Organic chemistry

• Tetravalency in carbon


  Carbon can form four bonds
• Orbital hybridization


  Mixing of atomic orbitals on the same atom to give new equivalent hybrid atomic orbitals
• Formation of C-H and C-C sigma bond

  1. In-phase overlap of half-filled orbitals

  2. Overlap along internuclear axis

  3. Gives a sigma (σ) bond
• Formation of C-C pi-bond

  1. Side-by-side overlap of half-filled p orbitals

  2. Gives a pi (π) bond
• Sigma (σ) bond


  • Covalent bond formed by end-on overlap of orbitals
  • Cylindrically symmetrical
• Pi (π) bond


  • Covalent bond formed by side-ways overlap of parallel orbitals
• Polar covalent bond


  Bond with unequal sharing of electrons due to electronegativity difference
• Methane has a tetrahedral structure with bond angles of 109.5° and bond distances of 110 pm
• Ground state electron configuration of carbon has two unpaired electrons in the 2s and 2p orbitals
• Excited state electron configuration of carbon

  Promotion of an electron from 2s to 2p orbital
• sp3 hybridization


  Mixing of 2s and three 2p orbitals to give four equivalent half-filled hybrid orbitals
• sp3 hybridized orbitals are consistent with four bonds and tetrahedral geometry
• Shape of sp3 hybrid orbitals

  • Tetrahedrally arranged
  • Reinforcement of electron wave in regions of same sign
  • Destructive interference in regions of opposite sign
• Formation of C-H sigma bond in methane

  In-phase overlap of half-filled 1s orbital of hydrogen with half-filled sp3 hybrid orbital of carbon
• Justification for sp3 hybridization in methane: consistent with tetrahedral structure, allows four sigma bonds, and gives stronger bonds than s-s or p-p overlap
• sp2 hybridization


  Mixing of 2s and two 2p orbitals to give three equivalent sp2 hybrid orbitals plus one unhybridized 2p orbital
• Ethane has a tetrahedral geometry at each carbon with C-H bond distance of 110 pm and C-C bond distance of 153 pm
• Formation of C-C sigma bond in ethane

  In-phase overlap of half-filled sp3 hybrid orbitals of two carbon atoms
• Ethylene has a planar structure with bond angles close to 120° and C=C bond distance of 134 pm
• Formation of C-C pi-bond in ethylene

  Side-by-side overlap of half-filled p orbitals of two carbon atoms
• sp2 orbitals


  Hybrid orbitals formed by mixing 2s and 2p orbitals
• sp2 orbitals


  • Three equivalent half-filled sp2 hybrid orbitals
  • One unhybridized p orbital
• Sigma (σ) bond


  Cylindrically symmetrical bond formed by the overlap of sp2 orbitals
• Pi (π) bond


  Bond formed by the overlap of p orbitals, not cylindrically symmetrical
• Formation of carbon-carbon double bond in ethene


  1. One sigma bond from overlap of sp2 orbitals

  2. One pi bond from overlap of p orbitals
• The pi bond prevents the double bond from rotating, making ethene a planar molecule
• sp hybridization


  Hybrid orbitals formed by mixing 2s and one 2p orbital
• sp hybridization


  • Two equivalent half-filled sp hybrid orbitals
  • Two unhybridized p orbitals
• Formation of carbon-carbon triple bond in acetylene

  1. One sigma bond from overlap of sp orbitals

  2. Two pi bonds from overlap of p orbitals
• The carbon-carbon triple bond in acetylene is described as one sigma and two pi bonds
• Electrons in covalent bonds are not necessarily shared equally by the two atoms
• Polar covalent bond


  Bond where the electron distribution is polarized due to one atom having a greater tendency to attract electrons
• Polar compounds


  • H2O
  • HCl
  • CH3Cl
• Non-polar compounds


  • CH4
  • C2H6
• Electronegativity


  The tendency of an atom to draw the electrons in a covalent bond toward itself
• Electronegativity increases across a row in the periodic table and decreases down a column
• Dipole


  Separation of positive and negative charge in a molecule
• Dipole moment


  The product of the charge and the distance between the centers of charge
• Polar compounds have stronger intermolecular forces and higher melting/boiling points than non-polar compounds
• Polar compounds are soluble in polar solvents, while non-polar compounds are soluble in non-polar solvents