3 Bonding

• Ionic bonding is the electrostatic force of attraction between oppositely charged metal and non-metal ions formed by electron transfer to form giant lattice structures
• Metal atoms lose electrons to form positive ions.

  Non-metal atoms gain electrons to form negative ions.
• Ionic bonding is stronger when the ions are smaller and/or have higher charges
• A covalent bond is a shared pair of electrons between 2 non-metal atoms
• A dative covalent bond forms when the shared pair of electrons in the covalent bond come from only one of the bonding atoms.
• Dative covalent bonding can also be called co-ordinate bonding
• In a dative covalent bond, the direction of the arrow goes from the atom that is providing the lone pair to the atom that is deficient.
• Metallic bonding is the electrostatic force of attraction between the positive metal ions and the delocalised electrons
• 3 main factors that affect the strength of metallic bonding:

  1. Number of protons/Strength of nuclear attraction
  2. Number of delocalised electrons per atom
  3. Size of ion (smaller = stronger)
• Covalent bonding can form:

  1. Simple molecular structures with intermolecular forces between molecules
  2. Giant molecular structures/Macromolecular
• Metallic bonding forms giant metallic lattice structures
• Ionic compounds:

  1. Have high melting and boiling points because of the giant lattice of ions with strong electrostatic forces between oppositely charged ions.
  2. Are soluble in water
  3. Have poor conductivity when solid as the ions are fixed and cannot move
  4. Have good conductivity when molten or dissolved in water as the ions can move
  5. Are crystalline solids
• Simple molecular compounds:

  1. Have low melting and boiling points because of the weak intermolecular forces between molecules e.g. Van der Waals
  2. Are usually not soluble in water
  3. Have poor conductivity as there are no ions or delocalised electrons to carry charge throughout the structure
  4. Are mostly gases and liquids at room temperature
• Giant molecular compounds:

  1. Have high melting and boiling points because of the many strong covalent bonds in the structure which take a lot of energy to break
  2. Are insoluble in water
  3. Diamond and sand have poor conductivity as the electrons cannot move but graphite has good conductivity as there are free delocalised electrons between layers
  4. Are solids at room temperature
• Metallic compounds:

  1. Have high melting and boiling points because of the strong electrostatic forces between the positive ions and the sea of delocalised electrons.
  2. Are insoluble in water
  3. Have good conductivity as the delocalised electrons can move through the structure
  4. Are shiny metals and malleable as all the ions are identical so the layers can slide easily over one another
• Electronegativity is an atom's ability to attract a pair of electrons in a covalent bond to itself.
• The most electronegative element is fluorine (4.0)
• A symmetrical molecule (all bonds identical and no lone pairs) will never be polar e.g. CO2
• (V + B - C) / 2

  V = Valency or group number of central atom

  B = Bonds around the central atom

  C = Charge
• Van der Waals forces occur between all simple covalent molecules and the separate atoms in noble gases
• Van der Waals forces are also called transient or induced dipole-dipole interactions
• Van der Waals

  • Since the electrons are constantly moving randomly, the electron density can fluctuate and parts of the molecule become more or less negative so small temporary dipoles form
  • These can cause dipoles to form in neighbouring molecules which are called induced dipoles
  • Induced dipoles are always the opposite sign to the original one
• Van der Waals:

  • The larger the molecule, the more electrons there are in the molecule, the higher the chance that temporary dipoles will form so the Van der Waals forces are stronger between the molecules and so boiling points will be higher.
• Long chain alkanes have a larger surface area of contact between molecules for Van der Waals to form than compared to spherical shaped branched alkanes and so have stronger Van der Waals
• Permanent dipole-dipole forces occur between polar molecules and are stronger than Van der Waals so the compounds have higher boiling points.
• Polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms so they have a permanent dipole
• Hydrogen bonding occurs in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of fluorine, oxygen and nitrogen, which must have an available lone pair of electrons.
• When drawing permanent dipole-dipole forces, draw the polarities too.

  
• When drawing hydrogen bonding, show the lone pair and the polarities:

  
• Hydrogen bonding is stronger than the other two types of intermolecular bonding.
• • The anomalously high boiling points of H20, NH3, and HF are caused by the hydrogen bonding between the molecules
  • The general increase in boiling point from H2S to H2Te is caused by increasing Van der Waals forces between molecules due to an increasing size of molecule and so increasing number of electrons

  
• Hydrogen bonds can form in alcohols, carboxylic acids, proteins and amides
• This is a giant ionic lattice structure:

  
• This is a giant metallic lattice structure:

  
• This is iodine, with weak Van der Waals forces between the molecules

  
• This is the structure of ice in a tetrahedral arrangement:

  
• The molecules in ice are held further apart than in liquid water so the density of ice is lower than that of water
• This is the macromolecular structure of diamond with a tetrahedral arrangement of 4 covalent bonds per carbon atom:

  
• This is the macromolecular structure of graphite in planar arrangements of 3 covalent bonds per carbon atom in layers. The 4th outer electron per atom is delocalised between layers.

  
• Positive ions are smaller compared to their atoms because it has one less shell of electrons and the ratio of protons to electrons has increased so there is greater net force on remaining electrons holding them more closely.