Bonding

• Metal atoms


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


  Gain electrons to form -ve ions
• Mg goes from


  1. 1s2 2s2 2p63s2

  2. to Mg2+ 1s2 2s2 2p6
• O goes from


  1. 1s2 2s2 2p4

  2. 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 a higher melting point than NaCl as the ions involved (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
• Within a group the size of the ionic radii

  Increases going down the group
• 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
• The negative ions formed from groups five to seven are larger than the corresponding atoms
• 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
• Ionic bonding


  The electrostatic force of attraction between oppositely charged ions formed by electron transfer
• Covalent bond


  A shared pair of electrons
• Dative covalent bond


  The shared pair of electrons in the covalent bond come from only one of the bonding atoms
• Common examples you should be able to draw that contain dative covalent bond

  • NH4+, H3O+, NH3BF3
• Metallic bonding


  The electrostatic force of attraction between the positive metal ions and the delocalised electrons
• Factors affecting the strength of metallic bonding

  • Number of protons/ Strength of nuclear attraction
  • Number of delocalised electrons per atom
  • Size of ion
• Ionic bonding is the electrostatic force of attraction between oppositely charged ions formed by electron transfer
• A covalent bond is a shared pair of electrons
• A dative covalent bond forms when the shared pair of electrons in the covalent bond come from only one of the bonding atoms
• Metallic bonding is the electrostatic force of attraction between the positive metal ions and the delocalised electrons
• 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)
• Ionic bond


  Compound containing elements of very different electronegativity and hence a very large electronegativity difference (> 1.7)
• Electronegativity is measured on the Pauling scale (ranges from 0 to 4)
• Fluorine is the most electronegative element and is given a value of 4.0
• Ionic and covalent bonding


  Extremes of a continuum of bonding type
• Van der Waals' forces

  • Also called transient, induced dipole-dipole interactions
  • Occur between all simple covalent molecules and the separate atoms in noble gases
  • Caused by fluctuations in electron density creating temporary dipoles that induce dipoles in neighboring molecules
• Factors affecting size of Van der Waals

  • 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
  • The shape of the molecule can also affect the size of the Van der Waals forces, with long chain alkanes having a larger surface area of contact between molecules for Van der Waals to form than compared to spherical shaped branched alkanes
• 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 addition to van der waals forces
  • Stronger than the other two types of intermolecular bonding
  • Causes the anomalously high boiling points of H2O, NH3 and HF
• Types of crystal structure

  • Ionic
  • Metallic
  • Molecular
  • Giant covalent (macromolecular)
• Ionic: sodium chloride


  • Giant ionic lattice showing alternate Na+ and Cl- ions
• Metallic: magnesium or sodium


  • Giant metallic lattice showing close packing magnesium ions