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Applied Chemistery-I
Session 2009-10 , UET Lahore
Period Classification of elements on the basis of structure of atoms
The periodiс table of the chemical elements (also Mendeleev's table, periodic table of the elements or
just periodic table) is a tabular display of the chemical elements. Although precursors to this table exist,
its invention is generally credited to Russian chemist Dmitri Mendeleev in 1869, who intended the
table to illustrate recurring ("periodic") trends in the properties of the elements. The layout of the table
has been refined and extended over time, as new elements have been discovered, and new theoretical
models have been developed to explain chemical behavior.
The periodic table is now ubiquitous within the academic discipline of chemistry, providing an
extremely useful framework to classify, systematize, and compare all of the many different forms of
chemical behavior. The table has found wide application in chemistry, physics, biology, and
engineering, especially chemical engineering. The current standard table contains 117 elements as of
July 2009 (elements 1–116 and element 118).
Structure of the periodic table
Group
1
#
Period
1
1
H
3
2
Li
11
N
3
a
4
5
6
7
2 3
5
6
7
8
9 10 11 12 13 14 15 16
17
18
8
O
9
F
2
He
10
Ne
13 14 15 16
Al Si P S
17
Cl
18
Ar
27 28 29 30 31 32 33 34
Co Ni Cu Zn Ga Ge As Se
35
Br
36
Kr
45 46 47 48 49 50 51 52
Rh Pd Ag Cd In Sn Sb Te
53 54
I Xe
4
Be
12
M
g
5
B
25
19 20 21 22 23 24
26
M
K Ca Sc Ti V Cr
Fe
n
37
42
38 39 40 41
43 44
R
M
Sr Y Zr Nb
Tc Ru
b
o
55 56 * 72 73 74 75 76
Cs Ba
Hf Ta W Re Os
77 78 79 80 81
Ir Pt Au Hg Tl
11
10 10 10 10 10 10 11 11 112
87 88 **
3
4 5 6 7 8 9 0 1 Uu
Fr Ra
Uu
Rf Db Sg Bh Hs Mt Ds Rg b
t
57
La
89
** Actinoids A
c
* Lanthanoids
1
4
58 59 60 61 62 63 64
Ce Pr Nd Pm Sm Eu Gd
95 96
90 91 92 93 94
A C
Th Pa U Np Pu
m m
6
C
7
N
82 83 84
Pb Bi Po
85
At
(117
114 115 116
)
Uu Uu Uu
(Uus
q p h
)
65 66 67 68 69 70 71
Tb Dy Ho Er Tm Yb Lu
97 98 99 100 101 102 103
Bk Cf Es Fm Md No Lr
Hasan Qayyum Chohan , Reg. No. 2009-CH-204
Chemical Engineering
UET Lahore
86
Rn
118
Uu
o
Applied Chemistery-I
Session 2009-10 , UET Lahore
This common arrangement of the periodic table separates the
lanthanoids and actinoids from other elements. The wide periodic table
incorporates the f-block. The extended periodic table adds the 8th and 9th periods,
incorporating the f-block and adding the theoretical g-block.
Element categories in the periodic table
Metals
Alkali
metals
Nonmetals
Inner transition
Alkaline
Transition Other Metalloids
Other
Noble Unknown
elements
earth
Halogens
elements metals
nonmetals
gases
metals Lanthanoids Actinoids
Borders show natural occurrence
Atomic number colors show state at
standard temperature and pressure (0 °C
and 1 atm)
Solids
Liquids
Gases
Unknown
Primordial
From
decay
Synthetic
(Undiscovered)
Other alternative periodic tables exist:Some versions of the table show a dark stair-step line along
the metalloids. Metals are to the left of the line and non-metals to the right. The layout of the periodic
table demonstrates recurring ("periodic") chemical properties. Elements are listed in order of increasing
atomic number (i.e., the number of protons in the atomic nucleus). Rows are arranged so that elements
with similar properties fall into the same columns (groups or families). According to quantum
mechanical theories of electron configuration within atoms, each row (period) in the table
corresponded to the filling of a quantum shell of electrons. There are progressively longer periods
further down the table, grouping the elements into s-, p-, d- and f-blocks to reflect their electron
configuration.In printed tables, each element is usually listed with its element symbol and atomic
number; many versions of the table also list the element's atomic mass and other information, such as
its abbreviated electron configuration, electronegativity and most common valence numbers.As of
2006, the table contains 117 chemical elements whose discoveries have been confirmed. Ninety-four
are found naturally on Earth, and the rest are synthetic elements that have been produced artificially in
particle accelerators. Elements 43 (technetium), 61 (promethium) and all elements greater than 83
(bismuth), beginning with 84 (polonium) have no stable isotopes. The atomic mass of each of these
element's isotope having the longest half-life is typically reported on periodic tables with parentheses.
Isotopes of elements 43, 61, 93 (neptunium) and 94 (plutonium), first discovered synthetically, have
since been discovered in trace amounts on Earth as products of natural radioactive decay processes.The
primary determinant of an element's chemical properties is its electron configuration, particularly the
valence shell electrons. For instance, any atoms with four valence electrons occupying p orbitals will
exhibit some similarity. The type of orbital in which the atom's outermost electrons reside determines
the "block" to which it belongs. The number of valence shell electrons determines the family, or group,
to which the element belongs.The total number of electron shells an atom has determines the period to
which it belongs. Each shell is divided into different subshells, which as atomic number increases are
filled in roughly this order (the Aufbau principle):
Subshell: S G F D P
Period
2
1
1s
2
2s
2p
Hasan Qayyum Chohan , Reg. No. 2009-CH-204
Chemical Engineering
UET Lahore
Applied Chemistery-I
Session 2009-10 , UET Lahore
3
3s
3p
4
4s
3d 4p
5
5s
4d 5p
6
6s
4f 5d 6p
7
7s
5f 6d 7p
8
8s 5g 6f 7d 8p
Hence the structure of the table. Since the outermost electrons determine chemical properties, those
with the same number of valence electrons are grouped together.
Progressing through a group from lightest element to heaviest element, the outer-shell electrons (those
most readily accessible for participation in chemical reactions) are all in the same type of orbital, with a
similar shape, but with increasingly higher energy and average distance from the nucleus. For instance,
the outer-shell (or "valence") electrons of the first group, headed by hydrogen, all have one electron in
an s orbital. In hydrogen, that s orbital is in the lowest possible energy state of any atom, the first-shell
orbital (and represented by hydrogen's position in the first period of the table). In francium, the heaviest
element of the group, the outer-shell electron is in the seventh-shell orbital, significantly further out on
average from the nucleus than those electrons filling all the shells below it in energy. As another
example, both carbon and lead have four electrons in their outer shell orbitals.
Note that as atomic number (i.e., charge on the atomic nucleus) increases, this leads to greater spinorbit coupling between the nucleus and the electrons, reducing the validity of the quantum mechanical
orbital approximation model, which considers each atomic orbital as a separate entity.
The elements ununbium, ununtrium, ununquadium, etc. are elements that have been discovered, but so
far have not received a trivial name yet. There is a system for naming them temporarily.
Classification
Groups
Main article: Group (periodic table)
A group or family is a vertical column in the periodic table. Groups are considered the most important
method of classifying the elements. In some groups, the elements have very similar properties and
exhibit a clear trend in properties down the group. These groups tend to be given trivial (unsystematic)
names, e.g., the alkali metals, alkaline earth metals, halogens, pnictogens, chalcogens, and noble gases.
Some other groups in the periodic table display fewer similarities and/or vertical trends (for example
Group 14), and these have no trivial names and are referred to simply by their group numbers
Periods
Main article: Period (periodic table)
A period is a horizontal row in the periodic table. Although groups are the most common way of
classifying elements, there are some regions of the periodic table where the horizontal trends and
similarities in properties are more significant than vertical group trends. This can be true in the d-block
(or "transition metals"), and especially for the f-block, where the lanthanoids and actinoids form two
substantial horizontal series of elements.
Blocks
Main article: Periodic table block
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Hasan Qayyum Chohan , Reg. No. 2009-CH-204
Chemical Engineering
UET Lahore
Applied Chemistery-I
Session 2009-10 , UET Lahore
Because of the importance of the outermost shell, the different regions of the periodic
table are sometimes referred to as periodic table blocks, named according to the
subshell in which the "last" electron resides. The s-block comprises the first two groups (alkali metals
and alkaline earth metals) as well as hydrogen and helium. The p-block comprises the last six groups
(groups 13 through 18) and contains, among others, all of the semimetals. The d-block comprises
groups 3 through 12 and contains all of the transition metals. The f-block, usually offset below the rest
of the periodic table, comprises the rare earth metals.
Other
The chemical elements are also grouped together in other ways. Some of these groupings are often
illustrated on the periodic table, such as transition metals, poor metals, and metalloids. Other informal
groupings exist, such as the platinum group and the noble metals.
Periodicity of chemical properties
The main value of the periodic table is the ability to predict the chemical properties of an element based
on its location on the table. It should be noted that the properties vary differently when moving
vertically along the columns of the table than when moving horizontally along the rows.
Periodic trends of groups
Modern quantum mechanical theories of atomic structure explain group trends by proposing that
elements within the same group have the same electron configurations in their valence shell, which is
the most important factor in accounting for their similar properties. Elements in the same group also
show patterns in their atomic radius, ionization energy, and electronegativity. From top to bottom in a
group, the atomic radii of the elements increase. Since there are more filled energy levels, valence
electrons are found farther from the nucleus. From the top, each successive element has a lower
ionization energy because it is easier to remove an electron since the atoms are less tightly bound.
Similarly, a group will also see a top to bottom decrease in electronegativity due to an increasing
distance between valence electrons and the nucleus.
Periodic trends of periods
Periodic trend for ionization energy. Each period begins at a minimum for the alkali metals, and ends at
a maximum for the noble gases.
Elements in the same period show trends in atomic radius, ionization energy, electron affinity, and
electronegativity. Moving left to right across a period, atomic radius usually decreases. This occurs
because each successive element has an added proton and electron which causes the electron to be
drawn closer to the nucleus. This decrease in atomic radius also causes the ionization energy to
increase when moving from left to right across a period. The more tightly bound an element is, the
more energy is required to remove an electron. Similarly, electronegativity will increase in the same
manner as ionization energy because of the amount of pull that is exerted on the electrons by the
nucleus. Electron affinity also shows a slight trend across a period. Metals (left side of a period)
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Hasan Qayyum Chohan , Reg. No. 2009-CH-204
Chemical Engineering
UET Lahore
Applied Chemistery-I
Session 2009-10 , UET Lahore
generally have a lower electron affinity than nonmetals (right side of a period) with the
exception of the noble gases.
History
Main article: History of the periodic table
In 1789, Antoine Lavoisier published a list of 33 chemical elements. Although Lavoisier grouped the
elements into gases, metals, non-metals, and earths, chemists spent the following century searching for
a more precise classification scheme. In 1829, Johann Wolfgang Döbereiner observed that many of the
elements could be grouped into triads (groups of three) based on their chemical properties. Lithium,
sodium, and potassium, for example, were grouped together as being soft, reactive metals. Döbereiner
also observed that, when arranged by atomic weight, the second member of each triad was roughly the
average of the first and the third. This became known as the Law of triads.[citation needed] German chemist
Leopold Gmelin worked with this system, and by 1843 he had identified ten triads, three groups of
four, and one group of five. Jean Baptiste Dumas published work in 1857 describing relationships
between various groups of metals. Although various chemists were able to identify relationships
between small groups of elements, they had yet to build one scheme that encompassed them all.
German chemist August Kekulé had observed in 1858 that carbon has a tendency to bond with other
elements in a ratio of one to four. Methane, for example, has one carbon atom and four hydrogen
atoms. This concept eventually became known as valency. In 1864, fellow German chemist Julius
Lothar Meyer published a table of the 49 known elements arranged by valency. The table revealed that
elements with similar properties often shared the same valency.
English chemist John Newlands published a series of papers in 1864 and 1865 that described his
attempt at classifying the elements: When listed in order of increasing atomic weight, similar physical
and chemical properties recurred at intervals of eight, which he likened to the octaves of music. This
law of octaves, however, was ridiculed by his contemporaries.
Portrait of Dmitri Mendeleev
Russian chemistry professor Dmitri Ivanovich Mendeleev and Julius Lothar Meyer independently
published their periodic tables in 1869 and 1870, respectively. They both constructed their tables in a
similar manner: by listing the elements in a row or column in order of atomic weight and starting a new
row or column when the characteristics of the elements began to repeat. The success of Mendeleev's
table came from two decisions he made: The first was to leave gaps in the table when it seemed that the
corresponding element had not yet been discovered. Mendeleev was not the first chemist to do so, but
he went a step further by using the trends in his periodic table to predict the properties of those missing
elements, such as gallium and germanium. The second decision was to occasionally ignore the order
suggested by the atomic weights and switch adjacent elements, such as cobalt and nickel, to better
classify them into chemical families. With the development of theories of atomic structure, it became
apparent that Mendeleev had inadvertently listed the elements in order of increasing atomic number.
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Hasan Qayyum Chohan , Reg. No. 2009-CH-204
Chemical Engineering
UET Lahore
Applied Chemistery-I
Session 2009-10 , UET Lahore
With the development of modern quantum mechanical theories of electron
configurations within atoms, it became apparent that each row (or period) in the table
corresponded to the filling of a quantum shell of electrons. In Mendeleev's original table, each period
was the same length. However, because larger atoms have more electron sub-shells, modern tables
have progressively longer periods further down the table.
In the years that followed after Mendeleev published his periodic table, the gaps he left were filled as
chemists discovered more chemical elements. The last naturally-occurring element to be discovered
was Francium (referred to by Mendeleev as eka-caesium) in 1939. The periodic table has also grown
with the addition of synthetic and transuranic elements. The first transuranic element to be discovered
was neptunium, which was formed by bombarding uranium with neutrons in a cyclotron in 1939.
Quantum number
Quantum numbers describe values of conserved quantities in the dynamics of the quantum
system. Perhaps the most peculiar aspect of quantum mechanics is the quantization of observable
quantities. This is distinguished from classical mechanics where the values can range continuously.
They often describe specifically the energies of electrons in atoms, but other possibilities include
angular momentum, spin etc. Since any quantum system can have one or more quantum numbers, it is
a rigorous job to list all possible quantum numbers
How many quantum numbers?
The question of how many quantum numbers are needed to describe any given system has no universal
answer, although for each system one must find the answer for a full analysis of the system. The
dynamics of any quantum system are described by a quantum Hamiltonian, H. There is one quantum
number of the system corresponding to the energy, i.e., the eigenvalue of the Hamiltonian. There is also
one quantum number for each operator O that commutes with the Hamiltonian (i.e. satisfies the relation
OH = HO). These are all the quantum numbers that the system can have. Note that the operators O
defining the quantum numbers should be independent of each other. Often there is more than one way
to choose a set of independent operators. Consequently, in different situations different sets of quantum
numbers may be used for the description of the same system.
These are conventionally known as


The principal quantum number (n = 1, 2, 3, 4 ...) denotes the eigenvalue of H with the J2 part
removed. This number therefore has a dependence only on the distance between the electron
and the nucleus (ie, the radial coordinate, r). The average distance increases with n, and hence
quantum states with different principal quantum numbers are said to belong to different shells.
The azimuthal quantum number (l = 0, 1 ... n−1) (also known as the angular quantum
number or orbital quantum number) gives the orbital angular momentum through the
relation
. In chemistry, this quantum number is very important, since it
specifies the shape of an atomic orbital and strongly influences chemical bonds and bond
angles. In some contexts, l=0 is called an s orbital, l=1, a p orbital, l=2, a d orbital and l=3, an
f orbital.

The magnetic quantum number (ml = −l, −l+1 ... 0 ... l−1, l) is the eigenvalue,
This is the projection of the orbital angular momentum along a specified axis.
.
Results from spectroscopy indicated that up to two electrons can occupy a single orbital. However two
electrons can never have the same exact quantum state nor the same set of quantum numbers according
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Hasan Qayyum Chohan , Reg. No. 2009-CH-204
Chemical Engineering
UET Lahore
Applied Chemistery-I
Session 2009-10 , UET Lahore
to Hund's Rules, which addresses the Pauli exclusion principle. A fourth quantum
number with two possible values was added as an ad hoc assumption to resolve the
conflict; this supposition could later be explained in detail by relativistic quantum mechanics and from
the results of the renowned Stern-Gerlach experiment.

The spin projection quantum number (ms = −1/2 or +1/2), is the intrinsic angular momentum
of the electron. This is the projection of the spin s=1/2 along the specified axis.
To summarize, the quantum state of an electron is determined by its quantum numbers:
name
symbol
orbital
meaning
principal
quantum
number
shell
azimuthal
quantum
number
(angular
momentum)
subshell (s
orbital is
listed as 0, p
orbital as 1
etc.)
magnetic
quantum
number,
(projection of
angular
momentum)
energy shift
(orientation
of the
subshell's
shape)
spin
projection
quantum
number
spin of the
electron (-1/2
= counterclockwise,
1/2 =
clockwise)
range of values
value example
for
:
for
:
for an electron, either:
Example: The quantum numbers used to refer to the outermost valence electron of the Carbon (C)
atom, which is located in the 2p atomic orbital, are; n = 2 (group 2), l = 1 or 0, ml = 1, or 0, or −1, ms =
−1/2 or 1/2.
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Hasan Qayyum Chohan , Reg. No. 2009-CH-204
Chemical Engineering
UET Lahore
Applied Chemistery-I
Session 2009-10 , UET Lahore
Note that molecular orbitals require totally different quantum numbers, because the
Hamiltonian and its symmetries are quite different.
Quantum numbers with spin-orbit interaction
For more details on this topic, see Clebsch-Gordan coefficients.
When one takes the spin-orbit interaction into consideration, l, m and s no longer commute with the
Hamiltonian, and their value therefore changes over time. Thus another set of quantum numbers should
be used. This set includes



The total angular momentum quantum number (j = 1/2,3/2 ... n−1/2) gives the total angular
momentum through the relation
.
The projection of the total angular momentum along a specified axis (mj = -j,-j+1... j), which is
analogous to m, and satisfies mj = ml + ms.
Parity. This is the eigenvalue under reflection, and is positive (i.e. +1) for states which came
from even l and negative (i.e. -1) for states which came from odd l. The former is also known
as even parity and the latter as odd parity
For example, consider the following eight states, defined by their quantum numbers:
1.
2.
3.
4.
5.
6.
7.
8.
n = 2, l = 1, ml = 1, ms = +1/2
n = 2, l = 1, ml = 1, ms = -1/2
n = 2, l = 1, ml = 0, ms = +1/2
n = 2, l = 1, ml = 0, ms = -1/2
n = 2, l = 1, ml = -1, ms = +1/2
n = 2, l = 1, ml = -1, ms = -1/2
n = 2, l = 0, ml = 0, ms = +1/2
n = 2, l = 0, ml = 0, ms = -1/2
The quantum states in the system can be described as linear combination of these eight states.
However, in the presence of spin-orbit interaction, if one wants to describe the same system by eight
states which are eigenvectors of the Hamiltonian (i.e. each represents a state which does not mix with
others over time), we should consider the following eight states:








8
j = 3/2, mj = 3/2, odd parity (coming from state (1) above)
j = 3/2, mj = 1/2, odd parity (coming from states (2) and (3) above)
j = 3/2, mj = -1/2, odd parity (coming from states (4) and (5) above)
j = 3/2, mj = -3/2, odd parity (coming from state (6))
j = 1/2, mj = 1/2, odd parity (coming from states (2) and (3) above)
j = 1/2, mj = -1/2, odd parity (coming from states (4) and (5) above)
j = 1/2, mj = 1/2, even parity (coming from state (7) above)
j = 1/2, mj = -1/2, even parity (coming from state (8) above)
Hasan Qayyum Chohan , Reg. No. 2009-CH-204
Chemical Engineering
UET Lahore