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 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
• This is because as one goes down the group 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 formed from groups five to seven


  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 the 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 the stronger metallic bonding
• Types of bonding and structure

  • Ionic
  • Covalent (simple molecular, macromolecular)
  • Metallic
• Properties of different bonding types

  • Boiling and melting points
  • Solubility in water
  • Conductivity (solid and molten)
  • General description
• Metals are malleable because the positive ions in the lattice are all identical, so the planes of ions can slide easily over one another
• Molecular shapes


  • Linear
  • Trigonal planar
  • Tetrahedral
  • Trigonal pyramidal
  • Bent
  • Trigonal bipyramidal
  • Octahedral
• Electronegativity


  The relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself
• F, O, N and Cl are the most electronegative atoms
• Electronegativity increases across a period as the number of protons increases and the atomic radius decreases
• Electronegativity decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases
• A compound can have intermediate bonding between ionic and covalent depending on the electronegativity difference between the atoms
• 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 90O
• Linear


  • Bond angle 180O
  • Bond angle ~89O (Reduced by lone pairs)
• Bent


  • Bond angles ~119 + 89O (Reduced by lone pair)
• 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
• Polar covalent bond


  A bond that forms when the elements in the bond have different electronegativities (of around 0.3 to 1.7)
• Ionic bond


  A bond that forms when the elements in the bond have 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
• 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

  • The more electrons there are in the molecule the higher the chance that temporary dipoles will form, making the Van der Waals stronger
• Permanent dipole-dipole forces


  Stronger than Van der Waals, occur between polar molecules
• 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
• Increasing boiling points of alkanes

  Caused by increasing number of electrons in bigger molecules, increasing Van der Waals
• Increasing boiling points of halogens down group 7

  Caused by increasing number of electrons in bigger molecules, increasing Van der Waals