CHAPTER NO 6

PERIODIC CLASSIFICATION

Introduction
When a very large number of elements become known to scientists , it was felt that they must be arranged in a systematic order .Because a systematic classification provides a clear idea and information about the properties of elements and make it easier to predict the properties of different elements. This classification provides the inter-relationship of scientific facts.
In old days, element were arranged in the ascending order of their atomic masses.
The orderly arrangement of elements is called "PERIOD CLASSIFICATION".
DOBEREINER’S TRIADS
A group of three elements, which have similar physical and chemical properties, is called "TRIADS".
In 1829, a German scientist made use of the relationship between atomic masses and the properties
of elements. He proposed,
"If three elements are arrange in ascending order of their atomic masses, such that the
atomic mass of middle element is Arithmetic mean of the first and third elements, then these element will show similar properties".
This is known as "Law of Triads". This rule is applicable only in a few cases.
For example
TRIADS
ARITHMETIC MEAN
Result
Li (7), Na (23), K (39)
7+39/2=23
atomic mass of Na
Ca (40), Sr (87), Ba (137)
40+137/2=88
nearly equal to the
atomic mass of Sr
S (32), SC (79), Te (128)
32+128/2=80
nearly equal to the
atomic mass of Sc
Drawback:This rule is not valid for all the elements.
NEWLAND’S LAW OF OCTAVES
In 1866, a British scientist, Newland, reported his "law of octave" by arranging elements according to increasing order of their atomic masses. He notified that "Every eight element, starting form any point, approximately has similar properties".
Newland’s arrangement was applicable only a few elements after that it was failed.
ADVANTAGES OF THE LAW
1. This law provided a basis for the classification of element into groups of elements having similar properties.

2. This law provided a wider scope to arrange all known elements into a tabular form

DISADVANTAGES OF THE LAW
1. Newland’s law is not applicable to all the elements.
2. This arrangement did not include NOBLE GASES because they were not discovered then.
3. Heavier elements could not be accommodated
LOTHER MEYER’S CLASSIFICATION
In 1864, a German Chemist Lothar Meyer published on incomplete periodic table .He includes about 56 elements arranged in a group from I to VIII. He plotted the values of different physical properties and obtained different curves .In these graphs, he observed that element with similar physical properties occupy similar positions in the curve
For example
Alkali metals occupy the peak of the curves.
Halogen occur on the ascending portions of the curve.

MENDELEEV’S PERIODIC TABLE

In 1869, a Russian chemist, Mendeleev, on the basis of physical and chemical properties discovered a relation known as "PERIODIC LAW".
Mendeleev’s Periodic Law
According to the law:
"The properties of element are the periodic function of their atomic masses".
Mendeleev arranged the known elements according to increasing order of their atomic masses. According to him, fundamental property of an element was atomic mass. He arranged these elements
in eight groups. They were further divided into sub-groups.
Mendeleev’s periodic table was very precise and provided the basis of modern periodic classification.
ADVANTAGES OF MENDELEEV’S PERIODIC TABLE:
Mendeleev’s periodic table offered the following advantages in understanding the properties of elements.
1. There was a regular gradation in the physical and chemical properties of element.
2. The group number of an element indicates highest oxidation state that it can attain.
3. There were many vacant spaces in table for the elements to be discovered. According to his prediction, he named them Eka-Boron, Eka-Aluminium and Eka-silikon. He also predicted the properties of these undiscovered elements including atomic masses. These elements were discovered as Sc ,Ga and Ge with same features as he predicted.
4 .Mendeleev’s arrangement helped to correct atomic masses of a number of elements.
DISCREPANCIES IN MENDELEEV’S PERIODIC TABLE
1- For placing the elements in proper groups, the order of the elements according to atomic mass was reversed in certain cases. He placed Iodine (127) after Tellurium (128) Potassium (39) and Ni (58)
after Co (59). Which is against his periodic law but correct according to properties.

2- Mendeleev’s periodic table does not provide clear idea about the structure of atom.

3- Lanthanide and Actinide have been assigned placed in the periodic table which is against the periodic law.

4- Alkali metal and coinage metals (Cu, Ag and Au) which differ widely in properties are placed into the same group.

5. There was no separate position for isotopes in his periodic table.


6. The change in atomic mass of two successive elements is not constant. Hence it is not possible to predict the number of missing elements by knowing the atomic masses of two known elements.
PERIODIC CLASSIFICATION
MODERN PERIODIC LAW
   In 1913, an English physicist MOSELEY, as a result of his work on characteristics of X-rays of elements    discovered a fundamental and most accurate relation between chemical properties and atomic number    of elements, known as "Modern periodic law".
   ACCORDING TO THE MODERN PERIODIC LAW:
   1. The chemical and physical properties of elements are the periodic function of their atomic    numbers".

   2. The properties of elements depends upon their electronic configuration which vary with    increasing atomic number in a periodic way".
   The modern periodic law provides a logical and scientific ground for the classification of elements.
LONG FORM OF PERIODIC TABLE
   Periods
   Periodic table is divided into seven horizontal rows of elements. Each row of elements is called    "PERIODS".All the elements of a period are different from each other. Period no of an element    indicates number of orbits in the atom of that elements.
   General features of a period
   1- Period number of an element represents number of energy levels in the element. For example,
ELEMENTS
PERIOD
NUMBER OF ENERGY LEVELS
Na (Z=23)
K (Z=39)
Rb( Z=55)
Third
Forth
Fifth
3 (K , L , M)
4 (K , L , M , N)
5 (K , L , M , N , O)
   2- Atomic size decreases in a period.
   3- Nuclear charge increases in a period.
   4- Ionization potential increases in a period.
   5- Electronegativity increases in period.
   6- Electro positivity and metallic character decreases in period.
   7- Each period starts with Alkali metal and ends on a Noble gas.
   For example:
   2nd period : Li .................................. Ne
   3rd period : Na ..................................Ar
   7-First element of each period is most reactive (electropositive), and last element is chemically inert.
   DESCRIPTION OF PERIODS
First period

   It is the shortest period and it contains only two elements (H and He). This period corresponds to filling    up of K-shell. Electronic configuration of the elements of this period is 1S1 and 1S2
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Second period

   It is called first short period and it contains eight elements. This period corresponds to filling up of L-    shell. In this period electrons occupy 2s and 2p orbital. This period contain 2 element of S-block and    six p-block
Third period

   It also contains eight elements and is known as second short period. It corresponds to filling up of M-    shell. It contains 2 S-block and 6 P-block elements. In this period 3S and 3P orbital are being filled.
Fourth period
   It contains 18 elements and is known as first long period. This period corresponds to the filling up of N-    shell. It contains 2 S-block, 6 P-block and 10 d-block (transition) elements. It start with filling 4S    orbital followed by 3d and 4p orbital.

Fifth period
Fifth period contain 18 elements and is known as 2nd long period. It corresponds to filling up of O-shell.    It starts with the filling of 5S-orbital followed by 4d and 5p orbital.
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Sixth period
   Sixth period contains 32 elements and it is the longest period.
   Among 32 elements
   2 elements of s-block with electronic configuration 6s1 and 6s
   26 elements of p-block with electronic configuration 6s2, 6p1 to 6s2, 6p6 elements 10 of d-block with    electronic configuration 6s2, 5d1 to 6s2, 5d10
   14 elements of f-block with electronic configuration 6s2, 5d1, 4f1 to 6s2, 5d1, 4f14
    In sixth period electrons starts filling 4f-orbital after 5d-orbital accommodate one electron.
Seventh period
   7th period is incomplete period.
   It includes:
   2 elements of s-block with electronic configuration 7s1, 7s2
   10 elements of d-block with electronic configuration 7s2, 6d1 to 7s2, 6d10
   14 elements of f-block with electronic configuration 7s2, 6d1, 5f1 to 7s2, 6d1, 5f14
Groups
   Modern periodic table is divided into eight vertical columns of elements. Each vertical column of    elements is called a "GROUP". The groups of the periodic table are further divided into two sub-groups    or families "A and B". Elements of sub-group "A" are known as Normal elements or Representative    elements. While the elements of sub-group "B" are known as Transition elements. Group number of an    element is represented by roman numerals such as IA, IIA, VIA etc. Group number of an element    indicates its highest oxidation state. Atomic size increases down the group. I.P decreases down the    group.
   General features of a group are as under

   1-Group number of an element shows number of electrons in the outermost shell of that element
ELEMENT                   NUMBER OF VALENCE ELECTRONS                        GROUP
Na                                                                1                                                                I-A
Cl                                                                 7                                                                VII-A
C                                                                  4                                                                Iv-A


   2. ATOMIC SIZE: Atomic size of element increases from top to bottom in a group.

   3. NUMBER OF SHELL: from top to bottom number of shell increases by one for each element.

   4. NUMBER OF VALENCE ELECTRONS: No of valence electron are constant.

   5. IONIZATION POTENTIAL: I.P decreases from top to bottom.

   6. ELECTRONEGATIVITY: Electronegativity decreases from top to bottom.

   7. METALLIC CHARACTER : Metallic character increases down the groups
APPLICATIONS OF PERIODIC TABLE
   Arrangement of elements in the form of periodic table is very useful for the proper study of elements.    Some important applications of periodic table are as follows:

Classification of elements into periods and groups is very useful and easy in the study of chemistry.
Suggestions for further research become available.
Prediction of new element is possible.
GROUP I-A
  1. Li, Na, K, Rb, Cs, Fr are the elements of group I-A.
  2. All the elements of group I-A contain one electron in their outermost shell.
  3. They are univalent.
  4. Elements of group I-A are metals.
  5. They have the largest atomic size.
  6. They easily loose electron.
  7. They have low ionization potential (I.P) and are very electropositive.
  8. They form ionic compounds.
  9. They are very reactive element
  10. They are electropositive elements.
  11. They are called alkali metals.
  12. They are good reducing agents.
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GROUP II-A
  1. Be, Mg, Ca, Sr, Ba and Ra are the members of this group.
  2. They have 2 electrons in their outermost shell.
  3. All these elements are metals.
  4. They are also known as alkaline earth metals.
  5. They are reactive elements.
  6. They form ionic compounds by loosing two electrons.
  7. They are electropositive elements.
  8. They form ionic compounds.
  9. They are also good reducing agents.
GROUP III-
  1. B, Al, Ga, In and Tl are the elements of group III-A.
  2. They have three electrons in their outermost shell.
  3. Except Boron, all the elements are metals.
  4. Boron is a metalloid.
  5. They are electropositive elements.
  6. They are chemically reactive.
  7. They have tendency to form covalent compounds.
GROUP VI-A

   1. O, S, Se, Te and Po are the elements of group VI-A.
   2. They have six electrons in their outermost shell.
   3. Their oxidation number is (–2).
   4. They have high values of electronegativity.
   5. Oxygen and sulphur are non-metals, Se and Te are metalloids, but Po is a metal.
   6. Elements of group VI-A show allotropy.
   7. Oxygen is a gas, but other elements are solids.
   8. Generally they are electro-negative elements.
GROUP VII-A

F, Cl, Br, I and At, are the members of group VII-A.
All the elements of group VII-A contain seven electrons in their outermost shell.
Generally they are present in (-1) oxidation state.
Elements of group VII-A are non-metals.
They are highly electronegative elements.
They have high I.P.
They can form ionic and covalent bond (compound).
They are very reactive elements
They exist as diatomic molecule. For example F2, Cl2, Br2, I2, At2
They are also called halogens.

TRANSITION ELEMENTS

Elements from group I-B to group VIII-B are known as transition elements.
For example Fe, Cu, Ag, Ni etc.
They are metals.
They are hard.
They have high melting point.
They exhibit variable oxidation states. For example +2 , Fe+3, Cu +1, Cu+2
They are paramagnetic.
They form complex (coordination) compounds.
Most of the compounds of transition elements are colored.
For example : CuSO4 (blue) , NiSO4 (green) , CoCl3 (pink)
Transition elements and their compounds are used as a catalyst.
For example : Pt , Ni , FeCl3 , V2O5 , MnO2, FeO etc.
PERIODICITY

   Physical and chemical properties of elements in periodic table are periodically repeated. This repetition of    properties is called periodicity. Atomic size, ionization potential, electronegativity, metallic character,heat    of hydration etc. vary in a periodic manner known as periodicity.