Classification of elements manik

PHARM 1101
Md. Imran Nur Manik
Lecturer
Department of Pharmacy
Northern University Bangladesh

Classification of Elements
Prepared By: Md. Imran Nur Manik; B.Pharm; M.Pharm Page 1
manikrupharmacy@gmail.com; Lecturer; Department of Pharmacy; Northern University Bangladesh.
Elements
Science has come a long 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 Mendeleef 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 ‘would be apparent. 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 (i.e., similar elements) are
repeated after definite regular intervals.
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. Mendeleef arranged the elements in the form of a table which is known
as Mendeleeff’s Periodic Table after his name.
Mosley’s Modern Periodic Law:
With the advancement or the knowledge about atomic structure
and discovery of new elements, in 1913, Mosley, a British
physicist, predicted that most of the defects of Mendeleef’s
periodic table disappear, if the basis of classification of elements
is changed to atomic number in place of atomic weight.
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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 definite intervals.
The Modern Periodic Table (Extended or Long form of Periodic Table):
The original Periodic Table suggested by Mendeleef has undergone many
modifications to remove the defects of Mendeleef’s periodic table, although the basic
features have been maintained in all the modified forms. Out of the various tables the
Extended Long Form of Periodic Table which is based on the Aufbau principle
(building up of the atomic electronic configurations) is the most simple and is widely
accepted.
Characteristics of Modern Periodic table:
In the modern periodic table 110 elements have been arranged according to the
increasing atomic number.
The main two parts of the periodic table are
1. Groups
2. Period
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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 group consisting of’ three triads
(Fe, Co, Ni; Ru, Rh, Pd; Os, Ir, Pt) and zero groups of inert gases. 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.
(b) 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,
all non-metals and metalloids.
(c) Elements of groups IB, IIB, IIIB (only Sc, Y, La and Ac), IVB, VB, VIB, VIIB and
VIII 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.
(d) 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.
(e) Two groups of 14 elements lying in group IIIB [Ce (Z=58) to Lu (Z=71) and
Th (Z=90) to Lr (Z=103)] have their three outermost shells incomplete. These are
called Lanthanides and Actinides respectively and have been placed at the bottom of
the table.
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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 H (Z= 1) and He (Z=2).
(b) 2nd
and 3rd
periods have 8 elements each.
2nd
period → Li (Z=3) to Ne (Z=10)
3rd
period → Na (Z=11) to Ar (Z=18)
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 → K (Z=19) to Kr (Z=36)
5th
period → Rb (Z=37) to Xe (Z=54)
6th
period → Cs (Z=55) to Rn (Z=86)
All these three periods are called long periods. 6th period also includes 14 rare
earths or lanthanides [Ce (Z=58) to Lu (Z=71)].
(d) 7th period is an incomplete period and at present it consists of 24 elements which
are Fr (Z=87) to Ds (Z=110). All these elements of this period are radioactive. This
period also includes 14 actinides [Th (Z=90) to Lr (Z=103)].
General Characteristics of Groups:
1. Number of valency electrons: On moving down a given group the number of
valence electrons does not change, i.e. remains 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
2. Valency1
: The valencies of all the elements of the same group are the same.
3. 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.
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4. 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
5. 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 VI A,
which begin with non-metals (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:
increasingcharacterMetallic:characterMetallic
MetalMetalloidsmetals-Non
BiSbAs,PN,VAgroupofElements :
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-
BasicAmphotericAcidic
OBiOSb,OAsOPON 3232325232 ,
6. 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 No. 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
Li (3) → 1s2
, 2s1
Na (11) → 1s2
, 2s2
, 2p6
, 3s1
K (19) → 1s2
, 2s2
, 2p6
, 3s2
, 3p6
, 4s1
Rb (37) → 1s2
, 2s2
, 2p6
, 3s2
, 3p6
, 3d10
, 4s2
, 4p6
, 5s1
Cs (55) → 1s2
, 2s2
, 2p6
, 3s2
, 3p6
, 3d10
, 4s2
, 4p6
, 4d10
,5s2
, 5p6
, 6s1
Fr (87) → 1s2
, 2s2
, 2p6
, 3s2
, 3p6
, 3d10
, 4s2
, 4p6
, 4d10
, 4f14
, 5s2
, 5p6
, 5d10
, 6s2
, 6p6
, 7s1
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General characteristics of periods:
1. Number of valency electrons: Number of valency electrons increases from 1 to 8
when we proceed from left to right in a period.
2. Valency: The valency of the elements with respect to (w.r.t) 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 w.r.t. Hydrogen 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 w.r.t. 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.
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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:
decreasingcharacterMetallic:characterMetallic
metals-NonMetals
ClS,P,Si,AlMg,Na,:period3ofElements rd
6. Acidity and Alkalinity: The gradual decrease of the metallic character from left to
right shows that, the oxides of the elements become less and less basic in the same
direction. For example:
Oxides
of the
elements
of 3rd
period
Na2O MgO Al2O3 SiO2 P2O5 SO3 Cl2O7
Strongly
basic
Basic Amphoteric Freely
acidic
Acidic More
acidic
Most
acidic
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 as shown below:
Elements of 2nd
period Li Be B C N O F Ne
Atomic number 3 4 5 6 7 8 9 10
Electronic configuration 2, 1 2, 2 2, 3 2, 4 2, 5 2, 6 2, 7 2, 8
No. of shells 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 the 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.
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Classification of elements of the periodic Table:
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 elements
Other metals.
Metalloids: The 5 elements classified as “metalloids” are located in groups 13 (IIIA),
14 (IVA), 15 (VA) and 16 (VIA) of the periodic table.
These elements have properties of metals and non-metals. Some are semi-conductors
and can carry an electrical charge making them useful in calculators and computers.
The metalloids are:
 Boron (B-5) [group 13]
 Silicon (Si-14) [group 14]
 Arsenic (As-33) [group 15)
 Selenium (Se-34) [group 16]
 Tellurium (Te-52) [group 16]
 Alkali metals: The 6 elements classified as ―alkali metals‖ are located in
group 1 (IA) of the periodic table.
The alkali metals are: Lithium (Li-3) ; Sodium (Na-11) ; Potassium (K-19) ;
Rubidium (Rb-37) ; Cesium (Cs-55) ; Francium (Fr-87)
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 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 2 of the Periodic Table.
The Alkaline Earth Metals are: Beryllium (Be-4); Magnesium (Mg-12); Calcium (Ca-20)
Strontium (Sr-38); Barium (Ba-56); Radium (Ra-88).
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The oxides of the three metals viz., 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, they
were named alkaline earth metals. Now this term is used to include all the
elements of Group II A.
Alkaline Earth Metals are all found in the Earth’s crust, but not in the elemental
form as they are so reactive. 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 3- l2 (group VIII and subgroup B) of the Periodic Table. The
d-block elements are called transition elements because they exhibit transitional
behaviour between highly reactive ionic compound forming s-block elements
(electropositive elements) on one side and mainly covalent compound forming
p-block elements (electronegative elements) on the other side.
Transition Metals are ductile, malleable, and conduct electricity and heat.
The Transition metals are:
Scandium (Sc-21) Yttrium (Y-39)
Titanium (Ti-22) Zirconium (Zr-40)
Vanadium (V-23) Niobium (Nb-41)
Chromium (Cr-24) Molybdenum (Mo-42)
Manganese (Mn-25) Technetium (Tc-43)
Iron (Fe-26) Ruthenium (Ru-44)
Cobalt (Co-27) Rhodium (Rh-45)
Nickel (Ni-28) Palladium (Pd-46)
Copper (Cu-29) Silver (Ag-47)
Zinc (Zn-30) Cadmium (Cd-48)
Lanthanum (La-57) Actinium (Ac-89)
Hafnium (Hf-72) Rutherfordium (Rf-104)
Tantalum (Ta-73) Dubnium (Db-105)
Tungsten (W-74) Seaborgium (Sg-106)
Rhenium (Re-75) Bohrium (Bh-107)
Osmium (Os-76) Hassium (Hs-108)
Iridium (Ir-77) Meitnerium (Mt-109)
Platinum (Pt-78) Darmstadtium (Ds-110)
Gold (Au-79)
Mercury (80)
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 Rare earth elements: The elements classified as “Rare earth elements” are located
in group 3 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. The Lanthanide and Actinide series of Rare Earth Elements are:
Lanthanide elements Actinide elements
Lanthanum (La-57) Actinium (Ac-89)
Cerium (Ce-58) Thorium (Th-90)
Praseodymium (Pr-59) Protactinium (Pa-91)
Neodymium (Nd-60) Uranium (U-92)
Promethium (Pm-61) Neptunium (Np-93)
Samarium (Sm-62) Plutonium (Pu-94)
Europium (Eu-63) Americium (Am-95)
Gadolinium (Gd-64) Curium (Cm-96)
Terbium (Tb-65) Berkelium (Bk-97)
Dysprosium (Dy-66) Californium (Cf-98)
Holmium (Ho-67) Einsteinium (Es-99)
Erbium (Er-68) Fermium (Fm-100)
Thulium (Tm-69) Mendelevium (Md-101)
Ytterbium (Yb-70) Nobelium (No-102)
Lutetium (Lu-71) Lawrencium (Lr-103)
 Other metals: The 10 elements classified as ―Other metals‖ are located in Groups 13
(IIIA), 14 (IVA), 15 (VA) and 16(VIA) of the Periodic Table.
All of these elements are solid, have a relatively high density and are opaque. The
―Other Metals‖ are:
Group 13 (IIIA) Group 14 (IVA) Group 15 (VA) Group 16 (VIA)
Aluminum (Al-13) Germanium (Ge-32) Antimony (Sb-51) Polonium (Po-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 1 (IA),
14 (IVA), 15 (VA) and 16 (VIA) of the Periodic Table.
Group 1 (IA) Group 14 (IVA) Group 15 (VA) Group 16 (VIA)
Hydrogen (H-1) Carbon (C-6) Nitrogen (N-7) Oxygen (O-8)
Phosphorus (P-15) Sulfur (S-16)
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 at room temperature, in two of the three states of matter: gases (such
as oxygen) and solids (such as carbon).
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 Halogens: The 5 elements classified as ―Halogens‖ are located in Group 17 (VIIA) of
the Periodic Table.
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. The Halogens are:
Fluorine (F-9)
Chlorine (Cl-17)
Gas
Bromine (Br-35) Liquid
Iodine (I-53)
Astatine (At-85)
solid
Astatine was discovered in 1940 and is an unstable element of radioactive origin. The
other four elements are stable and resemble each other in physical and chemical
properties.
 Noble gases or inert Gases: The 6 elements classified as ―Noble Gases‖ are located
in Group 18 (0) of the Periodic Table.
The outermost orbit in all these elements is completely filled. Therefore, these elements
are chemically inert. The Noble Gases on the periodic Table are:
- Helium (He-2) - Argon (Ar-18) - Xenon (Xe-54)
- Neon (Ne-10) - Krypton (Kr-36) - Radon (Rn-86)
Classification of elements on the basis of electronic configuration
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).
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 s2
p6
outer
configurations:
He (2) = 1s2
Ne (10) = 1s2
2s2
2p6
Ar (18) = 1s2
2s2
2p6
3s2
3p6
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 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
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
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 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.
Fig. 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.
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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
, 2s2
p6
, 3s2
p6
d10
, 4s2
p6
d10
f1
, 5s2
p6
d1
, 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
, 2s2
p6
, 3s2
p6
d10
, 4s2
p6
d10
f14
, 5s2
p6
d10
f1
, 6s2
p6
d1
, 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.
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Defects of Mendeleef’s Periodic Table:
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.
These two groups do not find its proper place in the table and have been placed all
together separately.
(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
(a) Ar (Z=18 at. wt.=40) Proceeds K (Z=19, at. wt. =39.0)
(b)Co (Z=27 at. wt.=59.9) Proceeds Ni (Z=28, at. wt. =58.6)
(c) Te (Z=52 at. wt.=127.6) Proceeds I (Z=53, at. wt. =126.9)
(d)Th (Z =90 at. wt.=232.2) Proceeds Pa (Z=91, at. wt. =231)
(iv) Similar elements are separated and dissimilar elements are placed in the
same group: 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.
(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.
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Mosley's Modern Periodic Law (Characteristics)
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 extranuclear 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.
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Merits of Long Form of Periodic Table over Mendeleef's Periodic Table
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) It relates the position of an element to its electronic configuration. Thus 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. ns 1 configuration and hence have similar properties.
Alkali
Metals
Atomic
No.
Complete Electronic
configuration
configuration
Li 3 2, 1 2s1
Na 11 2, 8, 1 3s1
K 19 2, 8, 8, 1 4s1
Rb 37 2, 8, 18, 8 1 5s1
(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. Such a location of the inert gases represents a logical completion 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. The
elements of group IA and IIA are active metals and are located at the extreme left of the
table.
The transition metals are found in the middle of the table. The elements lying to the right
of a dark line shown in the long form of periodic table in the form of ladder (i.e. steps)
are noble gases, metalloids and non-metals while those lying to the left of the line are
metals (active metals and transition metals).
(7) It is easier to remember, understand and reproduce.
Defects of the modern periodic table
Some defects of the periodic table are as follows
The problem of placing hydrogen remains unsolved.
It fails to accommodate the lanthanides and actinides in the main body of the table.
The arrangement is unable to reflect the electronic configuration of many elements.
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Position of hydrogen in the periodic table:
Placing hydrogen in group IA:
Electronic configuration: Hydrogen has the same electronic configuration as the
alkali metals (group IA).
Element
Atomic
No.
Complete Electronic
configuration
configuration
Li 3 2, 1 2s1
Na 11 2, 8, 1 3s1
K 19 2, 8, 8, 1 4s1
Rb 37 2, 8, 18, 8 1 5s1
2. Valency: Hydrogen has the valence of ―1‖ like the alkali metals.
3. Electro positivity: Hydrogen is an electropositive element like the alkali metals.
4. Reducing property: Like the alkali metals it also acts as a reducing agent.
5. Formation of stable compound: Hydrogen forms stable compounds with oxygen
and halogen like H2O, HX similar to the compounds Na2O and NaX.
6. Formation of M+
ions: Like alkali metals Hydrogen is a strong electropositive
element and has the tendency to lose its electron to form a unipositive cation.
H–e–
H+
L–e–
Li+
7. Affinity for non-metals: Both hydrogen and alkali metals have a strong affinity for
non-metals and little affinity for mortals.
Placing hydrogen in group VIIA:
1. Hydrogen should be placed just before Helium. Therefore it should be placed in
group VIIA.
2. Valency: Hydrogen has the valence of similar to the halogens.
3. Non Metal: Hydrogen is a non-metal like halogens.
4. Atomic state: Like the halogens hydrogen is a diatomic gas.
5. Combination with non-metals: Hydrogen forms compounds like CH4 and SiH4 with
non-metals similar to Chlorine forming CCl4 and SiCl4.
6. Formation of Negative ions: Like halogens, Hydrogen also gains one electron to
form negative ion.
H+e–
H–
F+e–
F–
7. Formation of hydride: Hydrogen forms hydride with some metals NaH, CaH2, like
halogens form the halides e.g. NaCl, KBr.
Thus, there is a debate whether hydrogen should place group IA or VIIA. But since it is
an s-block element it is placed with the alkali metals.
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Usefulness of periodic table
The followings are the important aspects of periodic table.
1. 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.
2. 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.
3. 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.
4. 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.
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1: Definition of Valence
The term valence (or valency) is often used to state the potential or capacity of an
element to combine with other elements. At one time, it was useful to define valence of
an element as: the number of hydrogen atoms or twice the number of oxygen atoms
with which that element could combine in a binary compound (containing two
different elements only).
In hydrogen chloride (HCl), one atom of chlorine is combined with one atom of hydrogen
and the valence of chlorine is 1. In magnesium oxide (MgO), since one atom of
magnesium holds one atom of oxygen, the valence of magnesium is 2.
As already stated, there are three different types of bonds that are known to join atoms
in molecules.
Although no precise definition of valence is possible, we can say that: Valence is the
number of bonds formed by an atom in a molecule.
Periodicity of Properties and Magic Number
The repetition of the elements with similar properties at certain regular intervals of
atomic number in the periodic table is termed periodicity of properties. In order to
understand the concept of periodicity of properties we may consider the properties of
the elements of groups I A (Alkali metals), zero (Inert gases) and VII A (Halogens)
given below
The examination of the properties of these elements will show that the elements
belonging to the same group have similar properties, In other words we can say that the
atomic number intervals at which the elements with similar properties reappear are 2, 8,
8, 18, 18, 32 and 32, i.e. we have to pass 2, 8, 8, 18, 18, 32, and 32 elements before we
come across an element with similar properties. The numbers 2, 8, 8 and 32 are called
magic numbers.
Foot Notes
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Shielding or Screening Effect of inner-shell Electrons on the Valence
shell Electron.
The decrease in the attractive force exerted by the nucleus on the valence shell
electron, which is obviously due to the presence of the electrons lying between the
nucleus and valence-shell electrons, (called intervening electrons) is called shielding
effect or screening effect. In other words, the intervening electrons screen or shield the
valence-shell electrons from the nucleus.
Factors Affecting the Magnitude of Shielding Effect
Following are the important factors on which the magnitude of shielding effect caused
by the inner-shell electrons on the valence-shell electron depends.
(i) No. of inner-shell electrons or inner shells. Greater is the number of inner-shell
electrons or inner shells, greater is the magnitude of shielding effect caused by the inner
electrons on the valence- shell electron. Thus as we move down a group, the number of
inner-shells or inner shell electrons increases and hence the shielding effect also
increases. For example in the elements of group IA.
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Effective nuclear charge
Effective nuclear charge, Z is defined as the actual nuclear charge, Z minus the
screening effect caused by the electrons intervening between the nucleus and the
outer electrons. It is due to the shielding effect of the inner electrons on the outer-
electrons that the valence electron experiences less attractive pull from the nucleus. The
decrease in the attractive force reduces the nuclear charge, Z represented by the atomic
number of the element.
This decreased nuclear charge is called effective nuclear charge and is represented by
Zeff. It is given by the relation:
Why the size of atom aka atomic radius increases from top to
bottom while decreases from left to right?
The atomic size decreases across the period because effective nuclear charge on the
valance shall electrons increases. In a period the valence shell remains same for a
elements of a period. Let us consider example of second period elements the valance
shell for these elements is 2s and 2p orbital as electrons are being filled in these orbitals.
The electrons are filled in firstly 2s orbital and then a2p orbital is being filled. As
distance from nucleus and valance shell remains constant as shown in figure.
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The concentration of positive charge in the nucleus
increases with increase in atomic number. The nucleus
of the atom gains protons moving from left to right,
increasing the positive charge of the nucleus and
increasing the attractive force of the nucleus upon the
electrons. Although the electrons are also added as
the elements move from left to right across a period,
but these electrons reside in the same energy shell
and do not offer increased shielding. Thus, moving
from left to right across a period, the atomic radius
decreases.
So in fluorine there are nine protons that are attracting
the electrons but in Lithium there are only 3 protons that
are attracting valance shell. So size of fluorine would be
less as valance shell would contract due to strong
nuclear attraction. Thus the new size of the fluorine
would be less than expected.
Atomic size increases down a group
The atomic radius increases moving down a group.
Each level (shell) only certain number of electron can occupy.
If the number of electron further increased then, levels are
increased. The new energy shells provide shielding, allowing
the valence electrons to experience only a minimal amount of
the protons' positive charge.
It happens because each succeeding element has an
additional level or shell of electrons. Each new level is
shielded from the pull of the nucleus by the layers below it, so
it is further out from the nucleus, making the atom bigger.
In other words, the number of shells increases as we go down
the group. The outermost electrons are repelled by the inner shell electrons (Screening
effect) and hence, the atomic size increases.
Atomic volume
Atomic volume is defined as the volume in c.c, occupied by one gram atom of the
element in the solid state and hence is commonly called gram atomic volume. It is
obtained by dividing the atomic weight of the element by its density; i.e.
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Why the atomic volume increases from top to bottom while first
decreases then increases from left to right?
Variation of Atomic Volume in a Period and a Group.
(a) In a group. Atomic volume increases more or less regularly in going down a group
(See). The increase in atomic volume in going down a group is due to the increase in
the number of shells. The larger the number of shells, the bigger is the volume of the
atom.
(b) In a period.
In going from left to right in a period, it varies cyclically, i.e., it decreases at first for
some elements, becomes minimum in the middle and then increases (See).The
variation of atomic volume in going from left to right in a period is influenced by the
following two factors
(i) Nuclear charge. We know that the nuclear charge (i.e. atomic number) increases by
one, as we move from left to right in a period. The increased nuclear charge attracts
each electron more strongly towards the nucleus, resulting in a decrease in the volume of
the atom.
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(ii) No. of valence electrons.
Towards the close of a period, due to an increase in the number of valence-electrons (i.e.
electrons in the valence shell) the volume of the atom increases, so that it may
accommodate all the electrons. These two factors, one causing an increase and the other
causing a decrease, combine to result that in a period atomic volume decreases at first
for some elements, becomes minimum in the middle and then increase.
Electronegativity
In a molecule A – B the electrons forming the covalent bond are attracted by atom A as
well as by B. This attraction is measured in terms of what we call electronegativity,
EN.
It may be defined as: The attraction exerted by an atom on the electron pair bonding
it to another atom by a covalent bond.
Trend in electronegativity
The variations in electronegativities of elements in the Periodic table are similar to those
of ionisation energies and electron affinities.
(1) Increase across a Period
The values of electronegativities increase as we pass from left to right in a Period.
Thus for Period 2 we have
This is so because the attraction of bonding electrons by an atom increases with
increase of nuclear charge (At. No.) and decrease of atomic radius. Both these factors
operate as we move to the right in a Period.
(2) Decrease down a Group
The electronegativities of elements decrease from top to bottom in a Group.
Thus for Group VII we have
The decrease trend is explained by more shielding electrons and larger atomic radius as
we travel down a Group.
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Importance of Electronegativity
The electronegativities of elements are widely used throughout the study of Chemistry.
Their usefulness will be discussed at appropriate places. The important applications of
electronegativities are listed below.
(1) In predicting the polarity of a particular bond. The polarity of a bond, in turn,
shows the way how the bond would break when attacked by an organic reagent.
(2) In predicting the degree of ionic character of a covalent bond.
(3) In predicting of inductive effects in organic chemistry.
(4) In understanding the shapes of molecules.
Electron Affinity
A neutral atom can accept an electron to form negative ion. In this process, in general,
energy is released.
Electron affinity (EA) of an element is the amount of energy released when an
electron is added to a gaseous atom to form an anion.
Ionization Energy
Amount of energy required to remove an electron from the ground state of a gaseous
atom or ion.
First ionization energy is that energy required to remove first electron.
Second ionization energy is that energy required to remove second electron, etc.
Group Trend
As you go down a column, ionization energy decreases.
As you go down, atomic size is increasing (less attraction), so easier to remove an
e-
.
Periodic Trend
As you go across a period (L to R), ionization energy increases.
As you go L to R, atomic size is decreasing (more attraction), so more difficult to
remove an e-
(also, metals want to lose e-
, but nonmetals do not).
It requires more energy to remove each successive electron.
When all valence electrons have been removed, the ionization energy takes a
quantum leap.
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Trends in First Ionization Energies
As one goes down a column, less energy is required to remove the first electron.
For atoms in the same group, Zeff is essentially the same, but the valence electrons are
farther from the nucleus.
Md.
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