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Periodic Table 5
 

Periodic Table 5

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    Periodic Table 5 Periodic Table 5 Presentation Transcript

    • PERIODIC TABLE
    • Periodicity
      • Recurrence of similar properties of elements at a regular intervals is called periodicity.
      • Some of the properties that show periodic variation are:
      • i. atomic radius & ionic radii
      • ii. Melting & boiling points
      • iii. Electrical conductivity
      • iv. Enthalpy of vapourisation
      • v. Electronegativity
      • vi. Ionisation energy
    • Atomic radii
      • Atom is a sphere. Associated with an atom, there are 3 radii to be considered:
      • atomic radius- distance from the centre of its nucleus to the edge of the atom. Bigger the atomic radius the bigger the atom
      • Covalent radius- half distance between the nuclei of two atoms bonded by a covalent bond in the same molecule
      • Van der waals radius- half the distance between the nuclei of the two atoms of adjacent molecules.
      • Atomic radius is influenced by two factors:-
      • Nuclear charge per atom
      • Electron shielding effect
    • Nuclear charge per atom
      • An increase in nuclear charge per atom increase the binding force on the valence electron by the nucleus. As such, the atomic radius is decreased.
      • Effective nuclear charge=
      • no. proton – inner shell electron
      • Example: Magnesium; 1s 2 ,2s 2 , 2p 6 , 3s 2
      • * 12-10= +2
    • Electron shielding effect
      • The electron shielding effect on valence electron of the X atom of electronic configuration 2.8.1 is stronger than Y atom of electronic configuration 2.1.
      • This is due to the fact that the X atom has 2 inner shells with 10 electrons which are able to provide a stronger shield to its valence electrons compared to the Y atom which has only one inner shell with 2 electrons.
      • Note: the effect of the electron shielding is opposite to the nuclear charge
    • Variation of atomic radius across period 2 & 3 from left to right
      • Nuclear per charge increases because the proton number increases.
      • The electron shielding effect of the inner shells remain constant because the number of inner shells are the same. The electrons are added to the outer shells of atoms and not to the inner shells.
      • Nuclear attraction on the valence electron of atom increases causing the electron cloud to be pulled closer to the nucleus.
      • Thus the atomic radius decreases.
    • Variation of atomic radius for transition elements
      • For transition elements, from left to right, the atomic radii are almost constant with no significant changes.
      • Nuclear per charge increases because the proton number increases.
      • The electron shielding effect increases because the electrons are added to the inner d-orbital.
      • The increase in electron shielding effects offsets the increase in the nuclear charge.
      • Overall the atomic radius does not show a significant change.
    • Variation of atomic radius down the groups
      • As we go down the group the electrons per atom increases, hence the number of electron shells increases.
      • This increases the number of inner electron shells, thus increase the electron shielding effect.
      • This offsets the increase in nuclear charge per atom due to increase in the number of protons.
      • Consequently the nuclear attraction on the valence electron weakens.
      • The valence electron tend to move outwards and increase the atomic radii.
    • Ionic radii
      • When an atom loses one or more electrons to form cations the positive charge exceed the negative charge. Thus the electrons become closer.
      • Cation are smaller in size compare to the atom which they are formed.
      • Size of atom Na= 186pm & size of Na + =99pm
      • When an atom gain one or more electron to form anions the nuclear charge remain constant but negative charge exceed the positive charge. In this case repulsion occurs and the size of the anions increases.
      • Size of Cl 2 =99pmm but size of Cl - - Cl - =181pm
    • isoelectronic
      • Ion that have equal number of electron in identical configuration.
      • Example: Na + = 1s 2 ,2s 2 ,2p 6
      • Mg 2+ = 1s 2 ,2s 2 ,2p 6
      • Mg 2+ is smaller than Na + because its nuclear charge is larger
      • For isoelectronic cations, the more positive the ionic charge the more smaller the ionic radii.
      • For isoelectronic anions, the more negative the charge, the larger the ionic radii.
    • Ionisation energy
      • Quantity of energy absorb to remove an electron.
      • The variation of ionisation energy across the period is analysed between:
      • A. The periods 2 and 3
      • B. For transition elements (block d- elements)
      • C. down the groups
    • Ionisation energy across the periods 2 and 3 from left to right
      • Ionisation energy increases because:-
      • The nuclear charge per atom increases
      • The atomic radius decreases
      • the distance between the nucleus and the valence electron shortens. Hence the valence electrons are closer to the nucleus.
      • The nuclear attraction on the valence electron increases.
      • The difficulty of removing the valence electron increases. Thus, Ionisation energy increases.
      • EXCEPTION 1:
      • Across the period 2, the ionisation energy of beryllium which is on left is higher than boron which is on the right.
      • Be: 1s 2 ,2s 2 B: 1s 2 ,2s 2 ,2p 1
      • Therefore the ionisation energy of B involves the removal of one valence electron from the 2p 1 while Be involves the removal of 2 electrons. Removing an electron from 2p 1 is easier compared to the 2s 2 configuration.
      • This is because the energy level of 2p is higher than 2s. The 2s 2 is more stable because the s orbital is completely filled and its more closer to the nucleus than 2p orbital. Hence the valence electron of the 2s orbital are more strongly bounded to the nucleus.
      • EXCEPTION 2:
      • Across the period 2, the ionisation energy of nitrogen which is on left is higher than oxygen which is on right.
      • This is because electronic configuration of N is 1s 2 ,2s 2 ,2p 3 while that of oxygen is 1s 2 ,2s 2 ,2p 4 . The ionisation energy involves the removal of valence electron from the 2p 4 while nitrogen from 2p 3 . Removing the electron from 2p 4 is easier than 2p 3 because the p orbital of 2p 3 are half filled orbitals which are more stable compare to 2p 4 which have incomplete p orbitals.
    • Ionisation energy down the group
      • As we go down the group, the ionisation energy decreases.
      • This is because:
      • The number of electron shells increases
      • Hence the distance of the valence electrons from the nucleus increase sequentially
      • The increase in the number of inner electron shells, cause electron shielding effect on the valence electron increases.
      • This factors decreases the force of attraction of the nucleus on the valence electron
      • This offsets the increasing nuclear charge per atom
      • Hence the valence electron is donates easily.
      • Consequently the ionisation energy decreases.
    • Ionisation energy for block d elements or transition elements
      • Involves the 1 st and 2 nd ionisation energy (removal of 1 st and 2 nd valence electron from the orbitals). Copper and chromium have higher ionisation energy from other block d elements.
      • Example: copper = [Ar] 3d 10 4s 1
      • chromium=[Ar] 3d 5 4s 1
      • The 1 st ionisation energy of both these elements involve the removal of valence electrons from the 4s orbital. The 2 nd ionisation energy of these elements involve the removal of valence electron from 3d orbitals while for other block d elements , it is from 4s orbital. Furthermore the 3d orbital for copper is completely filled and chromium is half-filled creating a more stable configuration. Hence the 2 nd ionisation energy for these two elements is more higher than the 1 st ionisation energy.
    • Electronaffinity, E A
      • E A is the measurement of the energy change when an atom gains an electron
      • F(g) + e - -> F - (g) ; E A = -328kJ/mol
      • This is an exothermic process where it releases energy and given in negative quantity.
    • Electronegativity
      • Electronegativity is a measure of attraction of atom for electrons in a covalent bond.
      • When the force of attraction of the nucleus increases on valence electron the electronegativity increases.
      • Across the period from left to right the electronegativity increases. This is because the decrease in atomic radius increase the force of attraction of the nucleus on valence electron.
      • However, as we go down any group the electronegativity decreases due to increasing in atomic radius down the group, decrease the force of attraction of the nucleus on valence electron.
    • Metallic properties
      • Across the period the metallic properties become non-metallic properties.
      • This is because the atomic radius decreases, the force of attraction of nucleus on the valence electron increases,
      • Difficulty to loose electrons (non-metallic)
      • Li, Be, B, C, N, O, F, Ne.
      • Li and Be show prominent metallic property with electrical conductivity.
    • Boiling & Melting Point
      • Depends on the strength of the intermolecular forces or bonds that must be weeken or destroyed before melting and boiling occurs.
      • Across the period 3 from Li to Ne, the melting point & boiling point increases until it peaks at C and then decreases sharply to N.
      • From N to Ne, the melting points & boiling points are about the same.
      • Metallic bond strength increases from Li to B causing the increase in melting & boiling points.
      • C is macromolecule with three dimensional giant structure held by strong covalent bonds.
      • N to Ne are simple molecules held by weak Van Der Waals forces.
    • Acid & Base Characteristics
      • Across the period elements Na, Mg, Al, Si, P, Cl & Ar, change from Base to acidic characteristics.
    • Acid-Base Nature Of Element Oxides
      • Li 2 O react with water to form hydroxide which exhibit a base character.
      • Li 2 O + H 2 O -> 2Li + + 2OH -
      • An anhydride base become base when water is added to it
      • Na 2 O & MgO yield base solution in water.
      • Cl 2 O, SO 2 , P 4 O 10 – acidic solution
      • SiO 2 doesn’t dissolve in water but dissolve slightly in base solution to produce sillicates and form an acidic oxide.
      • END