chp#03

• Valence Shell Electron Pair Repulsion (VSEPR) Theory


  Theory that the shapes or geometry of a molecule or ion depends on the number of shared pairs as well as the lone pairs of electrons around the central atom

  
• VSEPR theory was suggested by Sidgwick and Powell


  1940
• VSEPR Theory


  • Arrangement of electrons around central polyvalent atom is as far apart as possible to minimise repulsion
  • Non-bonding (lone) pairs occupy more space than bonding pairs
  • Both lone pairs and bond pairs determine molecular geometry
  • Magnitude of repulsion: lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
  • Multiple bonds behave like single bonds in determining geometry

  
• Central Atom


  Appears one time in formula, lone pairs determined for central atom
• How to find lone pairs

  Free electrons = Total valence electrons - Electrons used in bonding

  
• Examples of molecules using VSEPR Theory

  • CO2
  • H2O
  • SiCl2
  • SO3
  • NH3

  
• AB2 molecules without lone pair

  • Linear geometry with 180° bond angle
  • Multiple bonds act like single bonds

  
• AB2 molecules with lone pair

  • AB2E1: Angular/V-shaped with bond angle less than 120°

  AB2E2: V-shaped with bond angle around 104.5°

  
• AB3 molecules without lone pair

  • Planar triangular geometry with 120° bond angle
  • Multiple bonds don't affect geometry

  
• AB3 molecules with lone pair

  • Triangular pyramidal geometry with bond angle around 107.5°

  
• AB4 molecules


  • Regular tetrahedral geometry with 109.5° bond angle

  
• AB5 molecules


  • Trigonal bipyramidal geometry

  
• AB6 molecules


  • Octahedral geometry

  
• Valence Bond Theory (VBT)


  Theory based on wave-mechanical treatment of molecules, explains bond energies, lengths and shapes

  
• VBT was proposed by Heitler and London in 1927 and developed by Pauling
• Postulates of VBT


  • Bond formation by overlap of half-filled atomic orbitals
  • Electrons in overlapping orbitals have opposite spin
  • Overlapping orbitals must have same symmetry w.r.t. bond axis
  • Number of bonds = number of unpaired electrons in valence shell
  • Greater overlap = stronger bond
  • Single bond = 2 orbitals overlap, multiple bonds = additional orbitals overlap
  • Sigma (σ) bond = head-on overlap, Pi (π) bond = parallel/sideways overlap

  
• Sigma (σ) bond


  Bond formed by head-on overlap of atomic orbitals, region of highest electron density is symmetrically distributed around bond axis

  
• Pi (π) bond


  Bond formed by sideways/parallel overlap of already σ-bonded p-orbitals

  
• Sigma bond


  Bond where the region of highest electron density is symmetrically distributed around the bond axis

  
• The first bond formed between any two atoms is the sigma bond

  
• All single covalent bonds are sigma bonds (σ) and the electrons occupying a bond are called σ electrons

  
• Types of sigma bonds

  • S-S Overlapping
  • S-Px | S-Py | S-Pz
  • Px-Px Overlapping

  
• Pi bond


  Bond formed between two already (σ) bonded atoms by the sidewise overlap of their two half-filled p-atomic orbitals whose axes are parallel

  
• Pi bond formation


  • The two overlapping p-orbitals must be coplanar and their axes must be parallel
  • It is formed by side wise or lateral overlap between two p-orbitals which have their lobes perpendicular to the molecular axis

  
• Electron density in pi bond

  • The electron density is unsymmetrical around the bond axis
  • The probability of finding the electron is maximum in the region above and below the line joining the two nuclei

  
• Overlapping extent in pi bond

  • The overlap of p-orbitals in pi bond formation is not as good as in sigma bond. That's why sigma bond is always stronger than a pi bond

  
• Orbitals that make pi bonds

  • Py-Py
  • Pz-Pz
• In case of pi bond formation, the electron density is greatest above and below the line joining the two nuclei and this is also called nodal plane

  
• Only pure, parallel, co-planar, half-filled p-orbitals on adjacent atoms can form a pi bond

  
• Only one bond in any multiple bonds can be a sigma bond, the remaining bonds are pi bonds

  
• In case of hybridization, the overlapping of any hybrid-orbitals always produces a sigma bond

  
• Differences between sigma and pi bonds

  • Definition
  • Symbol
  • Number
  • Extent of Overlapping
  • Strength
  • Density
  • Types of overlapping

  
• Formation of H2 molecule


  1. Each H atom has the electron configuration 1st

  2. As two hydrogen atoms approach each other, their half-filled 1st orbitals overlap, giving H-H bond

  3. The overlap of orbitals provides a means for sharing electrons, thereby allowing each 1st to complete its valence shell

  4. The electron density is concentrated in the region along the line joining the two nuclei

  5. The bond formed is a sigma (a) bond

  
• Formation of Cl2 molecule

  1. Each Chlorine atom has one half-filled 3pz orbitals

  2. Sigma bond is formed between two Cl atoms by head on overlap of half-filled 3pz atomic orbital of each chlorine atom

  
• Formation of HF molecule


  The half-filled 1s orbital of H atom overlaps with the half-filled 2pz orbital of F to form a bond

  
• Formation of O2 molecule

  1. One bond is formed by the end-to-end overlap of half-filled 2px orbitals on each oxygen atom. This gives a σ bond

  2. The second bond is formed by the side-to-side overlap of half-filled 2py orbitals on each oxygen atom. This gives π bond

  3. Thus a double bond is formed between two oxygen atoms

  
• Formation of N2 molecule

  1. When the two nitrogen atoms approach each other, their 2px orbitals undergo end-to-end overlap, giving a σ bond

  2. The end-to-end overlapping brings the two nitrogen atoms so close together that their parallel 2p orbitals undergo side-to-side overlap to produce two pi bonds

  3. Thus a triple bond is formed between two nitrogen atoms; one is σ bond while the other two are π bonds

  
• Hybridization


  A process of mixing atomic orbitals of different energy and shape to form set of new orbitals of the same energy and same shape
• Hybridization


  • Mixing orbitals have slight difference in energy and different shapes
  • Hybridized orbitals have same energies and same shapes
  • In this process, the electron belonging to the ground state structure are promoted to the excited state. It increases number of unpaired electrons

  
• Types of hybridization


  • sp3
  • sp2
  • sp