The Periodic Table

The Periodic Table
History of the Periodic Table
• Anton Lavoisier wrote a textbook in 1789, and
listed the known elements in a table. He
grouped the elements based on similar
• Johann Wolfgang Dobereiner published a
paper in 1829 describing triads of elements
where groups of 3 elements could be arranged
together based on similar properties.
• Leopold Gmelin identified ten triads, three
groups of 4 elements, and one group of 5
elements in 1843.
• August Kekule in 1858 showed relationships
between elements and their bonding strength
which he called valency.
• Stannislao Cannizzaro in 1860 established a
universal method for measuring atomic mass.
• Julius Lothar Meyer in 1864 arranged the 49
known elements by valency. Elements with
similar properties had the same valency.
• John Newlands in 1864-1865 published a table
that arranged elements based on increasing
atomic mass. His periodic table repeated
properties after 8 elements, and he called this
periodicity “The Law of Octaves”, similar to
musical note periodicity of 8 notes.
• Dmitri Ivanovich Mendeleev published his
periodic table in 1869. He arranged the
elements in rows by increasing mass then
creating a new row when the properties
• Meyer also created a similar periodic table in
• Mendeleev however supported his table by
ignoring discrepancies in mass and leaving
gaps for missing elements.
• Meyer did not make any predictions that
would validate his reasoning for his table.
Mendeleev’s Periodic Table
• Arranged by increasing atomic mass
• Some elements had “switched positions”
where the mass placed them in different
columns than their properties indicated
– He reasoned that the switched elements masses
weren’t measured accurately enough and would
later be corrected by better measurements.
• Some gaps left in the table so elements fit into
better positions.
– These gaps were explained by Mendeleev as
elements that hadn’t yet been discovered.
– To support his table, Mendeleev made predictions
of properties for these missing elements
Mendeleev’s Predictions
Atomic Mass
68 amu
72 amu
6.0 g/cm3
5.5 g/cm3
Dark Grey
Melting Point
Oxide Formula
X2 O3
X O2
Chloride Formula
X Cl3
X Cl4
Actual Element Properties
Atomic Mass
69.72 amu
72.59 amu
5.9 g/cm3
5.35 g/cm3
Light Grey
Melting Point
29.78 °C
947 °C
Oxide Formula
Chloride Formula
The Modern Periodic Table
• Further measurements of atomic mass never
corrected the misplaced elements.
• Discovery of noble gases supported the
periodic table, by being able to be placed in a
group that connected halogens and alkali
• Mass problems became worse .
• Antonius Van den Broek in 1911, first
proposes that nuclear charge and electrons
responsible for element location on the
periodic table
• Henry Moseley tested Bohr’s hypothesis and
Van den Broek’s hypothesis, by experimentally
measuring the atomic numbers of Al to Au.
• Moseley then arranged elements based on
increasing atomic number and found the mass
misplacement on the periodic table is resolved
• Sadly, Moseley was killed in battle in WWI.
• Quantum model of the atom is responsible for
the current look of the table. The blocks are
arranged based on electron configurations.
Electron Configurations & The Periodic
• Sublevels can easily be identified on the
periodic table, and are called blocks.
– s-block: Groups 1 & 2 (including He)
– p-block: Groups 13-18 (excluding He)
– d-block: Groups 3-12
– f-block: Two rows at the bottom, Lanthanide &
Actinide series.
• The period of an element is determined from
its electron configuration.
– This is determined from the principle quantum
number for the s-sublevel & p-sublevel.
– Ex. Zr [Kr] 5s2 4d2 The 5 indicates the period!
– Ex. Pb [Xe] 6s2 4f14 5d10 6p2
6th period
• Group 1: The Alkali Metals
– Exception: Hydrogen, though in Group 1, is NOT
an alkali metal!
– Very reactive, cannot be found alone in nature.
– Very soft, can be cut with a knife.
– Low density, some float on water.
– Silvery metals with low melting points.
• Group 2: The Alkaline Metals
– Harder, denser, and higher melting points than
alkali metals.
– Very reactive, not as much as alkali metals, and
are also never found alone in nature.
• Hydrogen & Helium
1s1 & 1s2
– Though similar configurations to alkali and alkaline
metals, they don’t belong because of their unique
properties resulting from the small 1st energy level
– Hydrogen is not a metal, and has unique
properties that prevent it from being in any group.
– Helium has a configuration of a filled energy level, which
gives it extremely stable and unreactive properties, and
so is similar to the noble gases.
• Transition Metals, Groups 3-12
ns2 (n−1)d(1−10)
– Have electron configurations containing d-sublevels
filling and are therefore in the d-block.
– The d-sublevel filling has a quantum number 1 less than
the period where the element is found.
– Some elements have configurations different from the
aufbau principle due to shifts in stability.
– Transition elements are metals that are less reactive
than alkali & alkaline metals, and exhibit a larger range
of properties.
• The Halogens, Group 17
ns2 np5
– Nonmetals that are very reactive
– Nearly full s- & p-sublevels
• The Noble Gases: Group 18
ns2 np6 & He 1s2
– Nonmetals that are not reactive
– Full s- & p-sublevels completing an octet, within a
period (n). This configuration is known to be
extremely stable.
• The Main-Group Elements
– s-block & p-block of elements. (Groups 1,2 & 13-18)
These two show the periodic pattern quickly and with
common elements
• f-block of elements
ns2 (n−1)d10 (n−2)f(1−14)
• Filling f-sublevel of energy level 2 less than the
period number the element is located in
• Wedged between groups 3 & 4 in the d-block
• A.K.A. “The Rare Earth Metals”
• A.K.A. “Innertransition Metals”
– Lanthanide Series: Elements 58-71
• Shiny, dense, hard metals with similar reactivity to alkalines
– Actinide Series: Elements 90-103
• All are radioactive
• First 4, Thorium (Th) to Neptunium (Np) have been found
naturally on Earth
• The rest are all artificially created in laboratories
Periodic Trends
• Periodic Law:
–The physical & chemical properties of the
elements are periodic functions of their
atomic numbers.
• Element Properties & Their Periodic Trends:*
*The trends will be best demonstrated by the
main group elements. The trends, though
true for both the d-block & f-block, are not as
apparent as they are with the main group.
Periodic & Group Trends
• There are trends for all properties in both
Periodic (Horizontal) rows & Group (Vertical) rows
• Periodic trends in these notes will refer to the
trend as you move left to right across the table.
• Group trends in these notes will refer to the trend
as you move from top to bottom on the table.
– The trends would be the opposite if you move in
reverse of these assigned motions.
Atomic Radii
• I.O.W. The size of the atoms
• Atomic Radius is ½ the distance between two
identical atoms (same element) that are
bonded together.
• Periodic Trend: Decreases
• Group Trend: Increases
– P Trend Reason: Stronger nuclear charge pulls
electron energy levels in closer to the nucleus.
– G Trend Reason: Electrons are placed into energy
levels further away from the nucleus
Ionization Energy
• Ionization Energy is the energy required to
remove one electron from a neutral atom.
• A.K.A. The 1st Ionization Energy. The 2nd IE
would be energy to remove 2nd electron, etc.
• Periodic Trend: Increases
• Group Trend: Decreases
– P Trend Reason: Increasing Nuclear Charge-holds
the electrons more tightly
– G Trend Reason: Electrons located further away
from the nucleus, and more easily lost
Electron Affinity
• It is essentially the opposite of Ionization Energy
• Electron Affinity is the energy change that occurs
when an electron is acquired by a neutral atom.
• Periodic Trend: Decreases
• Group Trend: Increases
– P Trend Reason: Nucleus is closer to the outer edge
of the atom and requires less energy to bring new
electrons in to the atom
– G Trend Reason: Electrons added further away from
the nucleus, so requires more energy to bring new
electrons in to the atom
Ionic Radii
The size of ions
Slight difference between metals & nonmetals
Metals form Cations (+ charge ions)
Nonmetals form Anions (− charge ions)
Period Trend: Decrease*
Group Trend: Increase
– P Trend Reason: Stronger attraction from nucleus
– G Trend Reason: Electrons further away from nucleus
*Change when switching from metals to nonmetals
the ions suddenly swell, then decrease again
• Property of the ability of atoms to attract
electrons to their nuclei.
• Period Trend: Increases*
• Group Trend: Decreases*
– P Trend Reason: Nucleus is closer to outer edge of
the electron cloud
– G Trend Reason: Electron cloud is further away from
the nucleus
*Not including the Noble Gases!!!
• Most electronegative element is Fluorine, the
least electronegative element is Francium
• General property, typically defined by the amount
of energy required to separate the element from
• Metals: Periodic Trend: Decreases
Group Trend: Increases
• Nonmetals: Periodic Trend: Increases
Group Trend: Decreases
– P Trend Reason: The distance from nucleus to electrons
– G Trend Reason: The distance from nucleus to electrons
Remember-Metals want to lose electrons, Nonmetals
want to gain electrons!
Valence Electrons
• The outermost electrons of an atom are the
valence electrons.
• Main Group Elements:
Group #
Valence Electrons
Common Ionic Charge
+4 / −4
Electron Configurations of Ions
• Just like electron configurations of neutral atoms.
• Main Group elements gain or lose electrons to
have electron configurations that resemble the
nearest noble gas elements.
Ex. #1: O−2 1s2 2s2 2p6
Ne 1s2 2s2 2p6
Ex. #2: K+ 1s2 2s2 2p6 3s2 3p6 Ar 1s2 2s2 2p6 3s2 3p6
• The Transition Metals are a bit more complicated.
Many can lose multiple amounts of electrons to
reach differing levels of stability depending on
other variables the atoms experience.
Example of Transition Ions:
Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6
Fe+2 1s2 2s2 2p6 3s2 3p6 3d6
Fe+3 1s2 2s2 2p6 3s2 3p6 3d5
Fe [Ar]
Fe+3 [Ar]
    
    
    
2 4s electrons lost
2 4s & 1 3d electrons lost

similar documents