Chemrevise

• Ionic bonding


  The strong electrostatic force of attraction between oppositely charged ions formed by electron transfer
• Formation of ions


  1. Metal atoms lose electrons to form +ve ions

  2. Non-metal atoms gain electrons to form -ve ions
• Electron configuration changes


  • Mg goes from 1s2 2s2 2p63s2 to Mg2+ 1s2 2s2 2p6
  • O goes from 1s2 2s2 2p4 to O2- 1s2 2s2 2p6
• Factors affecting ionic bonding strength

  • Ionic bonding is stronger and the melting points higher 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


  Giant lattices of ions
• Ions with same electronic structure (of the noble gas Ne)


  • N3-
  • O2-
  • F-
  • Na+
  • Mg2+
  • Al3+
• Increasing number of protons


  Ions get smaller
• Going down a group

  Ionic radii increases
• Positive ions


  Smaller compared to their atoms
• Negative ions from groups 5-7

  Larger than the corresponding atoms
• Physical properties of ionic compounds

  • High melting points
  • Non conductor of electricity when solid
  • Conductor of electricity when in solution or molten
  • Brittle / easy to cleave apart
• Migration of ions in solution
• Covalent bonding


  A strong bond caused by the electrostatic attraction between the bonding shared pair of electrons and the two nuclei
• Significant electron density between atoms in covalent compounds
• Effect of multiple bonds

  • Greater electron density between nuclei
  • Greater force of attraction
  • Shorter bond length
  • Greater bond strength
• Dative covalent bond


  The shared pair of electrons comes from only one of the bonding atoms
• Common examples of dative covalent bonds


  • NH4+, H3O+, NH3BF3
• Dative covalent bond acts like an ordinary covalent bond when thinking about shape
• Example of dative covalent bonding


  • Two AlCl3 molecules join together through two dative bonds to form Al2Cl6
• Molecular shapes


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


  The relative tendency of an atom in a covalent bond to attract electrons to itself
• Across a period


  Electronegativity increases as atomic radius decreases
• Down a group


  Electronegativity decreases as atomic radius increases
• Polar covalent bond


  Bond with unequal distribution of electrons, producing a charge separation (dipole)
• Symmetric molecules are non-polar even if individual bonds are polar
• Ionic and covalent bonding are extremes of a continuum of bonding type
• London forces


  Instantaneous, induced dipole-dipole interactions occurring between all simple covalent molecules and noble gas atoms
• Factors affecting London forces


  • More electrons in molecule increases chance of temporary dipoles forming, making London forces stronger
  • Shape of molecule affects surface area of contact between molecules
• Permanent dipole-dipole forces


  Stronger than London forces, occur between polar molecules
• Hydrogen bonding


  Occurs between a hydrogen atom attached to N, O or F which has a lone pair of electrons