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 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 of dative covalent bond


  • NH4+, H3O+, NH3BF3
• Metallic bonding


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

  • Number of protons/Strength of nuclear attraction
  • Number of delocalised electrons per atom
  • Size of ion
• Mg has a higher melting point than Na due to its stronger metallic bonding
• Types of bonding and structure

  • Ionic: electrostatic force of attraction between oppositely charged ions
  • Covalent: shared pair of electrons
  • Metallic: electrostatic force of attraction between the metal positive ions and the delocalised electrons
• Examples of ionic, molecular, macromolecular, and metallic structures

  • Ionic: Sodium chloride, Magnesium oxide
  • Molecular: Iodine, Ice, Carbon dioxide, Water, Methane
  • Macromolecular: Diamond, Graphite, Silicon dioxide, Silicon
  • Metallic: Magnesium, Sodium
• Only use the words molecules and intermolecular forces when talking about simple molecular substances
• Properties of different bonding types

  • Ionic: High boiling/melting points, Generally good solubility in water, Poor conductivity when solid, Good conductivity when molten
  • Molecular (simple): Low boiling/melting points, Generally poor/insoluble, Poor conductivity
  • Macromolecular: High boiling/melting points, Insoluble, Poor conductivity (except graphite)
  • Metallic: High boiling/melting points, Good conductivity
• Metals are malleable as 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
• 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)