Bonding

• Metal atoms


  Lose electrons to form +ve ions
• Non-metal atoms


  Gain electrons to form -ve ions
• Mg


  Goes from 1s2 2s2 2p63s2 to Mg2+ 1s2 2s2 2p6
• O


  Goes from 1s2 2s2 2p4 to O2- 1s2 2s2 2p6
• Ionic bonding


  • Stronger and higher melting points when the ions are smaller and/or have higher charges
  • E.g. MgO has higher melting point than NaCl as the ions (Mg2+ & O2-) are smaller and have higher charges than those in NaCl (Na+ & Cl-)
• Ionic crystals


  Have the structure of giant lattices of ions
• Ionic radii


  • N3-
  • O2-
  • F-
  • Na+
  • Mg2+
  • Al3+
• N3-, O2-, F- and Na+, Mg2+, Al3+ all have the same electronic structure (of the noble gas Ne)
• There are increasing numbers of protons from N to F and then Na to Al but the same number of electrons
• The effective nuclear attraction per electron therefore increases and ions get smaller
• Ionic radii


  Increase going down a group because the ions have more shells of electrons
• Positive ions


  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
• Negative ions


  Larger than the corresponding atoms because the negative ion has more electrons than the corresponding atom but the same number of protons, so the pull of the nucleus is shared over more electrons and the attraction per electron is less, making the ion bigger
• Common examples of dative covalent bond


  • NH4+, H3O+, NH3BF3
• Factors affecting strength of metallic bonding

  • Number of protons/Strength of nuclear attraction
  • Number of delocalised electrons per atom
  • Size of ion
• Higher energy is needed to break bonds in Mg compared to Na due to stronger metallic bonding
• Electronegativity


  The relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself
• Most electronegative atoms


  • F
  • O
  • N
  • Cl
• Factors affecting electronegativity


  • Increases across a period as the number of protons increases and the atomic radius decreases
  • Decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases
• Purely covalent bond


  Compound containing elements of similar electronegativity and hence a small electronegativity difference
• Polar covalent bond


  Bond that forms when the elements in the bond have different electronegativities (of around 0.3 to 1.7)
• The element with the larger electronegativity in a polar compound will be the δ- end
• Ionic bond


  Compound containing elements of very different electronegativity and hence a very large electronegativity difference (> 1.7)
• A symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecular are polar
• The individual dipoles on the bonds 'cancel out' due to the symmetrical shape of the molecule. There is no net dipole moment: the molecule is non-polar
• Pauling scale


  Electronegativity is measured on this scale which ranges from 0 to 4, with fluorine being the most electronegative element at 4.0
• Ionic and covalent bonding

  The extremes of a continuum of bonding type. Differences in electronegativity between elements can determine where a compound lies on this scale.
• Van der Waals' forces

  Also called transient, induced dipole-dipole interactions. They occur between all simple covalent molecules and the separate atoms in noble gases.
• Van der Waals' forces

  • The more electrons there are in the molecule, the higher the chance that temporary dipoles will form, making the Van der Waals stronger between the molecules and so boiling points will be greater.
• Permanent dipole-dipole forces


  Stronger than Van der Waals and so the compounds have higher boiling points. Occurs between polar molecules that have a permanent dipole.
• Hydrogen bonding


  Occurs in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine, which must have an available lone pair of electrons.
• Hydrogen bonding is stronger than the other two types of intermolecular bonding.
• The anomalously high boiling points of H2O, 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 number of electrons.
• Substances that can form hydrogen bonds

  • Alcohols
  • Carboxylic acids
  • Proteins
  • Amides
• Four types of crystal structure

  • Ionic
  • Metallic
  • Molecular
  • Giant covalent (macromolecular)
• Ionic crystal structure


  • Sodium chloride
• Metallic crystal structure


  • Magnesium or sodium
• Molecular crystal structure


  • Iodine
  • Ice
• Macromolecular crystal structure


  • Diamond
  • Graphite