Why is the periodic table shaped like it is and how are
the elements arranged?
The Periodic Table - History
– Two scientists, Dimitri Mendeleev (Russia)
and Lothar Meyer (Germany)
• properties of elements did not change
smoothly with increasing atomic mas.
• Instead the properties of the elements repeated
periodically.
– Periodic Law: the properties of the
elements repeat periodically as the
elements are arranged in order of
increasing atomic number (# of protons)
– This periodic law is used to form the
Periodic Table
II. The Periodic Table
• John Alexander Newlands
– Arranged elements in order of increasing
atomic masses
– Noticed some properties recurring over and
over again – he called this the periodic law
• Dmitri Mendeleev
– Published the periodic table
of elements
– Very confident as he left spaces
empty – assumed those elements
were not yet discovered
Elements are arranged:
Vertically into Groups
Horizontally Into Periods
Groups of Elements
 Vertical columns on the periodic table
 Similar physical properties
 Similar chemical properties
Why?
Why?
If you looked at one
atom of every element
in a group you would
see…
Each atom has the same number of electrons in
it’s outermost shell.
Valence Electrons
for group A elements
s block
+1
ns1 +2
2
1A ns
Charge on stable
monatomic ion
p block
+3 ± 4 -3
-2
-1
8A
3A 4A 5A 6A 7A
2A
Group B
105
Db
107
Bh
Period number gives shell (n)
0
Elements are arranged according to atomic #
and e- configuration.
Li: 3 e-’s
1s2 2s1
Na: 11 e-’s 1s2 2s2 2p6 3s1
K: 19 e-’s
1s2 2s2 2p6 3s2 3p6 4s1
Paramagnetic or diamagnetic?
Valence orbitals: outer shell orbitals
beyond the closest noble-gas configuration
Valence electrons: “the ones that can
react” (located in the valence orbitals).
The other e-’s are called core electrons and don’t react.
2s2
3s2
4s2
5s2
6s2
7s2
2A
Be
Mg
Ca
Sr
Ba
Ra
Elements in a vertical row have the same
number of valence electrons.
2s22p5
3s23p5
4s24p5
5s25p5
6s26p5
F
Cl
Br
I
At
• The number of outer or “valence” electrons
in an atom effects the way an atom bonds.
• The way an atom bonds determines many
properties of the element.
• This is why elements within a group
usually have similar properties.
If you looked at an atom from each
element in a period
you would see…
Details of the Periodic Table Period
• The elements in the same horizontal
row are called a period
Hydrogen and Helium are in period 1
Lithium through Neon are in period 2
4
Periods on the Periodic Table
1
2
3
4
5
6
Each atom has the same number of
electron holding shells.
An example…
The period 4 atoms each have 4 electron
containing shells
4th Shell
K (Potassium)
Kr (Krypton)
Atom
Atom
Fe (Iron) Atom
Sublevel Blocks
s1 s2
1
2
3
4
5
6
p1 p2 p3 p4 p5 p6
d1 - d10
f1 - f14
A periodic table illustrating the
building-up order.
Electronic Structure
s block
p block
d block
105
Db
107
Bh
f block
Condensed Ground-State Electron
Configurations in the First Three Periods
Los Alamos National Laboratory's Chemistry Division
Periodic Table of the Elements
+1
+2
metals
typical charge on ion
atomic number
in binary compound
atomic mass
transition metals
non-metals
+3
http://pearl1.lanl.gov/periodic/default.htm
-3 -2 -1
noble
gases
Discovery dates of elements
Physical States at Room Conditions
gases
105
Db
liquids
107
Bh
Modified from http://www.cem.msu.edu/~djm/cem384/ptable.html
solids
Natural vs. Man-Made Elements
natural elements
not found on Earth
105
Db
107
Bh
man-made elements
Molecular Nature of Free Elements
diatomic
105
Db
tetratomic
octatomic
107
Bh
All other free elements are atomic in nature.
All Isotopes Radioactive
no stable forms
105
Db
107
Bh
Hydride Chemistry
(binary hydrogen compounds)
ionic
covalent
polymeric
salts with H-
molecules with +1
chains with H-to-H bonds
1A
2A
8A
3A 4A 5A 6A 7A
metallic
interstitial H2
Group B
105
Db
107
Bh
solids
mostly gases,
some liquids
Important Groups
Each group has distinct properties
• The periodic Table is divided into several
groups based on the properties of different
atoms.
Groups of Elements
• Alkali Metals:
– Group 1 metals
– Soft, silver coloured metals
that react violently with H2O
to form basic solutions
– Most reactive: cesium &
francium
Alkali Metals reacting with water:
•
•
•
•
•
Li (Lithium)
Na (Sodium)
K (Potassium)
Rb (Rubidium)
Cs (Cesium)
What would you expect
from Francium?!?!
Chemical Periodicity
Alkali metals
Soft, silvery
colored metals
Very reactive!!!
Potassium (K),
in Water (H2O)
Alkali Metal Family
Li
Na
K
Alkaline Earth Metals
Silvery-White Metals
Fairly reactive
Many are found in rocks in
the earth’s crust
Halogens:
– Group VII A, nonmetals, highly
reactive.
– Fluorine is the most
reactive
Noble Gases:
-Group VIII A
- Generally unreactive
The Halogen Family
Cl2(g)
I2(s)
Br2(l)
Halogens
Most are Poisonous
Fairly reactive
Chlorine Gas was
used as a chemical
weapon during World
War I.
It was used by the
Nazis in World War II.
Chlorine (Cl)
Bromine (Br)
and Iodine (I)
Jellyfish lamps made with noble
gases
•
Colors Noble Gases
produce in lamp
tubes(discharge
tubes):
Ne (Neon):
orange-red
• Hg (Mercury): light blue
• Ar (Argon): pale lavender
• He (Helium): pale peach
• Kr (Krypton): pale silver
• Xe (Xenon): pale, deep blue
• Transition Metals:
Most are good
Conductors of heat and electricity
Ductile and malleable
(easily bent/hammered into wires or sheets)
. Metalloids:
Metals very close to the
“staircase” line
They have properties of metals
and non-metals.
Si (Silicon) and Ge
(Germanium) are very
important “semi-conductors”
How many things can you think
of that have Transition Metals in
them?
What are semiconductors used in?
Nonmetals
. Brittle
. Do not conduct heat and electricity
Modern Periodic Table
Columns are called Groups or Families
Elements with similar chemical and
physical properties are in the same
column
Rows are called Periods
Each period shows the pattern of
properties repeated in the next period
Main Group (Representative Group) Groups IA - VIIIA
Transition Metals - Groups IB – VIIIB
Rare Earth Elements Lanthanides (Ce - Lu) and
Actinides (Th - Lr)
Metals
about 75% of all the elements
lustrous, malleable, ductile, conduct heat and
electricity
Nonmetals
dull, brittle, insulators
Metalloids
also know as semi-metals
some properties of both metals & nonmetals
Atomic Size
Smaller
S
m
a
l
l
e
r
Li
Na
K
Be
B
C
N
O
WHY?
F
Ionization Energy
Ionization Energy (IE) - The amount of energy needed to remove an electron
from an atom or ion. Each electron in any atom or ion has a specific
ionization energy.
First Ionization Energy - The amount of energy needed to remove an electron
from the outermost shell of a neutral (uncharged) atom.
The ionization energy is the energy that must be supplied to an
atom in the gas phase in order to remove an electron. If the
electron is the first one to be removed, one then refers to the first
ionization energy for that element.
The first ionisation energy of sodium is 494 kJ.mol-1.
The ionisation energy tells us how easy it is to convert an
element to a cation. The lower the first ionisation energy, the
easier it is to convert an element such as sodium, Na, to its
cation Na+.
Ionization Energy:
The energy required to completely remove an
e- from an atom in its gaseous state.
Mg(g)  Mg1+ + eMg1+(g)  Mg2+ + e-
1st ionization energy
2nd ionization energy
Question: Which of the above ionizations would have the
highest ionization energy and why?
electron being lost:
1st
2nd
3rd
4th
5th
6th
7th
The first ionisation energy of most elements are known to a good
degree of accuracy. It is interesting to see how these values vary with
atomic number Z:
The rare gases (He, Ne, Ar, Kr, Xe, Rn) appear at peak values of
ionization energy, which reflect their chemical inertness, while the
alkali metals (Li, Na, K, Rb, Cs) appear at minimum values of ionization
energy, in keeping with their reactivity and ease of cation formation.
Increases
Increases
Ionization energy decreases as you go down a group
Ionization energy increases as you go from left to right in a period
What is meant by metallic character?
Periodicity of ionic radii:
The size of ions differ markedly from the size of their parent atoms.
Take the case of sodium as an example:
The sodium atom (which has a single 3s electron), has an atomic
radius of 186 pm. Upon taking up energy (the ionization energy), it
loses this electron and is converted to the cation Na+, whose radius is
97 pm.
The radius of cations is always smaller than the radius of the
atoms from which they are derived, as shown in the figure
on the right, which applies to the Group I elements:
The main reason for this is that whenever metals are
converted to their cations, they always do so by losing the
electrons in their highest energy level.
Further, since the ion has less electrons than the atom from
which they are derived, there is less mutual repulsion
between these electrons, and the electron orbitals shrink to
some extent.
What happens in the case of anions? Let's take the case of chlorine:
The chlorine atom has a covalent radius of 99 pm. It can gain an extra 3p
electron (with release of energy, the electron affinity) in order to form the
chloride anion Cl-, with ionic radius 181 pm.
The radius of anions is always larger than
the covalent radius of the atoms from which
they are derived, as shown in the figure on
the right, which applies to the Group VII
elements.
The reason for this is that the additional
electron increases the mutual repulsion
between the electrons, and this causes an
expansion of the electron orbitals. Note that
the difference between the van der Waals
radius of an atom and the ionic radius of its
anion is not significant.
Periodicity of electronegativities:
When a covalent bond is formed between two identical atoms, such
as H-H or Cl-Cl, the pair of electrons which joins the atoms is evenly
shared between the two atoms. However, when two different atoms
are joined together by a covalent bond, the sharing of the electron
pair is not even, and the pair of electrons is shifted towards one of
the atoms:
The covalent bond is said to have been polarized, and one refers to
such a bond as a polar covalent bond.
The ability to polarize a covalent bond differs from one atom to another,
and is known as the electronegativity of the atom.
Electronegativities are measured on an arbitrary scale ranging from 0 (He)
to 4.1 (F).
The rare gases (He, Ne, Ar, Kr, Xe, Rn) have zero electronegativities, i.e.,
they only form covalent bonds, and this only in exceptional cases, and
have no tendency to attract electrons of that bond.
The halogens ( F, Cl, Br, I, At) appear at peak values of electronegativities.
Within a period, they have the highest tendency to polarize covalent bonds.
Note that within the group, the electronegativity decreases.
On the whole, electronegativities increase from left to right along a period
(see Li to F) and decrease from top to bottom within a group (see F to At).
Periodicity of electron affinity:
The electron affinity is the energy that is released when an atom
in the gas phase gains an electron and is thus converted to an
anion, also in the gas phase:
The electron affinity of chlorine is 349 kJ.mol-1.
Electron affinities are difficult to measure and there is no
reliable data available for most elements. However, the larger
the atom, the lower its electron affinity, as shown with Group
VII elements:
For reasons outside the scope of this discussion, the electron
affinity of fluorine is an exception to this trend.
Common Oxidation states: note the vertical similarities.