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Octet rule, lewis structure and formal charge (NOCB)

Octet rule, lewis structure and formal charge (NOCB)

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Octet rule and its Drawback
Rules to draw Lewis structure
Calculation of formal charge

Slide contain information on
Octet rule and its Drawback
Rules to draw Lewis structure
Calculation of formal charge

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Octet rule, lewis structure and formal charge (NOCB)

  1. 1. Nature of Chemical Bonding (NOCB) • Kossel and Lewis theory • Valance Bond theory(VBT) • Molecular Orbital theory(MOT)
  2. 2. KÖSSEL-LEWIS APPROACH TO CHEMICAL BONDING Lewis pictured the atom in terms of a positively charged ‘Kernel’ (the nucleus plus the inner electrons) and the outer shell that could accommodate a maximum of eight electrons This octet of electrons, represents a particularly stable electronic arrangement. Lewis postulated that atoms achieve the stable octet when they are linked by chemical bonds. E.g. F2, O2 etc
  3. 3. Lewis symbol • outer shell electrons take part in chemical combination and they are known as valence electrons. • The inner shell electrons are well protected and are generally not involved in the combination process. • Thus to represent the valence elcetrons lewis developed an representation method called the Lewis symbol
  4. 4. Significance • This number of valence electrons helps to calculate the common or group valence of the element. • The group valence of the elements is generally either equal to the number of dots in Lewis symbols or 8 minus the number of dots or valence electrons. • E.g. Na  has 7e- in valence shell so valency is 1
  5. 5. Kossel- Lewis • First time explained a type of bond that is the electrovalent/ionic bond • They explained how the highly electropositive Grp 1 elements bond with the group 17 electronegative elements. • They bond in order to achieve the stable noble gas configuration • The duplet or octet state in their valence orbitals. • Thus the bonding formed in order to achieve it is called electrovalent bond • The electrovalence is thus equal to the number of unit charge(s) on the ion
  6. 6. Octet rule Kössel and Lewis in 1916 developed an important theory of chemical combination between atoms known as electronic theory of chemical bonding. According to this, atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an octet in their valence shells. This is known as octet rule
  7. 7. Langmuir’s theory of Covalent bond • Each bond is formed as a result of sharing of an electron pair between the atoms. • Each combining atom contributes at least one electron to the shared pair. • The combining atoms attain the outershell noble gas configurations as a result of the sharing of electrons E.g. Formation of Cl2
  8. 8. • If two atoms share two pairs of electrons, the covalent bond between them is called a double bond. For example, in the carbon dioxide molecule • When combining atoms share three electron pairs as in the case of two nitrogen atoms in the N2 molecule and the two carbon atoms in the ethyne molecule, a triple bond is formed
  9. 9. Rules to write Lewis structure – E.g. CO32- 1. Write the symmetrical skeleton for the polyatomic ions O C O 2- C1s22s22p2 O1s22s22p4 O 2. Calculate the number of electrons available in the valence shell of all atoms (A) A= 1x4(C)+3x6(O)+2(for extra 2 e-)=24 e- 3. Calculate the total number of electrons needed by atoms to accqire the noble gas config(N) N=1x8+3x8=32 e- 4. Calculate total number of electrons shared(S) i.e S=N-A S=32-24=8 e- i.e 8/2=4 pairs of electrons
  10. 10. 5. Place the shared electrons in the skeleton. Use = and triple bonds wherever necessary [O:C::O]2- .. O
  11. 11. E.g. Ozone(O3) 1. Skeleton - O1s22s22p4 O O O 2. Calculate A A=3x6=18e- 3. Calculate N N=3x8=24e- 4. S=N-A=24-18=6e- i.e 6/2=3 electron pairs
  12. 12. Bond order • Bond Order is given by the number of bonds between the two atoms in a molecule • Isoelectronic molecules and ions have identical bond orders; for example, F2 and O2 have bond order 1. • N2 , CO and NO+ have bond order 3
  13. 13. Formal charge • The formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure • Formal charge (F.C.) on an atom in a Lewis structure = total number of valence electrons in the free atom — total number of non bonding (lone pair) electrons — (1/2) total number of bonding(shared) electrons • F.C = VE-LP-1/2 x BE
  14. 14. Example
  15. 15. Significance of formal charge • It helps to calculate the charge on the atoms in a Lewis structure • It helps to calculate the number of valence electrons • Formal charges help in the selection of the lowest energy structure from a number of possible Lewis structures for a given species
  16. 16. Drawbacks of octet rule The incomplete octet of the central atom Odd electron molecules Expanded octet
  17. 17. Incomplete octet • In some compounds, the number of electrons surrounding the central atom is less than eight. This is especially the case with elements having less than four valence electrons. Examples are LiCl, BeH2 and BCl3
  18. 18. 2. Odd electron molecules In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxide, NO2 , the octet rule is not satisfied for all the atoms
  19. 19. 3. Expanded octet • Elements in and beyond the third period of the periodic table have, apart from 3s and 3p orbitals, 3d orbitals also available for bonding. • In a number of compounds of these elements there are more than eight valence electrons around the central atom • examples of such compounds are: PF5 , SF6 , H2SO4 and a number of coordination compounds.
  20. 20. Other drawbacks • It is clear that octet rule is based upon the chemical inertness of noble gases. However, some noble gases (for example xenon and krypton) also combine with oxygen and fluorine to form a number of compounds like XeF2 , KrF2 , XeOF2 etc., • This theory does not account for the shape of molecules. • It does not explain the relative stability of the molecules being totally silent about the energy of a molecule.

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