Classification of Elements MANIK

Elements
• Science has come along way
since Aristotle’s theory of Air,
Water, Fire, and Earth.
• Scientists have identified 90
naturally occurring elements,
and created about 28 others.
• As more elements were
discovered in the 19th
century chemists started to
note similarities in their
properties.
• Early attempts to order the
elements in a regular fashion
were hampered by various
difficulties.

Mendeleef’s Periodic Law and Periodic Table
 Attempts were made to classify the elements in a number of
ways.
 In 1869, a Russian scientist, Dmitri Mendeleeff made
the most significant contribution towards the classification of
elements.
 Mendeleef observed that when all the 65 elements (known at
that time) were arranged according to increasing atomic weights,
similarities and differences in their properties are repeated at
regular intervals.

Mendeleeff’s Periodic Law and Periodic Table
This was enunciated in the form of a PERIODIC LAW which
was stated as-
‘ The physical and chemical properties of elements are a periodic
function of their atomic weights, i.e., if the elements are
arranged in the increasing order of their atomic weights, the
properties of the elements are repeated after regular
intervals’

Mendeleeff’s Periodic Law and Periodic Table
Mendeleef gave a detailed comparison of the physical and
chemical properties of the elements and set up a periodic
system in which the elements were arranged in horizontal
ROWS (series) and vertical COLUMNS (groups)
according to increasing atomic weights.
Mendeleeff arranged the elements in the form of a table
which is known as Mendeleeff’s Periodic Table after
his name.

With the advancement of the knowledge about atomic structure
and discovery of new elements, in 1913, Mosley, a British
physicist, predicted that most of the defects of
Mendeleeff’s periodic table disappear, if the basis of
classification of elements is changed to in
place of atomic weight. Accordingly Mosley put forward
Modern Periodic Law which is stated as follows:
The physical and chemical properties of the elements are
periodic function of their atomic numbers, i.e., if the elements
are arranged in the increasing order of their atomic numbers, the
properties of the elements (i.e. similar elements) are repeated
after regular intervals.

Groups in the periodic table
These columns are known as
GROUPS
There are 18
GROUPS
1 4 5 6 7 8 9
10 11 12 13 14 15 16 17 18
2 3
GROUPS are also known as
FAMILIES

In the modern periodic table 109 elements (118) have been
arranged according to the increasing atomic number.
Groups:
The vertical columns shown in the periodic table are called groups
or families or simply columns.
(a) There are nine groups in all including VIII (VIIIB) group
consisting of three triads (Fe, Co, Ni; Ru, Rh, Pd; Os, Ir, Pt) and
zero groups of inert gases.

Characteristics of Modern Periodic Table
(b) Groups I to VII are sub-divided into sub-groups A and B. Thus
there are 18 vertical columns which are: IA, IIA, IIIA, IVA, VA,
VIA, VIIA, Zero, IB, IIB, IIIB, IVB, VB, VIB, VIIB and three columns of
Group VIII (VIIIB).
(c) Elements of groups IA, IIA, IIIA, IVA, VA, VIA and VIIA have their
outermost shells incomplete while each of their inner
shell is complete. These elements are called normal or
representative elements. These elements consist of some
metals, non-metals and metalloids.

Electron Distribution in shell
11p+
Na atomwith1 valence e-

Electron Distribution in shell
17p+
Chlorine atomwith 7 valence e-

(d) Elements of groups IB, IIB, IIIB (only Sc, Y, La and Ac), IVB, VB, VIB,
VIIB and VIIIB have their two outermost shells incomplete.
These are called transition elements. These elements are placed
in the middle of the table. All these elements are metals.

(e) Elements of group zero have satisfied octet in their
outermost shell. These elements are called noble gases. These
are placed at the extreme right of the table.

(f) Two groups of 14 elements lying in group IIIB 58Ce to 71Lu
and 90
Th to 103Lw have their three outermost shells incomplete.
These are called lanthanides and actinides respectively and
have been placed at the bottom of the table.

Main periods in periodic table
1
4
5
6
7
8
9
2
3
The rows are known as PERIODS. There are 9 periods

Periods:
The horizontal rows shown in the periodic table are called periods or
simply rows. There are seven periods in the table.
(a) 1st period consists of 2 elements which are 1H and 2He
(b) 2nd and 3rd periods have 8 elements each.
2nd period → 3Li to 10Ne
3rd period → 11Na to 18Ar
Both these periods are called short periods.

(c) 4th and 5th periods have 18 elements each, while 6th period has
32 elements as shown below:
4th period  19K to 36Kr
5th period  37Rb to 54Xe
6th period  55Cs to 86Rn
All these three periods are called long periods. 6th period also includes 14
rare earths or lanthanides [58Ce to 71Lu]
(e) 7th period is an incomplete period and at present it consists of 23
elements which extend from 87Fr to 109Mt. All these elements of this
period are radioactive. This period also includes 14 actinides
[90
Th to 103Lr ].

1. Number of valence electrons and valency:
On moving down a given group, the number of valence electrons
does not change, i.e. remains the same.
The valencies of all the elements of the same group are the same.
Element
Atomic
No.
Electronic
configuration
Valence shell
electronic
configuration
Li 3 1s1 2s1 2s1
Na 11 1s2 2s2 2p6 3s1 3s1
K 19 1s2 2s2 2p6 3s2 3p6
3d10 4s1
4s1
Rb 37 1s2 2s2 2p6 3s2 3p6
3d10 4s2 4p6 5s1
5s1

General Characteristics of Groups
2. Properties of elements:
All the elements of a given group possess very similar physical
and chemical properties. There is a regular gradation in
their properties when we move from top to bottom in a group.
For example:
(a) The alkali metals (group IA) resemble each other and their
base-forming tendency increases from Li to Cs.
(b) The reactivity of halogens (group VIIA) decreases as we
pass from F to I.

3. Size of atoms:
Size of atoms increases on descending a group. For example in group
IA, atomic size increases from Li to Cs. Thus:
Li<Na<K<Rb<Cs
4. Metallic character:
The metallic character of the elements increases in moving from top to
bottom in a group.
This is particularly apparent in groups IVA, VA and VIA, which
begin with nonmetals (namely C, N and O respectively), and end
with metals (namely Pb, Bi and Po respectively).

For example, in group VA, N and P are non-metals, As and Sb are
metalloids and Bi is a typical metal.
Thus the metallic character of these elements increases from N to Bi
as shown below:
Elements of group VA : N, P As, Sb Bi
Non-metals Metalloids Metal
Metallic character: Metallic character increasing →
It is because of a gradual increase of the metallic character of the
elements from top to bottom that the oxides of the elements
become more and more basic in the same direction.
For example:
Oxides of the elements of group VA-
N2O3, P2O5 As2O3, Sb2O3 Bi2O3
Acidic Amphoteric Basic

5. Number of electron shells:
In going down a group the number of electron shells increases by
one at each step and ultimately becomes equal to the number
of the period to which the element belongs as shown below
for the elements of Group IA.
Elements Electronic
configuration
Number of shells
Li (3) 2, 1 2
Na (11) 2, 8, 1 3
K(19) 2, 8, 8, 1 4
Rb(37) 2,8.18,8,1 5
Cs(55) 2, 8, 18, 18, 8, 1 6
Fr(87) 2, 8, 18, 32, 18, 8, 1 7

 Gr-1 (Period 2 )
 Each element in this period has two
regions of space
around it for
electrons

 Each element in this period has 3
regions of space
around it for
electrons

regions of space
around it for
electrons

regions of space
around it for
electrons
The Lanthanides are part of period 6

 Gr-1 (Period 7)
regions of space
around it for
electrons
The Actinides are part of period 7

1. Number of valence electrons:
Number of valence electrons increases from 1 to 8 when we
proceed from left to right in a period.

General characteristics of periods
2. Valency:
The valency of the elements with respect to hydrogen in each short period
increases from 1 to 4 then falls to one while the same with respect
to oxygen increases from 1 to 7 as shown below for the elements of 2nd
and 3rd period-
Elements of 2nd
period
Li Be B C N O F
Hydrides of the
elements
LiH BeH2 BH3 CH4 NH3 H2O HF
Valency of the
elements with
respect to H
1 2 3 4 3 2 1
Elements of 3rd
period
Na Mg Al Si P S Cl
Oxides of the
elements
Na2O MgO Al2O3 SiO2 P2O5 SO3 Cl2O7
Valency of the
elements with
respect to oxygen
1 2 3 4 5 6 7

3. Size of atoms:
Size of atoms decreases from left to right in a period. Thus alkali metals
have the largest size while the halogens have the smallest size.
4. Properties of elements:
 The properties of the elements of a given period differ considerably.
 But the elements in the two adjacent periods show marked similarity
between them. For example, when we consider the elements of 2nd and
3rd periods, we find that Na resembles Li, Mg resembles Be, Al
resembles B, Si resembles C, P resembles N, S resembles O, Cl
resembles F and Ar resembles Ne.

5. Metallic character:
On moving from left to right in a period the metallic character of the elements
decreases. For example in 3rd period, Na, Mg and Al are metals while Si, P,
S and Cl are non metals as shown below:
Elements of 3rd period: Na, Mg, Al Si, P, S, Cl
Metals Non-metals
Metallic character: Metallic character decreasing →
6. Acidity and Alkalinity:
It is because of the gradual decrease of the metallic character from left to right
that the oxides of the elements become more and more acidic in the same
direction. For example:
Oxides of
elements
of 3rd
period
Na2O MgO Al2O3 SiO2 P2O5 SO3 Cl2O7
Strogly
basic
Basic Amphoteric Freely
acidic
acidic More
acidic
Most
acidic
Basic character decreasing

7. Number of shells:
In going from left to right, in a period the number of electron shells remains the
same and the number of a period corresponds to the number of the
shells found in the elements of that period.
e.g. all the elements of 2nd period have the electrons only in first two shells
Elements of
2nd period Li Be B C N O F Ne
Atomic
number 3 4 5 6 7 8 9 10
Electronic
configuration
1s2
2s1
1s2
2s2
1s2
2s22p1
1s2
2s22p2
1s2
2s22p3
1s2
2s22p4
1s2
2s22p5
1s2
2s22p6
Shell number 2 2 2 2 2 2 2 2

8. Diagonal relationship:
Sometimes, an element in the periodic table shows similarity of properties
with another element of the next group and next period diagonally. Such
types of relationship between two elements are known as diagonal
relationship. The following are the important examples of diagonal
relationship found in the periodic table.
i) Li-Mg diagonal relationship
ii) Be-Al relationship
iii) B-Si relationship
Diagonal relationship is the resemblance of the properties of the elements
of 2nd period with their diagonally opposite members lying in 3rd period.
Group IA IIA IIIA IVA
2nd period Li Be B C
3rd period Na Mg Al Si

Families on the Periodic Table
• Columns are also grouped into
families.
• Families may be one column, or
several columns put together.
• Families have names rather
than numbers. (Just like your
family has a common last
name.)

• Columns of elements are called
groups or families.
• Elements in each family have
similar but not identical
properties.
• For example, lithium(Li),
sodium(Na), potassium(K),
and other members of family
IA are all soft, white, shiny
metals.
• All elements in a familyhave
the same number of valence
electrons.
• Each horizontal row of
elements is calleda period.
• The elements in a periodare
not alike in properties.
• In fact, the properties
change greatly across even
given row.
• The first element in a period
is always an extremely
active solid. The last
element in a period, is
always an inactive gas.

Sizes of Atoms
The bonding atomic
radius is defined as
one-half of the
distance between
covalently bonded
nuclei.

Ionic Radius Trend
 Metals – lose e-, which means more p+ than e-
(more attraction) SO…
Ionic Radius < Neutral Atomic Radius
 Nonmetals – gain e-, which means more e-
than p+ (not as much attraction) SO…
Ionic Radius > Neutral Atomic Radius

Sizes of Ions
Ionic size depends
upon:
 Nuclear charge.
 Number of
electrons.
 Orbitals in which
electrons reside.
 Metals – lose e-,
which means more
p+ than e- (more
attraction) SO…
Ionic Radius <
Neutral Atomic
Radius

Sizes of Ions
 Ions increase in size
as you go down a
column.
 Due to increasing
value of n.

Metals versus Nonmetals
 Metals tend to form cations.
 Nonmetals tend to form anions.

Sizes of Ions
 Cations are
smaller than their
parent atoms.
 The outermost
electron is
removed and
repulsions are
reduced.

Cation Formation
11p+
Na atom
1 valence electron
Valence e- lost in
ion formation
Effective nuclear
charge on remaining
electrons increases.
Remaining e- are
pulled in closer to
the nucleus. Ionic
size decreases.
Result: a smaller sodium cation,
Na+

Sizes of Ions
 Anions are larger
than their parent
atoms.
 Electrons are
added and
repulsions are
increased.
 Nonmetals – gain
e-, which means
more e- than p+
(not as much
attraction) SO…
 Ionic Radius
> Neutral Atomic
Radi

Anion Formation
17
p+
Chlorine
atom with 7
valence e-
One e- is added
to the outer
shell.
Effective nuclear charge is
reduced and the e- cloud
expands.
A chloride ion is
produced. It is
larger than the
original atom.

The elements displayed on the Periodic Table
are classified as:
Metalloids Non-metals
Alkali metals Halogens
Alkaline earth
metals
Noble gases
Transition metals Rare earth
elementsOther metals
Classification of elements on the
Periodic Table

Groups in the periodic table
1 4 5 6 7 8 9
10 11 12 13 14 15 16 17 18
2 3
Some GROUPS and PERIODS have
other common names
Group 1 metals are
known as the
ALKALI metals
Group 2 metals are known
as the ALKALINE earth
METALS
Groups 3 to 12 are
known as the
TRANSITION metals
Group 17 is
known as the
HALOGENSThe elements
touching this
staircase are
known as the
METTALOIDS
Group 18 is
known as the
NOBLE
GASES
Period 8 is known
as the
LANTHANIDE
metals
8 Period 9 is known as
the ACTINIDE metals
9

Alkalai
Metals
Group IA
Sodium

Group IIA
Alkaline Earth
Metals
Calcium

Lanthanides & Actinides
Lanthanide
Actinide

Group 8A
Noble Gases
K r y p t o n

Lead (Pb) is found in what family?
A. Transition Elements
B. Alkalai Metals
C. Noble Gases
D. Alkaline Earth Metals
A. Transition Elements

Group 1A is also called?
A. Metalloids
B. Halogens
C. Alkalai Metals
D. Transition Elements
C. Alkalai Metals

Which of the following set of
elements are all noble gases?
A. Cl, Fl, & Ne
B. Ar, Ne, & He
C. Be, B, & Ar
D. He, Kr, & Mg
B. Ar, Ne, & He

All Actinides are…
A. Alkaline Earth
Metals
B. Group IA Elements
C. Radioactive
D. Halogens
C. Radioactive

Calcium (Ca) & Magnesium (Mg)
are both found in this family.
A. Lanthanides
B. Noble Gases
C. Alkaline Earth Metals
D. Group 8A
C. Alkaline Earth Metals

Metalloids
 The 5 elements classified as “Metalloids” are located in Groups III A,
IV A, V A and VI A of the Periodic Table.
 These elements have properties of both metals and non-metals.
 Some are semiconductors and can carry an electrical charge making
them useful in calculators and computers.
 The metalloids are:
Boron (B-5) [Group IIIA] - Selenium (Se-34) [Group-VIA]
Silicon (Si-14) [Group IVA] - Tellurium (Te-52) [Group-VIA]
Arsenic (As-33) [Group VA]

Alkali Metals
 The 6 elements classified as “Alkali Metals” are located in
Group IA of the Periodic Table.
 The elements contain 1 electron in their outer most orbital.
 These elements are collectively called alkali metals since they
form strongly alkaline oxides and hydroxides.
 Alkali metals are very reactive metals that do not occur freely
in nature.
 They are soft, shiny, malleable, ductile, and good
conductors of heat and electricity.
 Fr is a radioactive element.

Alkaline earth Metals
 The 6 elements classified as “Alkaline Earth Metals” are located
in Group IIA of the Periodic Table.
 The elements contain 2 electrons in their outer most orbital.
 The oxides of the three metals i.e. Ca, Sr and Ba were known
much earlier than the metals themselves and were called
alkaline earths since they were alkaline in character and
occurred in nature as earths [lime (CaO), strontia (SrO) and
baryta (BaO)]. Later, when Ca, Sr and Ba were discovered.
 Alkaline Earth Metals are all found in the Earth’s crust but not
in the elemental form as they are so reactive.

Alkaline earth Metals
 These elements are very reactive, but less reactive than alkali
metals
 Instead, they are widely distributed in rock structures.
 Although Ra has similar properties as alkaline earth metals, it is
a radioactive element.

Transition
Metals
 The elements classified as
“Transition Metals” are located in
Groups IB to VII B and VIII of
the Periodic Table.
 The d-block elements are called
transition elements because they
exhibit transitional behavior
between highly reactive ionic
compound and mainly covalent
compound on the other side.

Transition Metals
 Transition Metals are ductile, malleable, and conduct electricity
and heat. The Transition metals are-

Other Metals
 The 10 elements classified as “Other Metals” are located in Groups
III A, IV A, V A and VI A of the Periodic Table.
 All of these elements are solid, have a relatively high density and are
opaque.
 The “Other Metals” are:
Group III A Group IV A Group V A Group VI A
Aluminum (Al-13) Germanium (Ge-32) Antimony (Sb-51) Polonium (84)
Gallium (Ga-31) Tin (Sn-50) Bismuth (Bi-83)
Indium (In-49) Lead (Pb-82)
Thallium (Ti-81)

Non-Metals
 The 7 elements classified as “Non Metals” are located in Groups I A,
IV A, V A and VI A of the Periodic Table.
 Non-metals are not easily able to conduct electricity or heat and do
not reflect light.
 Non-metallic elements are very brittle and cannot be rolled into wires
or pounded into sheets.
 They exist as solids (such as carbon) and gases (such as oxygen) at
room temperature .
Group 1 Group IV A Group V A Group VI A
Hydrogen (H-1) Carbon (C-6) Nitrogen (N-7) Oxygen (O-8)
Phosphorous (P-15) Sulfur (S-16)

Halogens
 The 5 elements classified as “Halogens” are located in
Group 7 of the Periodic Table.
 These elements have 7 electrons in their outer shell.
 The elements are highly reactive and are poor conductors
of heat & electricity.

Halogens
 The term “Halogen” means “salt-former” and
compounds containing halogens are called “salts”.
 The halogens exist, at room temperature, in all three
states of matter -gases such as fluorine & chlorine,
solids such as iodine and astatine and liquid as bromine.
 Astatine is an unstable element of radioactive origin.

Noble gases or Inert Gases
 The 6 elements classified as “Noble Gases” are
located in Group zero of the Periodic Table.
 The outermost orbit in all these elements has
satisfied octet.
 Therefore, these elements are chemically inert.
2He
10Ne
18Ar
36Kr
54Xe
86Rn

Rare Earth Elements
 The elements classified as “Rare Earth Elements” are located in
Group III B of the Periodic Table and in the 6th and 7th periods.
 The Rare Earth Elements are of the Lanthanide and Actinide
series.
 They are hardly being found in earth.
 Most of the elements in the Actinide series are synthetic or man-
made.

Lanthanide Elements Actinide Elements
Lanthanum (La-57)
Cerium (Ce-58)
Praseodymium (Pr-59)
Neodymium (Nd-60)
Promethium (Pm-61)
Samarium (Sm-62)
Europium (Eu-63)
Gadolinium (Gd-64)
Terbium (Tb-65)
Dysprosium (Dy-66)
Holmium (Ho-67)
Erbium (Er-68)
Thulium (Tm-69)
Ytterbium (Yb-70)
Lutetium (Lu-71)
Actinium (Ac-89)
Thorium (Th-90)
Protactinium (Pa-91)
Uranium (U-92)
Neptunium (Np-93)
Plutonium (Pu-94)
Americium (Am-95)
Curium (Cm-96)
Berkelium (Bk-97)
Californium (Cf-98)
Einsteinium (Es-99)
Fermium (Fm-100)
Mendelevium (Md-101)
Nobelium (No-102)
Lawrencium (Lr-103)
Rare Earth Elements

The Periodic Table
MetalsTransition MetalsRare Earth MetalsMetalloidsNon-Metals

Organization of Periodic Table
Alkali MetalsAlkaline Earth MetalsHalogensNobel Gasses

According to the electronic configurations, the elements may
be divided into four types such as:
1. The Inert Gases (Elements of 0 group).
2. The Representative Elements
(s and p block elements).
3. The Transition Elements (d block elements).
4. The Inner Transition Elements (f block elements).
Classification of elements on the basis of
electronic configuration

Classification of elements on the basis of electronic configuration
The Inert Gases:
 The noble or inert gases (zero group elements) have been placed at the
end of each period in the periodic Table. It appears that all these
elements have satisfied octet in their outermost orbitals.
 Helium has 2s2 stable arrangement and all other inert gases have s2p6
outer configurations:
He (2) = 1s2
Ne (10) = 1s2 2s2 2p6
Ar (18) = 1s2 2s2 2p6 3s2 3p6

 The configuration shows duplet and octet in the outermost energy
levels. Therefore, they are chemically inert. As a result, their valencies
are zero.
 Therefore, the position of the noble gases should be in zero groups.
 It may be noted that no atom has a complete energy level except helium
and neon.
 These elements are colorless gases.

The Representative Elements (s and p block elements):
 These elements generally belong to A sub-group of
the Periodic Table.
 s-block elements: The elements in which the last
electron(s) enters the s-orbital of their outermost
energy layer are called s-block elements.
 Thus the alkali metals (Group IA), alkaline earth
metals (Group IIA) are s block or s orbital elements.
Example: Na (11) – 1s2 2s2 2p6 3s1

 p-block elements: The elements in which the last electron(s)
enters to the p-orbital of their outermost energy layer are called
p-block elements.
 The valence electrons of all the elements
from boron to halogens
(groups IIIA to VIIA vertically) occupy
p orbitals. Hence these elements are
called p block or p orbital elements.
Example:
Al (13) : 1s2 2s2 2p6 3s2 3p1
Cl (17) : 1s2 2s2 2p6 3s2 3p5

Classification of elements on the basis of electronic
configuration
The Transition Elements (d block elements):
 The elements in which the last electron(s) enters to the d-orbital
which is inner to the outer-most shell are called d-block elements.
 The elements of group VIII and sub-group B are generally the d-
block elements.
 These elements contain two incomplete energy levels because of the
building up of the inner d electrons.
Example: Sc (21) : 1s2 2s2 2p6 3s2 3p6 3d1 4s2
Fe (26) : 1s2 2s2 2p6 3s2 3p6 3d6 4s2

Classification of elements on the basis of
electronic configuration
 Elements which have normally the same number of electrons in the
outermost level but have a progressively greater number of
electrons in an inner level (such as d level) are called “Transition
Elements”.
 In the Periodic Table we come across four such transition series in which
the additional electrons enter the 3d, 4d, 5d and 6d orbitals.

Classification of elements on the basis of electronic
configuration
List of Transition Elements (d block elements):

 The first transition series of elements involving the completion of 3d
level start from Sc (21) to Zn (30).
 The second series of transition elements start from Y (39) up to Cd
(48) involving 4d energy level.
 The third group of transition metals starts from La (57) but with a
break from Ce (58) to Lu (71) which are classified as inner
transition metals and proceed up to Hg (80) involving 5d energy
level.

Properties of Transition Elements:
a) All the elements are of high melting points, electropositive and heavy
metals.
b) These metals have almost the same atomic and ionic sizes. There is only
slight increase in the ionization energy of the formation of M+2 ions.
c) All these elements show positive oxidation states of +2 and +3 generally
and form mostly ionic compounds. Higher oxidation states are also
exhibited in some cases.
d) As a general rule, the transition elements form colored compounds.
e) These elements are also effective catalytic agents.
f) All these form quite a large number of complex compounds.
These properties are due to the influence of the incomplete inner d orbitals in
the transition elements. The properties are similar in the case of inner
transition elements where f orbitals are being completed.

The inner transition elements (f block elements):
 The elements in which the last electron(s) enters into the (n-2)f-orbital
are called f-block elements.
 These elements are located in group IIIB and have three incomplete
outer levels.
 Since (n-2)f orbital lies comparatively deep within the kernel (being
inner to the penultimate shell), these elements are also called inner-
transition elements.
 The f-block elements consist of two series of elements which are placed
in two rows at the bottom of the periodic table.

 The first series of 14 elements (atomic numbers 58 to 71) in
which 4f level is being build up follows lanthanum (57) and
are called Lanthanides.
 Example:
Ce (58) → 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f1 5d1
→ 1s2, 2s2p6, 3s2p6d10, 4s2p6d10f1, 5s2p6d1, 6s2
→ 2, 8, 18, 19, 9, 2

 Another series of 14 elements (atomic numbers 90 to 103) in which 5f
level is being filled follows actinium and is known as Actinides.
Example:
Th (90) → 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6
7s2 5f1 6d1
→ 1s2, 2s2p6, 3s2p6d10, 4s2p6d10f14, 5s2p6d10f1, 6s2p6d1, 7s2
→ 2, 8, 18, 32, 19, 9, 2
The inner transition elements (lanthanides and actinides) are all metals
and show variable oxidation states. Their compounds are highly
colored.

Although the Mendeleef’s Periodic Table was the first successful attempts
for the classification of elements but it suffers from the following defects:
(i) Position of hydrogen: Hydrogen resembles both the alkali metals
(group IA and halogen (group VIIA) in properties. Therefore, its position
in the periodic table is anomalous.
(ii) Position of lanthanides and actinides: A group of 15 elements (At.
No. 57 to 71) which is called rare earths or lanthanides does not find its
proper place in the table and has been placed at one place in group III
and period 6. Similarly, another group of 15 elements (At. No. 89 to
103) called actinides does not find its proper place and has been put at
one place in group III and period 6.
Defects of Mendeleef’s Periodic Table

(iii) Existence of four anomalous pairs of elements.: The order of
increasing atomic weight has been ignored in case of four pairs of
elements in order to place them in a position justified by their properties.
Thus elements of higher atomic weights precede those of lower atomic
weight at four places as shown below
(iv)
Elements with similar properties like Cu
and Hg, Ag and Th, Ba and Pb are separated while dissimilar elements
like Cu, Ag, and Au are grouped along with the alkali metals. Mn is
grouped with the halogens.
(a)Ar (Z=18 at. wt.=40) Proceeds K (Z=19, at. wt. =39.0)
(a)Co (Z=27 at. wt.=59.9) Proceeds Ni (Z=28, at. wt. =58.6)
(a)Te (Z=52 at. wt.=127.6) Proceeds I (Z=53, at. wt. =126.9)
(a)Th (Z =90 at. wt.=232.2) Proceeds Pa (Z=91, at. wt. =231)

Defects of Mendeleef’s Periodic Table
(v) Position of isotopes: If the elements are arranged in the order of their
increasing atomic weights, it is not possible to accommodate large
number of isotopes in the periodic table.
(vi) Group does not represent valency: Excepting osmium, elements
placed in group eight do not show a valency of 8. Also the elements
lying in the middle of long periods show two or more valencies.
e.g. Cr, Mn etc.

With the replacement of atomic weight by atomic number as the basis
of classification of elements, many of the irregularities in the
Mendeleef's table disappear as shown below:
1. Position of hydrogen.
The dual role of hydrogen is explained by the fact that it has one
electron in its outer orbit. It has equal tendency of gaining or losing one
electron for assuming a stable configuration. When it loses one electron
to give H+, it resembles alkali metals (which give Li+, Na+, K+, etc. ions)
while when it gains one electron to give H–, it resembles halogens
(which give Cl–, Br– etc.).

2. Anomalous pairs of elements.
This anomaly disappears altogether and the pairs
Ar—K, Co—Ni, Te—I and Th—Pa are found
arranged in the table in the order in increasing
atomic numbers as shown below:
Pairs of elements Ar K Co Ni Th I Ta Pa
Atomic numbers 18 19 27 28 52 53 90 91
Atomic weights 40 39 59.9 58.6 127.6 126.9 232.12 231

3. Position of rare earths.
The arrangement of extra nuclear electrons in all the
rare earth elements can be represented as 2, 8, 18
(18 + x), 9, 2, where x varies from 0 (for La) to 14
(for Lu). With this general arrangement of electrons,
all of them possess the same valency and similar
chemical properties. This justifies their grouping at
the same place.
4. Position of isotopes.
Since isotopes of the same element possess the
same atomic number, all of them should occupy one
and the same place in the periodic table.

5. Justification for dissimilar elements
being placed together.
The length of the periods is determined by arrangement of
electrons in different orbits. The end of every period results
from the completion of the last orbit (last number is always an
inert gas). Different periods carry 2, 8, 18 and 32 elements.
When 18 elements are to be distributed among 8 groups;
groups 1-7 get two elements each while group 0 gets only
one. The three elements which cannot be arranged elsewhere
are placed in a special group VIII. This lack of space is
enough justification for group VIII.
Out of the two elements which every long period adds to a
group, one resembles the typical element, the other does not.
This gives rise to the formation of subgroups. This explains
why dissimilar elements have been grouped together.

The long form of periodic table has a number of merits over the
Mendeleef's periodic table in the following respects:
(1) The classification of the elements is based on a more
fundamental property viz., atomic number.
(2) Each group contains elements with similar electronic
configuration and hence similar properties. For example, all the
alkali metals have similar valence-shell electronic configuration
viz. ns1 configuration and hence have similar properties.

(3) It explains the similarities and variations in the properties
of the elements in terms of their electronic configurations and
brings out clearly the trends in chemical properties across the
long periods.
(4) The inert gases having completely filled electron shells
have been placed at the end of each period.
(5) In this form of the periodic table, the elements of the two
sub-groups have been placed separately and thus dissimilar
elements do not fall together.
(6) It provides a clear demarcation of different types of the
elements like active metals, transition metals, non-metals,
metalloids, inert gases, lanthanides and actinides.
(7) It is easier to remember, understand and reproduce.

Although the long form of the Periodic Table is superior to Mendeleeff’s
Periodic Table in many respects still it retains some defects and
limitations.
(i) Position of hydrogen: Hydrogen resembles both the alkali metals
(group IA and halogen (group VIIA) in properties. The problem of the
position of hydrogen still remains unsolved.
(ii) Position of lanthanides and actinides: A group of 14 elements in
each of the lanthanide series and actinide series have been shown
separately below the periodic table. Like the Mendeleeff’s table, it fails
to accommodate the lanthanides and actinides in the main body of the
table.

Defects of Long form of (Modern) Periodic Table
(iii) Inconsistency in group VIII: In group VIII, three elements have
been placed one after another in the same period. Such as Fe (26), Co
(27) and Ni (28) have been placed in the same group (VIII). This
arrangement is against the general pattern of the periodic table.
(iv) Similar elements are separated and dissimilar elements are
placed in the same group: Certain elements which possess similar
properties are separated in the periodic table, for example Cu and Hg,
Ag and Tl, while many dissimilar elements have been grouped together.
For example, Cu, Ag and Au are grouped along with the alkali metals
though there is little resemblance between them.

Defects of Long Form of Periodic Table
Although the long form of the periodic table is superior
to Mandeleef's periodic table in many respects, it
retains some of the defects as such.
(i) The problem of the position of hydrogen still
remains unsolved.
(ii) It fails to accommodate the lathanides and
actinides in the main body of the table.
(iii) The arrangement is unable to reflect the electronic
configuration of many elements.

Why is the Periodic Table important
to me?
• The periodic table is the
most useful tool to a
chemist.
• You get to use it on
every test.
• It organizes lots of
information about all the
known elements.

Pre-Periodic Table Chemistry …
• …was a mess!!!
• No organization of
elements.
• Imagine going to a grocery
store with no organization!!
• Difficult to find information.
• Chemistry didn’t make
sense.

 Classification of elements
The classification of elements of similar
properties into groups, simplified their study.
For example Na a member of alkali metals ,
reacts with water vigorously giving hydrogen
gas and forming NaOH, which is a strong
base. The other alkali metals also react with
water in a similar manner.

 Prediction of undiscovered elements
At present all the elements from atomic number 1 to
109 have been discovered and their properties are
more or less known. But a very remarkable use of the
periodic table was made by the Mendeleev in
predicting a number of undiscovered elements, which
were shown by a number of gaps in the periodic table.
Mendeleev’s table contained only 65 elements with a
large number of vacant places. Mendeleev predicted
the existence and properties of 6 elements
corresponding to the gaps. These elements have
since been discovered and are Sc, Gallium,
Germanium, Technium, Rhenium, and Polonium.

 Correction of atomic weight
Atomic weight of some of the elements at the time of
Mendeleev gave a wrong position in the periodic table.
The properties of these elements required their
placement somewhere else.
For instance the element
indium was placed in a
vacant place in the periodic
table between Cd (112.4)
and Sn (118.7) and indium
with weight of about 114
fitted very well in between
Cd and Sn

 Periodic table in
industrial research
The periodic table has been
found to be quite useful in
industrial researches. Several
of the light metals and their
alloys used in modern
mechanical equipment’s, jet
engines and air crafts were first
studied in detail because of
their position in the periodic
table.

References
 Le May, H., Beall, H., Robblee, K., & Brower, D. (1996).
Chemistry: Connections to Our Changing World. Upper Saddle
River: Prentice Hall.
 What Do You Know about the Periodic Table?. Discovery
Channel School (2004). Retrieved March 23, 2009, from
Discovery Education: http:streaming.discoveryeducation.com/
 Helmenstine, Todd (2008). Printable Periodic Table. Retrieved
March 23, 2009, from About.com: Chemistry Web site:
http://chemistry.about.com/b/2008/01/26/printable-periodic-
table.htm
 All images found on google images

Classification of Elements MANIK

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