Intermolecular Forces of Attraction - Forces of attraction that exist between molecules

Network (Covalent) Bonds - Strongest strong bond

• In a network solid, atoms are covalently bonded in a lattice structure.
• Network solids are very hard and their strong bonds result in very high melting and boiling point temperatures.
• Network solids have localized electrons which are in fixed positions in the covalent bonds.  This makes network solids poor conductors of electricity.
• Examples of network solids include
  • Diamond - a form of carbon in which its electrons have an sp³ hybridization and the atoms are arranged in a tetrahedral network.
  • Quartz - SiO₂
  • Silicon carbide - SiC

Ionic Bonds (between ionic molecules) - Medium strong bond

• In an ionic solid, adjacent ions have electrostatic attractions and are arranged in a lattice structure.
• Ionic solids have strong bonds resulting in high melting and boiling point temperatures.
• Ionic solids have localized electrons which are in fixed positions around the atoms.  This makes ionic solids poor conductors of electricity.
  • Ionic liquids however (e.g. molten NaCl) do conduct electricity because the electrons are free to move about.
• Coulomb's Law states that the force of attraction between two objects is equal to the product of their charges  (q₁ and q₂) divided by the square of the distance between them (r):

  

  • Therefore, the charges on the ions in an ionic solid are directly proportional to the strength of the ionic bonds.
  • The sizes of the ions in an ionic solid are inversely proportional to the strength of the ionic bonds

    e.g.  The ionic bonds in MgO (charges +2 and -2) are stronger than those in NaF (charges +1 and -1)
    e.g.  The ionic bonds in LiF are stronger than those in KBr (because LiF is much smaller than KBr)

Metallic Bonds (bonds between metal atoms) - Weakest strong bond

• Metallic solids are often described as "a group of nuclei surrounded by a sea of mobile electrons."  In other words, the electrons in metallic solids are delocalized and are free to move about, making metals good conductors of electricity.
• Metallic bonds are still relatively strong forces of attraction and most metals have high melting and boiling point temperatures.
  • Mercury is the only metal that is not solid at room temperature.
• Though most metals are very hard, the unrigid structure of the electrons makes them malleable and ductile.
• Like in ionic solids, Coulomb's Law can be used to describe the relative strengths of metallic solids:

  

  • The size of the metal atoms is inversely proportional to the strength of the metallic bonds.  In other words, the smaller the atoms, the stronger the force.  Smaller size allows for the positively charged nucleus of one metal atom to be closer to the negatively charged electrons of another, increasing the strength of the attraction between them.

Hydrogen Bonding - Strongest weak bond

• In hydrogen bonds, the positively charged hydrogen end of one molecule is attracted to the negatively charged end of another molecule which must be an extremely electronegative element (fluorine, oxygen, or nitrogen - FON)
  • e.g.  H₂O, HF, and NH_3.
  • Alcohols can also form hydrogen bonds with water in which the O in the OH group of the alcohol bonds to the positively charged H end of the water molecule and the H in the OH group of the alcohol bonds to the negatively charged O of the water molecule.

  
  

• Hydrogen bonds are the strongest weak bond because the H atom essentially gives its single electron to form a bond and is therefore left unshielded.  The relatively strength of hydrogen bonds results in higher melting and boiling point temperatures than those in molecules with other van der Waals forces of attraction.
• Hydrogen bonding can explain why water is less dense in the solid phase than it is in the liquid phase (contrary to most other substances).  The hydrogen bonds between water molecules in ice to form a crystal structure, keeping them further apart than they are in the liquid phase.

Dipole-Dipole Forces - Medium weak bond

• Dipole-dipole forces exist between neutral, polar molecules where the positive end of one molecule is attracted to the negative end of another molecule.

• The greater the polarity (difference in electronegativity of the atoms in the molecule), the stronger the dipole-dipole attraction.

• Dipole-dipole attractions are very weak and substances held together by these forces have low melting and boiling point temperatures. Generally, substances held together by dipole-dipole attractions are gases at room temperature.

London Dispersion Forces - Weakest weak bond

• London dispersion forces (LDF) occur between neutral, nonpolar molecules.  LDF occur due to the "random motion of electrons."  At any moment, one atom may be surrounded by an extra electron from a neighboring atom, resulting in an instantaneous polarity on the atom.  During that instant, the "polarized" atom will act as a very weak dipole.
• Since LDF are dependent upon the random motion of electrons, the more electrons an atom or molecule has, the greater the LDF between them.
• LDF as a whole are extremely weak, so substances held together by these forces have extremely low melting and boiling point temperatures.  These substances tend to be gases at room temperature.

Next:  "Gases:  Properties of Gases"