Covalent Bonding

• Is a bond in which bonding atoms share \(\ce{e−}\), rather than transferring them
  • One \(\ce{e−}\) from each bonding atom is contributed, with more \(\ce{e−}\) changing the properties/bond number/BO
• AOs are distorted into a shape which will allow them to contribute to both nuclei
• Can be defined by \(\Delta\chi>2\), however this has the same issues as ionic bonding in that bonding character is not binary
• Covalent compounds (molecules) are held together by intramolecular forces

• Typically a bond that occurs between non-metals

  • H can both ionically and covalently bond

Bond Enthalpy (\(\Delta H^\circ\))

• The energy required to break a particular bond in 1 mol of gaseous molecules
  • The molecules have to already be in a gaseous state, otherwise intermolecular forces will contribute to stabilising the bond, increasing the energy
    

• Bond energies can be directly measured for diatomic molecules, as it can be related to the thermal decomposition of the molecule itself
• The more bonds (double, triple), the stronger the bond energy Can also be effected by the size/\(\Delta\chi\) of the atoms involved

Reactions

• When a reaction occurs:

\[ \Delta E=\sum \text{ bonds broken }−\sum\text{ bonds formed} \]

• E.g. \(\hskip{1cm}\ce{2H2_{(g)} + O2_{(g)} -> 2H2O{(l)}}\hskip{2cm}\Delta H=−483.6\:kJ\cdot mol^{−1}\)
  • In this reaction, the bonds of 2 moles of \(\ce{H2}\) and 1 mole of \(\ce{O2}\) are broken
  • However two H-O bonds are formed
  • The excess of energy released causes the exothermic reaction
  • Using the standard bond enthalpies table, this is estimated at \(\sim−450\:kJ\cdot mol^{−1}\)

Types of covalent bonds

• Covalent bonds can either be polar or none polar
• Non polar bonds equally share their electrons between the two atoms and only really happens in homodiatomic molecules
• Polar bonds do not share electrons equally, which is typically based on \(\Delta\chi\)

Types of intermolecular forces

Bond Type	Energy (\(kJ\cdot mol^{-1}\))
Dispersion	0.05-40
Dipole-induced dipole	2-10
Ion-induced dipole	3-15
Dipole-dipole	5-25
H-Bond	10-40
Ion-Dipole	40-600
For Comparison
Covalent bond	150-1100
Ionic bond	400-4000

Dipole Moment (unit: Debye)

• A dipole moment is a charge difference, based on the distribution of an electron cloud
  
• It consists of two point charges and is a vector quantity (both magnitude and direction)
• This can be over the entire molecule, or just between individual bonds

Dipole-dipole forces

• Are simple electrostatic forces between molecules that help them stick together
  

Induced dipoles

• Since dipoles have a charge, they can interact with \(\ce{e−}\) of non polar compounds
• This can distort their \(\ce{e−}\) distribution enough to form a temporary dipole of their own
• This can also happen with ions, causing a stronger dipole to be formed, since an ionic charge is a full, formal charge, instead of a partial dipole charge

H-bonds

• Are just particularly strong dipoles and occur when a particularly electronegative atom with lone pairs interacts with a hydrogen atom

Dispersion forces

• Dispersion forces are the weakest of all the bonding forces, but can be the only thing keeping a nonpolar species intact (pentane, liquid gasses)
• Since atoms are not uniformly charged, even non charged, inert species can influence the \(\ce{e−}\) distibution around them. This minute, momentary interaction can cause a tiny dipole to exist, that can create a tiny dipole in surrounding atoms that will allow for intermolecular attraction to exist
• They exist between all particles, but are almost always overshadowed by other forces
• The presence of dispersion forces can be effected by:
  • The polarizability of the compound (how easily the electrons can move around)
    • Increases with molecular weight and atomic radii
  • The size of the molecule (larger molecules allow electrons to move more freely, allowing them to build up on one side of the molecule)