionization energy tends to increase as atomic number
increases in any horizontal row (or period).
In any column or group, there is a gradual decrease in the first
ionization energy as atomic number increases.
Metals: low first ionization energies.
Nonmetals: high first ionization energies
Two factors which tend to lower the first ionization energy as you go
down a particular group are:
1. Shielding effect- where inner electrons block the attraction of the
nucleus for the outer electrons.
2. Increased distance of the outer electrons from the nucleus.
The first ionization energy increases as you go across the periodic
table. This is a result of an increasing nuclear charge
Factors Affecting Ionization Energy
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Nuclear Charge - ionization energy is proportional to the nuclear charge
Shielding Effect - Ionization energy is inversely proportional to the shielding
effect
Radius - Ionization energy is inversely proportional to the distance between
the nucleus and the outer electrons
Sublevel - an electron from a full or half-full sublevel requires additional
energy to be removed
4. Oxidation Numbers
Group 1
Group 2
Group 312
Group 13
Group 14
Group 15
Group 16
Group 17
Group 18
lose 1
electron
lose 2
electrons
multiple
gain/loss
lose 3
electrons
lose/gain 4
electrons
gain 3
electrons
gain 2
electrons
gain 1
electron
stable
+1
+2
+3
±4
-3
-2
-1
0
5. Electronegativity
Electronegativity - The relative tendency of an atom to attract
electrons to itself when it is bonded to another atom
Influenced by the same factors which affect ionization energy and
electron affinity
1. Size
2. Shielding effect
3. Nuclear charge
The Trends (in the Periodic Table) are the same
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increases from left to right
increases from bottom to top
The most active metals have the lowest electronegativity (Francium).
The most active nonmetals have the highest electronegativity
(Fluorine).
Many chemical properties of the elements can be organized in terms of
electronegativity.
6. Electron Affinity
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the attraction an isolated atom has for an additional electron
shows the same trend as first ionization energy (increases from
left to right, and decreases from the top down).
Bond Strength - Two elements will react with each other depending on
their relative attraction for electrons