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Classification Of Elements And Periodicity In Properties

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Classification of Elements and Periodicity in Properties PDF Notes, Important Questions and Synopsis

SYNOPSIS

  • Johann Dobereiner classified elements into a group of three elements called triads.
  • In Dobereiner’s triad, the atomic weight of the middle element is very close to the arithmetic mean of the other two elements. This relationship is known as law of triads.
  • Because Dobereiner’s law of triads worked only for few elements, it was dismissed.
  • John Newland arranged the elements in increasing order of atomic weight and noted that the properties of every eighth element are similar to the first one. This relationship is known as law of octaves.
  • Lothar Meyer proposed that when elements are arranged in increasing order of atomic weights, similarities appear at a regular interval in physical and chemical properties.
  • According to Mendeleev’s periodic law, the physical and chemical properties of elements are periodic functions of their atomic weights.
  • Merits of Mendeleev’s periodic table
    • Mendeleev’s periodic table helped in correcting the atomic masses of some of the elements such as gold, beryllium and platinum based on their positions in the periodic table.
    • Mendeleev could predict the properties of some undiscovered elements such as scandium, gallium and germanium and left gaps for these undiscovered elements in his periodic table.
  • Demerits of Mendeleev’s periodic table
  • Position of hydrogen is not correctly defined in the periodic table. It is placed in group 1 though it resembles both group 1 and 17.
  • In certain pairs of elements, increasing order of atomic masses was not obeyed.
  • Isotopes were not given separate places in the periodic table although Mendeleev’s classification is based on the atomic masses.
  • Some similar elements are separated, while dissimilar elements are grouped together.
  • Mendeleev did not explain the cause of periodicity among the elements.
  • Lanthanoids and actinoids were not given a separate position in the table.
  • Moseley performed experiments and studied frequencies of the X-rays emitted from the elements. With these experiments, he concluded that atomic number is essentially the fundamental property of an element than its atomic mass.
  • After Moseley’s experimental results, Mendeleev’s periodic law was modified to the modern periodic law.
  • According to the modern periodic law, the physical and chemical properties of the elements are periodic functions of their atomic numbers.
  • Modern periodic table is also referred to as the long form of the periodic table.
  • Horizontal rows in the periodic table are called periods, and vertical columns in the periodic table are called groups.
  • In the modern periodic table, there are 7 periods and 18 groups.
  • The periodic number corresponds to the highest principal quantum number of elements.
  • In the modern periodic table, 14 elements of both sixth and seventh periods, i.e. lanthanoids and actinoids, respectively, are placed separately at the bottom of the periodic table.
  • Elements with atomic number greater than 92 are called transuranic elements.
  • According to IUPAC, until a new element’s discovery is proved and its name is officially recognised, it is given a temporary name. This nomenclature is based on latin words for their numbers.
  • The interim names of the newly discovered elements are derived by combining together the roots in the order of digits which make up the atomic number and -ium is added at the end.
  • Notation for the IUPAC nomenclature of elements:

    Digit

    Name

    Abbreviation

    0

    Nil

    N

    1

    Un

    U

    2

    Bi

    B

    3

    Tri

    T

    4

    Quad

    Q

    5

    Pent

    P

    6

    Hex

    H

    7

    Sept

    S

    8

    Oct

    O

    9

    Enn

    E


    The distribution of electrons into the orbitals of an atom is called its electronic configuration.
  • The electrons in an orbital are filled according to the n + l rule.
  • The number of elements in each period is twice the number of atomic orbitals available in the energy level that is being filled.
  • On moving down a group in the periodic table, the number of shell increases from 1 to 7.
  • Value of the principal quantum number for the valence or outermost shell gives the period.
  • The first period has the principal quantum number n = 1, contains two elements and corresponds to the K-shell.
  • The second period has principal quantum number n = 2, contains eight elements and corresponds to the L-shell.
  • The third period has principal quantum number n = 3, contains eight elements and corresponds to the M-shell.
  • The fourth period has principal quantum number n = 4 and contains eighteen elements.
  • Elements from scandium (Z = 21) to zinc (Z = 30) are called 3d transition series of elements or first transition series.
  • The fifth period has principal quantum number n = 5 and contains eighteen elements.
  • Elements from yttrium (Z = 39) to cadmium (Z = 48) are called 4d transition series of elements or second transition series.
  • The sixth period has principal quantum number n = 6 and contains 32 elements.
  • Elements from lanthanum (Z = 57), hafnium (Z = 72) to mercury (Z = 80) are called 5d transition series of elements or third transition series.
  • Fourteen elements from cerium (Z = 58) to lutetium (Z = 71) are called elements of inner transition series or lanthanoid series.
  • Fourteen elements from thorium (Z = 90) to lawrencium (Z = 103) are called elements of 5f inner transition series or actinoid series.
  • The modern periodic table is divided into four main blocks—s-block, p-block, d-block and f-block depending on the type of orbital that are being filled with an exception of hydrogen and helium.
  • The elements in which the last electron enters the s-orbital of their outermost energy level are called s-block elements.
  • The s-block consists of two groups such as group 1 and group 2.
  • The elements of group 1 are called alkali metals and have ns1 as the general outer electronic configuration.
  • The elements of group 2 are called alkaline earth metals and have ns2 as the general outer electronic configuration.
  • The elements in which the last electron enters the p-orbital of their outermost energy level are called p-block elements.
  • The p-block elements constitute elements belonging to group 13 to 18.
  • Elements of s-block and p-block are collectively called representative elements.
  • The outermost electronic configuration of p-block elements varies from ns2np1 to ns2np6.
  • Elements of group 18 having ns2np6 configuration are called noble gases.
  • Elements of group 17 are called halogens.
  • Elements of group 16 are called chalcogens.
  • Number of valence electrons in group = Group number − 10 for elements belonging to group 13 to 18.
  • Elements in which the last electron enters d-orbitals of penultimate energy level constitute d-block elements.
  • Elements of group 3 to 12 in the centre of the periodic table constitute the d-block elements.
  • General outer electronic configuration of d-block elements is (n − 1)d1–10 ns1–2.
  • The d-block elements constitute transition series elements. 
  • Elements in which the last electron enters f-orbitals are called f-block elements.
  • Elements of lanthanoid series have general outer electronic configurations of 4f1–14 5d0–1 6s2.
  • Elements of actinoid series have general outer electronic configuration of 5f1–14 6d0–1 7s2.
  • Elements in lanthanoid and actinoid series are called the inner transition series.
  • Metals comprise more than 78% of all known elements and appear on the left-hand side of the periodic table.
  • Non-metals are placed on the right-hand side of the periodic table.
  • The cause of periodicity of properties of elements is because of the repetition of similar electronic configuration of their atoms in the outermost energy shell after certain regular intervals.
  • Covalent radius for a homonuclear molecule is defined as half of the distance between the centres of the nuclei of two similar atoms bonded by a single covalent bond.
  • For a heteronuclear molecule, covalent radius may be defined as the distance between the centre of the nucleus of atom and mean position of the shared pair of electrons between the bonded atoms.
  • Metallic radius is defined as half of the internuclear distance between two neighbouring atoms of a metal in a metallic lattice.
  • For simplicity, the term atomic radius is used for both covalent and metallic radius depending on whether the element is a non-metal or metal.
  • Atomic radius decreases with increase in atomic number on going from left to right in a period.
  • Atomic radius of elements increase from top to bottom in a group.
  • Van der Waals radius is half of the distance between two similar atoms in separate molecules in a solid.
  • Ionic radius may be defined as the effective distance from the nucleus of the ion upto which it has an influence in the ionic bond.
  • On moving from top to bottom in a group in the periodic table, ionic radius increases.
  • On moving from left to right in a period in the periodic table, ionic radius decreases.
  • A quantitative measure of the tendency of an element to lose electrons is given by ionization enthalpy. It represents the energy required to remove an electron from an isolated gaseous atom in the ground state.
  • On moving from left to right along a period in the periodic table, ionization enthalpy increases. On moving along a period, successive electrons are added to orbitals in the same quantum level and shielding of the nuclear charge by the inner core of electrons does not increase to an extent to compensate for increased attraction of electrons to the nucleus. Thus, increasing nuclear charge outweighs shielding across a period. Eventually, more energy is required to remove the outermost electron.
  • On moving from top to bottom in a group in the periodic table, ionization enthalpy decreases. On moving down a group, successive shells are added and the outermost electron moves further away • The d-block elements constitute transition series elements.
  • Elements in which the last electron enters f-orbitals are called f-block elements.
  • Elements of lanthanoid series have general outer electronic configurations of 4f1–14 5d0–1 6s2.
  • Elements of actinoid series have general outer electronic configuration of 5f1–14 6d0–1 7s2.
  • Elements in lanthanoid and actinoid series are called the inner transition series.
  • Metals comprise more than 78% of all known elements and appear on the left-hand side of the periodic table.
  • Non-metals are placed on the right-hand side of the periodic table.
  • The cause of periodicity of properties of elements is because of the repetition of similar electronic configuration of their atoms in the outermost energy shell after certain regular intervals.
  • Covalent radius for a homonuclear molecule is defined as half of the distance between the centres of the nuclei of two similar atoms bonded by a single covalent bond.
  • For a heteronuclear molecule, covalent radius may be defined as the distance between the centre of the nucleus of atom and mean position of the shared pair of electrons between the bonded atoms.
  • Metallic radius is defined as half of the internuclear distance between two neighbouring atoms of a metal in a metallic lattice.
  • For simplicity, the term atomic radius is used for both covalent and metallic radius depending on whether the element is a non-metal or metal.
  • Atomic radius decreases with increase in atomic number on going from left to right in a period.
  • Atomic radius of elements increase from top to bottom in a group.
  • Van der Waals radius is half of the distance between two similar atoms in separate molecules in a solid.
  • Ionic radius may be defined as the effective distance from the nucleus of the ion upto which it has an influence in the ionic bond.
  • On moving from top to bottom in a group in the periodic table, ionic radius increases.
  • On moving from left to right in a period in the periodic table, ionic radius decreases.
  • A quantitative measure of the tendency of an element to lose electrons is given by ionization enthalpy. It represents the energy required to remove an electron from an isolated gaseous atom in the ground state.
  • On moving from left to right along a period in the periodic table, ionization enthalpy increases. On moving along a period, successive electrons are added to orbitals in the same quantum level and shielding of the nuclear charge by the inner core of electrons does not increase to an extent to compensate for increased attraction of electrons to the nucleus. Thus, increasing nuclear charge outweighs shielding across a period. Eventually, more energy is required to remove the outermost electron.
  • On moving from top to bottom in a group in the periodic table, ionization enthalpy decreases. On moving down a group, successive shells are added and the outermost electron moves further away • The d-block elements constitute transition series elements.
  • Elements in which the last electron enters f-orbitals are called f-block elements.
  • Elements of lanthanoid series have general outer electronic configurations of 4f1–14 5d0–1 6s2.
  • Elements of actinoid series have general outer electronic configuration of 5f1–14 6d0–1 7s2.
  • Elements in lanthanoid and actinoid series are called the inner transition series.
  • Metals comprise more than 78% of all known elements and appear on the left-hand side of the periodic table.
  • Non-metals are placed on the right-hand side of the periodic table.
  • The cause of periodicity of properties of elements is because of the repetition of similar electronic configuration of their atoms in the outermost energy shell after certain regular intervals.
  • Covalent radius for a homonuclear molecule is defined as half of the distance between the centres of the nuclei of two similar atoms bonded by a single covalent bond.
  • For a heteronuclear molecule, covalent radius may be defined as the distance between the centre of the nucleus of atom and mean position of the shared pair of electrons between the bonded atoms.
  • Metallic radius is defined as half of the internuclear distance between two neighbouring atoms of a metal in a metallic lattice.
  • For simplicity, the term atomic radius is used for both covalent and metallic radius depending on whether the element is a non-metal or metal.
  • Atomic radius decreases with increase in atomic number on going from left to right in a period.
  • Atomic radius of elements increase from top to bottom in a group.
  • Van der Waals radius is half of the distance between two similar atoms in separate molecules in a solid.
  • Ionic radius may be defined as the effective distance from the nucleus of the ion upto which it has an influence in the ionic bond.
  • On moving from top to bottom in a group in the periodic table, ionic radius increases.
  • On moving from left to right in a period in the periodic table, ionic radius decreases.
  • A quantitative measure of the tendency of an element to lose electrons is given by ionization enthalpy. It represents the energy required to remove an electron from an isolated gaseous atom in the ground state.
  • On moving from left to right along a period in the periodic table, ionization enthalpy increases. On moving along a period, successive electrons are added to orbitals in the same quantum level and shielding of the nuclear charge by the inner core of electrons does not increase to an extent to compensate for increased attraction of electrons to the nucleus. Thus, increasing nuclear charge outweighs shielding across a period. Eventually, more energy is required to remove the outermost electron.
  • On moving from top to bottom in a group in the periodic table, ionization enthalpy decreases. On moving down a group, successive shells are added and the outermost electron moves further away from the nucleus. Because electrons are present in the inner shells, shielding of nuclear charge increases. Thus along a group, shielding outweighs increasing nuclear charge. Eventually, less energy is required to remove the outermost electron.
  • Among various groups in the periodic table, group 18 elements have the highest ionization enthalpy because of stable electronic configuration.
  • When an electron is added to a neutral gaseous atom to convert it into a negative ion, enthalpy change accompanying the process is called electron gain enthalpy.
  • In general, electron gain enthalpy becomes more negative from left to right in a period. The effective nuclear charge increases from left to right across a period. Thus it is easier to add an electron to a smaller atom because the added electron would be on average closer to the positively charged nucleus.
  • In general, electron gain enthalpy becomes less negative as we go from top to bottom in a group. This is because the added electron would be farther away from the nucleus.
  • The ability of an atom in a chemical compound to attract shared electrons to itself is called electronegativity.
  • The electrons present in the outermost shell are called valence electrons, and these electrons determine the valence of atom.
  • Valence of representative element is usually equal to
    • The number of electrons in the valence shell
    • 8(the number of electrons in the valence shell)
  • On moving down a group because the number of valence electrons remains the same, all elements exhibit the same valence.
  • The oxidation state of an element in a particular compound gives the charge acquired by its atoms on the basis of electronegativity consideration from other atoms in the molecule.
  • It is observed that some elements of the second period show similarities with elements of the third period diagonally present though they belong to different groups. This similarity in the properties of elements diagonally present is called diagonal relationship.
  • Lithium is diagonally related to magnesium, beryllium is diagonally related to aluminium and boron is diagonally related to silicon.
  • The anomalous behavior of the first element of s- and p-block elements of each group as compared to other group members is because of the following reasons:
    • Small size of atom
    • Large charge/radius ratio
    • High electronegativity
    • Non-availability of d-orbitals in their valence shell
  • In the periodic table, there is high chemical reactivity at two extreme ends and chemical reactivity is lowest at the centre. Maximum chemical reactivity at the extreme left (alkali metals) is shown by easy loss of electrons forming cations, and maximum chemical reactivity at the extreme right (among halogens) is shown by gain of electrons forming anions.
  • Tendency of an element to lose or gain electrons is also related to metallic or non-metallic characters.
  • Elements at the extreme left of the periodic table have a tendency to lose electrons and become positively charged. Hence they show metallic character.
  • Elements at the extreme right of the periodic table have a tendency to gain electrons. Hence they show non-metallic character.
  • Metallic character decreases along a period on moving from left to right in the periodic table. Because elements at the extreme left of the periodic table show metallic character, oxides formed by them are basic.
  • Non-metallic character increases along a period on moving from left to right in the periodic table. Because elements at the extreme right of the periodic table show non-metallic character, oxides formed by them are acidic.
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