You all are welcome to read this article. We shall learn Classification of Elements as a notes of periodic table 2024 in a new style. We have given proper attention towards current revised syllabus and included all the important topics of this chapter. This chapter is very useful for detail study of chemistry. This chapter includes all the essential terms applied in any faculty of chemistry, Hence we start with the following topics.
This chapter will help us to understand how to decide the position of all the elements. We shall also learn periodicity in properties of elements and anomalous behaviour of some elements in their groups.
READ MORE:- Notes of p-Block elements Group-16
A brief history of classification of elements:
Doebereiner’s Triads:
In this classification, similar elements in groups of three were arranged and showed that their atomic weights were either neatly the same or the atomic weight of the middle element was nearly the arithmetic mean of the other two. These groups of three elements were called Doebereiner’s Triads. This concept of triads could be applied only to a limited number of elements. Therefore, it was dismissed.
Telluric screw or helix:
This classification of elements was drafted by A.E.B. de Chancourtois. He arranged the elements in order of increasing atomic weights and made a cylindrical table of elements to display the periodic repetition of properties. He observed that the elements with similar properties fell on the same vertical lines drawn from the centre of the spiral. However, this also did not receive much attention.
Newlands’Law of Octaves:
In this classification of elements, the elements were arranged in order of their increasing atomic weights, the properties of every eighth element were similar to those of the first one like the eighth tone of a musical scale. This generalization was also discarded since it was applicable to only lighter elements having atomic weights up to 40 u.
Lothar-Meyer arrangement:
In this classification of elements, he plotted a graph between the atomic Volumes and atomic weights of the elements and observed that the elements with similar properties occupied similar positions on the curve. For example.
- The most strongly electro- positive alkali metals (Li, Na, K, Rb) occupy the peaks on the curve.
- The less strongly electro- negative alkaline earth metals (Be, Mg, Ca, Sr) occupy the descending positions on the curve.
- The most electronegative elements (F, Cl, Br, I) occupy the ascending positions on the curve.
Mendeleev’s periodic Law:
According to this law-
“The physical and chemical properties of elements are a periodic function of their atomic weights”.
Mendeleev’s periodic table: In this classification of elements, elements were arranged in tabular form in which elements of similar properties occupy the same position when arranged in increasing order of atomic weights. This repetition of properties of elements after certain regular intervals is called periodicity of properties.
Characteristics of Mendeleev’s periodic table:
This table was formulated in the following ways:
(1) There are nine vertical columns called groups. These are designated as 0, I, II, III, IV, V, VI, VII, VIII. Except for group 0 and VIII, each group is further divided into two sub-groups designated as A and B.
(2) There are six horizontal rows called series. These are numbered from 1 to 6.
(3) The elements which lie on the left hand side of each group constitute sub group A while those placed on the right hand side form sub group B.
(4) Group VIII contains nine elements in three sets each containing three elements.
(5) Group 0 has no subgroups. It contains only one vertical column of inert gases.
Significance of Mendeleev’s periodic table in classification of elements
Some important contributions of the Mendeleev’s periodic table for the classification of elements are followings:
(1) Systematic study of the elements. This periodic table simplified and systemized the study of the elements and their compounds since their properties could now be studied as groups or families rather than individuals.
(2) Prediction of new elements: During arranging the elements in its periodic table according to their properties, Mendeleev left some blank spaces or gaps and predicted the properties of some unknown elements on the basis of their positions. For example, gallium (eka-aluminium) and germanium (eka-silicon).
(3) Correction of doughtful atomic weights: This periodic table has helped in correcting the doughtful atomic weights of some elements by improving in their valencies.
Defects in the Mendeleev’s periodic table:
- Anomalous position of hydrogen: Hydrogen is placed in group IA. However, it resembles the elements of both the group IA (alkali metal) and group VIIA (halogens).
- Anomalous pairs of elements: Some elements with higher atomic weights precede the elements with lower atomic weights. For example, Ar (at.wt. 39.9) precedes K ((at.wt. 39.1).
- Position of isotopes. Isotopes of any element were not given any position in the periodic table.
- Some dissimilar elements are grouped together while some similar elements are placed in different groups. For example, alkali metals Li, Na, K, etc (group IA) are grouped together with coinage metals such as Cu, Ag and Au (group IB) though their properties are quite different. At the same time, certain chemically similar elements like Cu (group IB) and Hg (group IIB) have been placed in different groups.
- Position of elements of group VIII. No proper place has been allotted to nine elements of group VIII which have been arranged in three triads without any justification.
Modern periodic law: for the classification of elements.
Moseley in 1912, justified by his experiment that the square root of the frequency (v) of the prominent X-rays by a metal is proportional to the atomic number.
√v = a(z – b)
Hence, this law was given as:
The physical and chemical properties of the elements are a periodic function of their atomic numbers.
Modern periodic table: the classification of elements.
This table was designed just like Mendeleev’s periodic table but it was based on increasing order of atomic numbers. The position of elements in groups and periods were same.
Cause of periodicity:
In modern periodic table, elements were arranged in increasing order of atomic number that occurs a repeat of similar outer electronic Configuration after certain regular intervals. As a result elements show periodicity in properties.
Long form of periodic table for the classification of elements:
Long form of periodic table is based upon the electronic Configuration of elements and this table is also known as Bohr’s table.
Structural features of long form of periodic table: This table consists of eighteen vertical columns and seven horizontal rows. These have been arranged in order of increasing atomic numbers in such a way that the elements with similar electronic Configuration are placed under each other in the same vertical columns.
Groups: There are eighteen vertical columns in the long form of periodic table known as groups. These groups are numbered from 1 to 18.
Periods: There are seven horizontal rows in this table are known as periods. They are numbered from 1 to 7. The first period consists of two elements, second and third periods of eight elements, fourth and fifth periods of eighteen elements and the sixth period of 32 elements. The seventh period is incomplete but can have 32 elements. Read more the Classification of Elements.
Blocks : There are four blocks in the long form of periodic table. They are designated as s, p, d and f. s-block consists of two groups as 1 and 2, p-block of 6 groups from 13 to 18, and d-block of 10 groups from 3 to 12. f- block elements have been placed in the 3rd group of sixth and seventh periods as Lanthanoids and actinoids respectively.
IUPAC nomenclature of elements with atomic number > 100
The following rules are taken into consideration in the IUPAC nomenclature of the elements having atomic numbers greater than one hundred:
- The names are derived directly from the atomic numbers using numerical roots for 0 and numbers from 1– 9 and adding the suffix ‘ium’.
- In certain cases, the names are shortened. For example, bi + ium = bium, tri + ium = trium and enn + nil = ennil.
- The symbol of these elements are written using the first letter of the root of each digits of atomic number. For example, if atomic number is 100, its recommended name will be Unnilunium and symbol will be Unn.
General electronic configuration of block elements:
s-block elements: ns1-2, where n=1-7.
p-block elements: np1-6 ns2. where n=2-7.
d-block elements: (n-1)d1-10 ns0-2
where n=4-7.
f-block elements: (n-2)f0-14 (n-1)d0-2 ns2 where n=6-7.
How to determine the position of elements in periodic table:
To determine the position of elements in periodic table, we have to follow the following facts.
(1) The outermost shell containing electrons of any element indicates its period.
(2) The subshell in which last electron enters in electronic Configuration of any element indicates its block.
(3) Determination of group is done according to the block of the elements.
- For s-block elements, total electrons present in the outermost shell are equal to the group.
- For p-block elements, total electrons present in the outermost shell plus 10 are equal to the group.
- For d-block elements, total electrons present in (n-1)d and ns subshells are equal to the group.
- For f-block elements, the group number is always 3rd either of 6th period or 7th period.
Periodic trends in properties of elements:
Atomic radius:
The distance from the centre of the nucleus to the outermost shell containing electrons is called atomic radius.
Types of atomic radii. There are three types of atomic radii.
(a) Covalent radius: one-half the distance between the nuclei of two covalently bonded atoms of the same element in a molecule is called covalent radius.
(b) Van der Waals radius: one-half the distance between the nuclei of two identical non- bonded isolated atoms or two adjacent identical atoms belonging to two neighboring molecules of an element in the solid state is called van der Waals radius.
(c) Metallic radius: one-half the intermolecular distance between the two adjacent metal ions in the metallic lattice is called metallic radius.
Comparison of atomic radii: van der Waal radius > metallic radius > covalent radius.
Variation of atomic radii in the periodic table:
In general, the atomic radii decrease with increase in atomic number as we move from left to right in a period. The atomic radii of elements increase with increase in atomic number as we move from top to bottom in a group.
Ionic radii: in the classification of elements.
It may be defined as the effective distance from the centre of the nucleus of the ion upto which it exerts its influence on its electronic cloud.
Comparison of the Ionic radii with corresponding atomic radii:
(1) The radius of the cation is always smaller than of its parent atom. Due to decrease in the number of shells and increase in the effective nuclear charge. For example, Na+ < Na
(2) The radius of the anion is always larger than that of its parent atom mainly because of decrease in the effective nuclear charge. For example, Cl- > Cl
(3) Ions of the different elements which have the same number of electrons but different magnitude of the nuclear charge are called isoelectronic ions.
The Ionic radii of isoelectronic ions decrease with the increase in the magnitude of the nuclear charge. For example, Al3+ < Mg2+ < Na+ < F– < O2- < N3-.
In general like atomic radii, Ionic radii also increase down the group and decrease along the periods from left to right.
Ionisation enthalpy: During classification of elements
The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to convert into gaseous cation is called its Ionisation enthalpy. Its unit is eV or KJmol-1.
Successive Ionisation enthalpy:
The Ionisation enthalpies required to remove first, second, third etc. electrons from an isolated gaseous atom are called successive Ionisation enthalpy. The order of successive Ionisation enthalpy is ΔiH3 > ΔiH2 > ΔiH1
Factors affecting the Ionisation enthalpy:
(1) Nuclear charge: The Ionisation enthalpy increases with increase in nuclear charge.
(2) Atomic size or Radius: Ionisation enthalpy decreases with the increase in the atomic size.
(3) Penetration effect of the electrons: Ionisation enthalpy increases as the penetration effect increases. This effect means how much electrons of valence shell are attracted towards the nucleus. The order of penetration effect is s > p > d > f.
(4) Shielding or screening of the inner shell electrons: As the shielding or the screening effect of the inner electrons increases, the Ionisation enthalpy decreases. In multi-electron atoms, the repulsive from the electrons in the inner shells experienced by valence shell electrons is called shielding effect.
Zff = Z – σ
(5) Electronic Configuration: If an atom contains exactly half-filled or completely filled orbitals, then such an arrangement has extra stability. Therefore, the removal of an electron from such an atom requires more energy than expected. That is why, Be and N have higher Ionisation enthalpy than B and O respectively.
Variation of Ionisation enthalpy in periodic table:
In group, generally, Ionisation enthalpy decreases down the group.
In period, generally, the ionisation enthalpy increases from left to right.
Order of 1st Ionisation enthalpy :
Li < B < Be < C < H < O < N < F < Ne < Ar.
Electron gain enthalpy: In the classification of elements
The amount of energy released when a neutral isolated gaseous atom accepts an extra electron to form the gaseous negative ion is called electron gain enthalpy. It is denoted by ΔegH.
This process may be represented as:
X(g) + e —– X– (g); ΔH = ΔegH
Factors on which the electron gain enthalpy depends on:
(1) Atomic size: As the size of the atom increases, the force of attraction between the nucleus and the incoming electron decreases and hence the electron gain enthalpy becomes less negative.
(2) Nuclear charge: As the nuclear charge increases, the force of attraction between the nucleus and the incoming electron increases and hence the electron gain enthalpy becomes more negative.
(3) Electronic Configuration: Elements having exactly half-filled or completely filled orbitals are very stable. As a result, energy has to be supplied to add an electron. Hence, electron gain enthalpy have large positive value since they do not accept the additional electron so easily.
Variation of electron gain enthalpies in the periodic table:
Within a group: In general, the election gain enthalpy become less negative as we down a group.
The electron gain enthalpy of some of the elements of second period, I.e., O and F are, however less negative than the corresponding elements S and Cl of the third period.
Along a period: In general, electron gain enthalpy becomes more and more negative from left to right in a period.
Halogens have the most negative electron gain enthalpies. The electron gain enthalpy of noble gases is positive.
Electronegativity:
The tendency of the atom of an element to attract the shared pair of electrons towards itself in a covalent bond is called electronegativity. It is denoted by X.
Electronegativity depends on the following factors:
(1) State of hybridization: For example, a sp-hybridized carbon is more electronegative than sp2- hybridized carbon, which in turn is more electronegative than a sp3- hybridized carbon.
(2) Oxidation state of the element: As the oxidation state of the element increases, the electronegativity increases.
(3) Nature of the substituents linked with atoms: For example, the carbon atom in CF3I acquires greater positive charge than in CH3I. Therefore carbon atom in CF3I is more electronegative than in CH3I.
Trends of electronegativity in periodic table:
The electronegativity increases along a period from left to right and decreases down a group. Thus Fluorine is most electronegative element and is given a value of 4.0 (Pauling scale) and cesium is the least electronegative element.
Some important formula for Electronegativity 
Periodicity of valence or oxidation state: In the Classification of Elements.
Valence: The number of electrons present in the outermost shell of an atom is called the valence electrons and the number of electrons needs to acquire stable Configuration is called its valence.
Oxidation state: The charge assigned on an atom, when bonded with the atom of other element is called its oxidation state. The variation of valence electrons and valence takes place in following ways.
Anomalous behaviour:
In fact, this type of behaviour is generally found in the elements of 2nd group. Some elements of 2nd period show similarities with the elements of third period placed diagonally to each other. This type of behaviour is known as diagonal relationship. For example, Li with Mg, Be with Al and B with Si show similarities in properties.
Reasons for Diagonal relationship:
(1) Small size
(2) Large charge/radius ratio
(3) High electronegativity
(4)Absence of d-orbitals