Bonding, Structure, Properties and Energy

• Lewis diagram


  Represents the arrangement of valence electrons in a molecule

  
• Octet rule


  • Every atom has 8 electrons (a full shell)
  • H atom has only 2 electrons in its shell
• Bonded pairs of electrons


  • Shared electrons that contribute to the covalent bond
  • Can be marked with a straight line (each line represents 2 electrons) or a pair of dots / crosses
• Lone pairs of electrons

  • Valence electrons not involved in covalent bonding but must be included in the diagram
  • Marked with a pair of dots /crosses
• Valence electrons in a stable configuration - such as a molecule - are always paired. Therefore an octet of electrons will have 4 sites / positions around an atom.
• To draw Lewis diagrams

  1. Count the total number of valence electrons in the molecule

  2. Connect outer atoms to the central atom using single bonds (a pair of electrons)

  3. Place the remaining electrons as lone pairs around atoms to complete octets, starting with outer atoms

  4. If you run out of electrons before completing all octets, move lone pairs to make double or triple bonds – thus ensuring each atom gets 8 electrons

  5. Remember H only has 2 electrons and never has lone pairs

  
• Boron and Beryllium molecules


  • When boron is the central atom there will always be an incomplete octet with only 6 electrons around B
  • When beryllium is the central atom there will always be an incomplete octet with only 4 electrons around B

  
• Halogens (group 17) atoms never form double bonds. Do NOT move lone pairs to make double bonds with F, Cl, Br etc.

  
• H2O


  • 2 x 1 + 6 = 8 electrons
  • H – O – H
  • H – O – H
  • O is the central atom
  • Each single bond represents 2 electrons. 4 more electrons need to be added as lone pairs
  • O has 8 electrons – 2 lone pairs and 2 bonding pairs

  
• NCl3


  • 5 + 3 x 7 = 26 electrons
  • Cl – N – Cl
  • Cl
  • Cl – N – Cl
  • Cl
  • N is the central atom
  • Each single bond represents 2 electrons. 20 more electrons need to be added as lone pairs
  • N has 8 electrons – 1 lone pair and 3 bonding pairs
  • Each Cl has 8 electrons – 3 lone pairs and 1 bonding pair

  
• Lewis diagrams


  Used to determine the shape of molecules
• Regions of negative charge around an atom

  • Can correspond to a lone pair of electrons or bonded electrons (either single, double or triple covalent bonds)
  • Repel with maximum separation around an atom to minimise repulsions (valence shell electron pair repulsion theory - VSEPR)
• Molecular shape


  Determined by the distribution of regions of negative charge around the central atom in the molecule (due to repulsion) and how many are bonding / lone pairs

  
• Drawing the shape of a molecule

  Use a 3-D perspective with correct representation of bond angles
• CHCl3


  • Tetrahedral shape
• HOCl


  • Bent shape
• HF


  • Trigonal pyramid shape
• PH3


  • Trigonal pyramid shape
• BF3


  • Trigonal planar shape
• O3


  • Bent shape
• Lewis diagram


  Used to determine molecule shape
• Bonding regions and lone pairs

  • Determines the shape of the molecule
• Molecular shape


  The 3D arrangement of atoms in a molecule
• Predicting molecular shape


  1. Use Lewis diagram to determine molecule shape

  2. Describe geometry of electron regions - number of and repulsions between regions of negative charge

  3. State number of bonding regions and lone pairs

  4. Name the shape of the molecule
• Number of regions of negative charge around central atom


  Determines geometry - repelled with maximum separation to minimise repulsions
• SO3 has 3 regions of negative charge around central atom (S)
• SO3 has trigonal planar geometry with 120° bond angles
• All 3 regions in SO3 are bonding
• The shape of SO3 is trigonal planar with 120° bond angles
• CO2 has 2 bonding regions of electrons
• H2O has 2 bonding regions and 2 lone pairs of electrons
• Electronegativity


  The ability of an atom to attract bonding electrons to itself
• The higher the electronegativity of an atom the more strongly it attracts bonding electrons
• Electronegativity values


  • H (2.1)
  • Li (1.0)
  • Be (1.5)
  • B (2.0)
  • C (2.5)
  • N (3.0)
  • O (3.5)
  • F (4.0)
  • Na (0.9)
  • Mg (1.2)
  • Al (1.5)
  • Si (1.8)
  • P (2.1)
  • S (2.5)
  • Cl (3.0)
  • K (0.8)
  • Ca (1.0)
• Remember the order: F > O > N/Cl > S/C > H
• Polar covalent bond


  A covalent bond where the two atoms have different electronegativities, resulting in one end being slightly positive and the other slightly negative
• A covalent bond becomes more polar when the difference between the electronegativities of the two atoms increases
• Polar covalent bond


  • H-F
• Non-polar covalent bond


  • F-F
• Molecular polarity


  Determined by the symmetry of polar covalent bonds around the central atom