Bonding

• Metal atoms


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


  Gain electrons to form -ve ions
• Mg electronic configuration changes

  Mg goes from 1s2 2s2 2p63s2 to Mg2+ 1s2 2s2 2p6
• O electronic configuration changes

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


  • Stronger and higher melting points when ions are smaller and/or have higher charges
  • E.g. MgO has higher melting point than NaCl as Mg2+ & O2- are smaller and have higher charges than 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 as 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 as 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 bonds


  • 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
• Higher energy is needed to break the stronger metallic bonds in Mg compared to Na, hence Mg has a higher melting point
• Types of bonding and structure

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

  • Boiling and melting points
  • Solubility in water
  • Conductivity when solid
  • Conductivity when molten
  • General description
• 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
• Trigonal bipyramidal


  Variation of the 5 bond pair shape
• Square planar


  • Bond angle 90°
• Linear


  • Bond angle 180°
• Bent


  • Bond angle ~119° + 89° (Reduced by lone pair)
• This makes a total of 10 electrons made up of 4 bond pairs and 1 lone pair
• Bond angle ~89° (Reduced by lone pair)
• Electronegativity is 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
• Electronegativity increases across a period

  As the number of protons increases and the atomic radius decreases, the electrons in the same shell are pulled in more
• Electronegativity decreases down a group


  The distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases
• Purely covalent bond


  When a compound contains elements of similar electronegativity and hence a small electronegativity difference
• Polar covalent bond


  When a bond is formed between elements with different electronegativities (of around 0.3 to 1.7)