periodic table

• William proust suggested that atomic masses of all elements are simple multiple of atomic mass of hydrogen
• doberiner observed that several groups of closely related elements in which atomic mass of central elements is mean of atomic mass of two other element.
• newland arranged elements according to increasing atmoci mass and noticed every eighth element had similar properties.
• Mendeleev's arrangement was based on atomic weight (atomic mass) rather than atomic number
• mendeleev left gaps in his table where he predicted the existence of undiscovered elements, these were later discovered by others
• modern periodic table law "the physical and chemical property of elements are periodic function of their atomic numbers"
• periods

  • 1 is shortest n=1
  • 2 and 3 are short n=2 or 3
  • 4 and 5 long period
  • 6 longest
• groups

  • 18 groups
  • from left to right
  • IA, IIA, IIIB, IVB, VB, VIB, VIIB, VIIIB, IB, IIB, IIIA, IVA, VA, VIA, VIIA, VIIIA
  • 1 2 3 4 5 6 7 (10,11,12) 8 9 13 14 15 16 17 18
  • VIIIB consist of subgroups belong to iron triad
  • IA= alkali metal
  • IIA=alkaline earth metal
  • VIA=chalcogen metal (many elements can be extracted by oxides
  • VIIA=halogene
  • VIIA= inert


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• Elements with atomic no 58 to 71 are grouped with lanthanum and are called lanthanides or rare earth Element.

  Elements with atomic no 90 to 103 are grouped with actinium and are called actinides.

  placed outside periodic table from being unconveniently long.
• S block elements are in group IA and IIA.

  • soft silver metal
  • highly electropositive (tendency to lose electron)
  • Impart colour
• P block elements are of group IIIA to VIIIA

  • mostly solid, few gases (N2, O2, F2, Cl2)
  • metal or nonmetal
• D block elements are transition element

  (n-1)d 1 to 10 ns2
• The repulsion provided by the inner shell electron to the outermost electron against nuclear attraction is called shielding.
• Effective nuclear charge= Nuclear charge - shielding constant
• Increase of no. of electrons decreases effective nuclear charge(Z*)
• Atomic radii

  • Distance from nucleus to the point where electron cloud density is effectively zero.
  • 1A=10^-8
  • 1nm=10^-7
  • As n increases atomic radius increases
• Atomic radii

  • increase of effective nuclear charge decreases size of atom
  • Atomic radii decreases across s and p block element
• Halogen has smallest size instead of inert gas because inert gas has saturated valence shell and it doesnt form covalent bond but forms vanderwaal's bond in which electron cloud is more diffusable.
• In a group atomic radii increases with in increase in atomic no
• Isoelectronic ions are those ions which have similar electronic configuration and no of electron but different nuclear charge. for eg; S2- and Cl- have nuclear charge +16 and +17 but same no of electron 18
• ionic radii of isoelectronic ions decrease with increase in magnitude of nuclear charge.
• The size of atom is greater than cation.

  The size of atom is smaller than anion.
• The minimum amount of energy needed to remove loosely bound electron from an isolated gaseous atom is called ionization energy
• First I.E<Second I.E because effective nuclear charge increases
• • Increase of atomic radius and n, increases ionization energy
  • Increase of Effective nuclear charge, increases ionization energy
  • Half filled or compelety filled orbitals are stable with high ionization energy
• • I.E increase from left to right bc increase in z*
  • I.E of beryllium>boron bc it is easier to remove electron from p orbital of B than s orbial of Be because penetrating power to nucleus is more for s than p
  • I.E of N>O bc in O two electrons are paired and in pair repulsion increase energy and it is easier to remove electron.
• • I.E decreases from top to bottom in group due to increase of n.
  • S-block element have lowest I.E and p block element have highest I.E.
  • Increase of +ve charge increases I.E
• Electron affinity is the measure of capacity of an element to form negative ion.

  It is the amount of energy released or absorbed whn an electron is added to an isolated gaseous atom.
• • As size increases, electron affinity decreases.
  • As nuclear charge increases, electron affinity increases.
  • Elements having half filled or completely filled orbitals are very stable so they have low or zero electron affninity
• • On top to bottom in a group E.A decreases due to increase in size.
  • On left to right in a period E.A increases due to decrease in size.
• The tendency of an atom to attract shared paired of electrons to itself is called electronegativity
• • Smaller the size greater is the electronegativity.
  • Increase of effective nuclear charge increase EN.
  • A cation attracts pair more readily due to smaller size.
  • A anion doesnt readily attract due to larger size.
• Metallic character decreases across the period from left to right due to increase of effective nuclear charge.

  Metallic character increases from top to bottom due to increase in n and decrease in i.e
• EN of elements decreases top to bottom in a group due to increase in atomic radius

  EN of elements increases from left to right due to decrease in atomic radius.
• What are the elements present in group VIIA called and why?

  halo-salt gen-origin, salt of alkali