Genesis of Periodic Classification and Mendeleev Periodic Table:

Genesis of Periodic Classification and Mendeleev Periodic Table: Classification of elements and periodicity in properties: Periodic table helps in the systematic study of most of the elements found in nature. The elements are classified into different groups and periods in the periodic table which helps to study about the compounds formed by those elements. Moreover the analysis of the ionization energy, electronegativity, electron affinity etc is also possible. Dobereiner triad rule: Dobereiner pointed out that the atomic weight of an element present in the middle in a group of three elements is equal to the mean of the first and third element only if they possess the same physical and chemical property. Representation of Dobereiner triads: Triad elements Li Na K Ca Sr Ba Atomic weight 7 23 39 40 88 137 Mean value 23 88.5 Drawback: His rule was unable to arrange all the known elements in a triad. Law of Octaves: Newlands founded this law in 1865 in which the elements were arranged in order of increasing atomic weights. Every eighth element was found to possess similar properties to that of the first element. This law was given this name because of its resemblance with the octaves of music where the eighth note resembles the first. Drawback: His law seemed to be correct only for elements up to Calcium. Mendeleev s periodic table: Mendeleev published the periodic law for the first time which states as the properties of the elements are a periodic function of their atomic weights. Characteristics: i)it is based on the atomic weight. ii)63 elements were known and noble gas were not discovered. iiithe elements were arranged in a horizontal row called periods and vertical column as groups in the periodic table in order of increasing atomic weights. iv) There were 7 periods and 8 groups in Mendeleev s periodic table. v)each group upto 7 th is divided into two subgroups A and B. A subgroups are normal elements and B subgroups are termed as transition elements. vi)the 8 th group consist of 9 elements in three rows. 2

vii)the properties of elements belonging to same grop possess same properties. Merits: i)study of the properties of the elements became easier when the elements were classified in groups according to their properties. ii)his law encouraged for the prediction of new elements. iii)correction were done in atomic weights of some elements. Atomic weight= valency x equivalent weight. Demerits: i)uncertainty in the position of hydrogen.lear. ii)isotopes were not given separate positions. iii)lanthanides and actinides were related to IIA or IIB was not clear. iv) the order of increasing atomic weights were not followed strictly. Modern Periodic Law and present form of Periodic Table : Modern periodic table: I) Moseley proposed the Modern periodic table. II) Atomic number is the basis of this law. III) A high speed electron was bombarded on different metal surface to obtain X-rays.It was found out that v is proportional to Z where v is the frequency of X-ray and Z = atomic number. Modern periodic law: This law is based on the assumption that the physical and chemical properties of elements are the periodic function of their atomic number. Characteristics: i) The 9 vertical columns present were called groups ii) 1 st to VIIIth group+ group of inert gases. iii) Ramsay introduced the inert gases. iv) The 7 horizontal series termed as periods. Long form of periodic table: i) Based on concept of Bohr Bury electronic configuration and atomic number. ii) Rang and Werner proposed this model. iii) 7 periods and 18 columns. iv) The 18 vertical columns are named as 1 st to 18 th group according to IUPAC v) The outermost shell of the elements present in the same group possess equal number of electrons. Li, Be, B,the second period elements shows diagonal relationship with Mg,Al,Si 3

of third period elements. Na,Mg,Al,Si,P,S,Cl of the 3 rd period represents properties of other elements in the respective group. Atomic number of the last inert gas is 86. Number of gaseous elements present =11 Number of liquid elements=6 Number of of solid elements =95 The 2 nd period possess maximum number of gaseous elements (N,O,F,Ne respectively) Electronic Configuration of Elements s block elements: 1.The elements in the periodic table whose last electron enters the s orbital are termed as the s block elements. A maximum of two electrons can be accommodated in the s orbital. ns1 and ns2 are the general formula of the elements of this block. n varies from 1 to 7. The first group elements are known as alkali metals as the reaction with water leads to the formation of alkali.the second group elements are known as the alkaline earth metal because of the reaction of their oxides with water to form alkali. 14 elements are present in the s block. H and He are gaseous elements whereas Fr87 and Ra88 are radioactive. liquid elements belonging to s block are Cs and Fr. p block elements: The elements in the periodic table whose last electron enters the p orbital are termed as the p block elements A maximum of six electrons can be incorporated in the p-orbital. General formula for the p;block elements is ns2p1-6. Where n=2 to 6. The elements having the formula ns2p6 are inert as their energy level is filled. Excluding he, a total of 30 elements are present in the p block Ga and Br are liquids The step like structures in the periodic table seperates the elements into metals, non metals and metalloids. d block elements: The elements in the periodic table whose last electron enters the d orbital are termed as the d block elements. The electrons gets filled up in the penultimate shell in the d orbital Between s and p block elements, the d block elements lie. The general formula of these elements is ns2p6d1-10 where n=4 to 7. The elements present in this block are metals. Mercury is the only liquid element present in d block. f block elements: The elements in the periodic table whose last electron enters the f orbital are termed as the f block elements 4

Atomic number 58-71 and from 90-103 represents the f block elements.elements from 58-71 are called lanthanides and that from 90-103 are called actinides. The lanthanides are present in very low abundance and are so called the rare earth elements. 28 f block elements are present in the periodic table. The actinide elements are radioactive. The general formula of these elements is (n-2)s 2 p 6 d 10 f 1-14 (n-1)s2p6d10ns2 where n =6, 7. Metals, non metals and metalloids: Apart from the classification of elements into s, p, d, f block, they are further classified into metals non metals and metalloids. 78% of the known elements are metals and are present in the side of the periodic table. Properties of metals: Metals are solids at room temperature except mercury. They have high melting and boiling point They are good conductors of heat and electricity. They are malleable and ductile Properties of non metals: Non metals are usually solids or gases at room temperature. They possess low boiling and melting point except carbon and boron. They are poor conductors of heat and electricity. They are brittle. Metalloids possess characteristics of both the metals and non-metals. They are also called semi metals. Periodic Trends in Properties of Elements: Trends in physical properties: The variation in the numerous physical properties such as atomic or ionic size, melting and boiling point, ionization energy, electronegativity, electron affinity etc shows periodic variations. Effective nuclear charge (Z eff ): A decrease in the force of attraction on valence electron is due to inner shell is called the screening effect or the shielding effect. Due to screening effect, the valence shell electron experience less force of attraction exerted by the nucleus. This is called the effective nuclear charge or Z eff. Z eff =Z- where Z= atomic number, is the screening constant. For 1s electron, =0.30, 0.35 for ns and np electron, 0.85 for (n-1) penultimate orbit s,p,d electrons, 1 for (n-2) and inner all the electrons present in s,p, d, f. Atomic radius: 5

The size of an atom increases from top to bottom in a group and decreases from left to right in a period in the periodic table.the atomic radii depends on the size of an atom. With increasing size, the atomic radius too increases.some other ways of calculating the size of an atom are: Covalent radius: The internuclear distance between two bonded atoms is termed as the covalent radius. The covalent radius of an atom A in a molecule A 2 is: r A d A A 2 Metallic radius: it is the one half of the internuclear distance between two closest metals in a metallic crystal. Metallic radius> Covalent radius Vander Waal s radius: it is the one half of the distance between the nuclei of two adjacent atoms that belongs to two neighbouring molecules of a compound in the solid state. Vander Waals radius > metallic radius> covalent radius Ionic radius: By loosing an electron or gaining an electron, a neutral atom may convert to a cation or an anion.the intermolecular distance between two ions helps in the determination of the ionic radii of the ion. Radius of cation: Radius of the cation is lower than the corresponding atom.this is because the atom looses an electron to form a cation and thus on decreasing te number of electrons on the atom, the size too decreases. Radius of an anion: Radius of the anion is higher than the corresponding atom.this is because the atom gains an electron in the formation of an anion. On increasing the number of electrons on the atom, the size of the atom increases due to repulsion between the electrons. Factors affecting atomic radius are: Atomic radius decreases with the increase in +ve charge. Atomic radius increases with the increase in -ve charge. Atomic radius increases with the number of shells Atomic radius decreases with increase in bond energy. Periodic variation : Ionic radii increases from top to bottom in a group whereas it decreases from left to right across a period. Exception is observed in the lanthanides due to lanthanide contraction. The nuclear charge in lanthanides and actinides increases by +1 due to which atomic size decreases. This is termed as lanthanide contraction. 6

Another exception is observed in the atomic size of Ga, Al and B is in their The trend of increasing atomic size observed is B<Al Ga instead of B<Al<Ga. This is due to increase in nuclear charge by 0.15 because of which the size remains almost constant. Isoelectronic species: Isoelectronic species are those that contains same number of electrons. In such species such as N 3-, O 2-,Ne, Na +,Mg 2+ etc, the atomic size decreases with the increase in positive charge. Ionization energy or Ionization potential: It is the minimum amount of energy that is required when an electron is removed from the outermost shell of a neutral gaseous atom. For an atom M, M+E 1 -- M + + e - E 1 =1 st ionization potential M + +E 2 -- M 2+ + e - E 2 =2 nd ionization potential M 2+ +E 3 -- M 3+ + e - E 3 =3 RD ionization potential Trend in successive ionization energies are as follows: 1 st ionization potential <2 nd ionization potential <3 rd ioni ionization potential Ionization energy is affected by: Atomic size: increase in atomic size decreases the ionization potential. This is because the outermost shell remains far apart from the nucleus. Effective nuclear charge: ionization potential increases with increase in nuclear charge. Penetration power of subshell: Penetration power of subshell is s>p>d>f. as the subshell is closer to the nucleus, the ionization potential required will be more. Half-filled and fully filled orbitals: Half-filled and fully filled orbitals are very stable because of which the ionization potential increases. Periodic trend: The Ionization energy increases down a group and decreases along a period due to decrease in atomic size along the period Electron affinity: The amount of energy that is released when an electron is added to the outermost shell of a neutral atom is the electron affinity. Factors affecting electron affinity: Atomic size decreases with increasing electron affinity Electron affinity increases with decrease in screening effect. Electron affinity increases with effective nuclear charge. Periodic trend: The Electron affinity increases down a group and decreases along a period due to decrease in atomic size along the period. The electron affinity of Cl>F Electronegativity: The tendency of an atom to attract the shared pair of electron towards itself is the electronegativity of that Atom. Factors affecting electronegativity: 7

Electronegativity decreases with increasing atomic size Electronegativity increases with effective nuclear charge. Electronegativity increases with increase in %s character. Electronegativity increases with increase in oxidation state. Electronegativity: Periodic trend: The Electronegativity decreases down a group and increases along a period due to increase in the non-metallic character along the period. Non-metallic elements have strong tendency to gain electrons. The electronegativity of F>Cl. Electronegativity can be measured by different scales. They are as: 1. Pauling scale of electronegativity: X 0. 208 where AB =23.06(X A X B ) 2 and X A, X B are the electronegativity of A and A X B A B B. 2. Alred Rochow scale of electronegativity: 2 Z eff 0.359 Z eff Electronegativity = = 0. 744 2 2 r r Z eff = Z- where Z = atomic number, = shielding constant. 3. Mullikan s scale of electronegativity: Electronegativity= (electron affinity +ionization potential)/2 Chemical properties: Chemical reactivity: The ionization energy in the extreme left of a periodic table is the least and electronegativity in the extreme right is the highest. So, it results in high chemical reactivity at the two extremes. Electronegativity: Acidic strength of hydrides increases with increasing electronegativity of metal Hydroxides are basic if electronegativity of the metal is less than 1.7 Example : NaOH ClOH 0.9 (basic) 3(acidic) Nature of oxides increases along a period as well as down the group. Na 2 O MgO are basic oxides Al 2 O 3 SiO 2 are amphoteric P 4 O 10 SO 3 Cl 2 O 7 are acidic. Ionization potential : Ionization potential decreases with increase in metallic character. Metallic: low ionization potential Non-metallic : high ionization potential 8

Ionization potential decreases with increase in reducing character. Stability of oxidation state: 1. If the difference between two successive ionization potential 16eV elements with lower oxidation states are stable. eg: Na----- > Na + 1 st Ionization potential Na + ------> Na 2+ 2 nd Ionization potential Difference of ionization potential =42.7 ev which is greater than 16eV and so Na + is stable. 2. If the difference between two successive ionization potential 11eV elements with higher oxidation states are stable. Mg ---- > Mg + Mg+ ----- > Mg 2+ 1 st Ionization potential 2 nd Ionization potential Difference of ionization potential =7.4.eV which is smaller than 11eV, so Mg 2+ is stable. Boiling and melting points: Trend along a period: In a period boiling and melting point increases first and then decreases For alkali metals, there exists low boiling and melting point. For metals, tungsten possess the highest melting point (3410 0 C) and mercury has the lowest (-38 0 C) For non- metals, carbon in the form of diamond possess highest meting point of 3727 0 C and the lowest for helium of -270 0 C. Trend in a group: s block elements boiling and melting point decreases down a group due to weak cohesive energy of metallic bond. In p block elements, from IIIA to IVA boiling and melting point decreases and increases from VA to 0 (noble gas) Monoatomic molecules have low boiling and melting point than diatomic molecules. Atomic solid like C, Si, B have higher boiling point and melting point due to strong covalent bonding Molecular solids have low boiling and melting point due to weak Vander Waals force among molecules. In d block elements, melting and boiling point increases down the group due to lanthanide contraction which increases the bond energy. All the best! Team Gradeup Attempt the free Mock Test here Download Gradeup, the best IIT JEE Preparation App 9

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