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 is stronger and the melting points higher


  When the ions are smaller and/ or have higher charges
• MgO has a higher melting point than NaCl
• 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


  Because as one goes down the group the ions have more shells of electrons
• 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


  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
• 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
• The three main factors that affect the strength of metallic bonding

  • Number of protons/ Strength of nuclear attraction
  • Number of delocalised electrons per atom
  • Size of ion
• Mg has stronger metallic bonding than Na and hence a higher melting point
• Molecular shape


  • 10 electrons made up of 4 bond pairs and 1 lone pair
  • Variation of the 5 bond pair shape (trigonal bipyramidal)

  
• Square planar


  • Bond angle 90°
• Linear


  • Bond angle 180°
• Bent


  • Bond angle ~119° + 89° (Reduced by lone pair)
• Electronegativity


  The relative tendency of an atom in a covalent bond 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 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)
• Symmetric molecule


  • All bonds identical and no lone pairs
  • Individual dipoles on the bonds 'cancel out' due to the symmetrical shape of the molecule
  • No net dipole moment: the molecule is non-polar
• Pauling scale


  Electronegativity is measured on this scale (ranges from 0 to 4)
• Ionic and covalent bonding


  Extremes of a continuum of bonding type
• Intermediate bonding


  Differences in electronegativity between elements can determine where a compound lies on this scale
• Van der Waals' forces

  Transient, induced dipole-dipole interactions that occur between all simple covalent molecules and the separate atoms in noble gases
• Factors affecting size of Van der Waals

  • More electrons in the molecule increases the chance of temporary dipoles forming, making the Van der Waals stronger and boiling points 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 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
• Substances that can form hydrogen bonds

  • Alcohols
  • Carboxylic acids
  • Proteins
  • Amides