Bonding and structure

• Ionic bonding


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


  1. Metal atoms lose electrons to form +ve ions

  2. Non-metal atoms gain electrons to form -ve ions
• Ionic compounds


  • Mg2+ 1s2 2s2 2p6
  • 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-
• Giant ionic lattice


  Regular 3D pattern of ions in an ionic solid
• Ionic compounds


  • High melting points due to strong electrostatic attractive forces
  • Non-conductors of electricity when solid as ions are held tightly in lattice
  • Good conductors of electricity when in solution or molten as ions are free to move
• Covalent bond


  Strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
• Dative covalent bond


  Shared pair of electrons comes from only one of the bonding atoms
• Compounds with dative covalent bonds


  • NH4+
  • H3O+
  • NH3BF3
• Average bond enthalpy


  Measurement of covalent bond strength, larger value means stronger bond
• Types of bonding


  • Ionic
  • Covalent
• Structures


  • Giant ionic lattice
  • Simple molecular
• Only use the words molecules and intermolecular forces when talking about simple molecular substances
• Properties of ionic vs molecular compounds


  • Ionic: High melting/boiling points, generally good solubility in water, poor conductivity when solid, good conductivity when molten
  • Molecular: Low melting/boiling points, generally poor solubility in water, poor conductivity
• Molecular shapes


  • Linear: 2 bonding pairs and 0 lone pairs. BOND ANGLE: 180. Example: CO2, BeCl2
  • Trigonal planar: 3 bonding pairs and 0 lone pairs. BOND ANGLE: 120. Example: BCl3, BF3
  • Tetrahedral: 4 bonding pairs and 0 lone pairs . BOND ANGLE: 109.5. Example: CH4, SiCl4
  • Trigonal pyramidal: 3 bonding and 1 lone pair. BOND ANGLE: 107. Example: NH3, PF3, ClO3
  • Bent: 2 bonding pairs and 2 lone pairs. BOND ANGLE: 104.5. Example: H2O, OCl2,
  • Octahedral: 6 bonding pairs and 0 lone pairs. BOND ANGLE: 90. Example: SF6
• Electronegativity


  Relative tendency of an atom in a covalent bond to attract electrons to itself
• F, O, N and Cl are the most electronegative atoms
• Polar covalent bond


  Bond with unequal distribution of electrons, producing a charge separation (dipole)
• The element with the larger electronegativity in a polar compound will be the δ- end
• Symmetric molecules with all bonds identical and no lone pairs will not be polar even if individual bonds are polar
• Induced dipole-dipole interactions


  Temporary dipoles in molecules can induce dipoles in neighbouring molecules
• Factors affecting induced dipole-dipole interactions


  • Larger molecules with more electrons have stronger interactions
  • Long chain alkanes have larger surface area for interactions
• Permanent dipole-dipole forces


  Stronger than induced dipole-dipole, occur between polar molecules
• Hydrogen bonding


  Strong intermolecular force occurring between H atom attached to N, O or F which has a lone pair
• Hydrogen bonding occurs in addition to induced dipole-dipole interactions
• Compounds that can form hydrogen bonds


  • Water
  • Alcohols
  • Carboxylic acids
  • Proteins
  • Amides
• Water can form two hydrogen bonds per molecule due to its high electronegativity and two lone pairs
• Molecular structures


  • Iodine - covalent bonds between I2 molecules, held together by weak induced dipole-dipole interactions
  • Ice - water molecules held further apart than in liquid water
• Ionic compounds are oppositely charged ions held by strong electrostatic attractions
• Ionic compounds


  Oppositely charged ions held by strong electrostatic attractions
• Determining ionic compound formula

  1. Write ions

  2. Swap charges

  3. Drop charges

  4. Simplify ratios
• Ionic compounds


  • Have giant ionic structures
  • Dissolve in water, conduct electricity when molten or dissolved
  • Have high melting points due to strong electrostatic forces
• Covalent bonding


  Sharing of electrons to achieve full outer shells
• Single, double, triple bonds


  Different numbers of shared electron pairs
• Dative covalent (coordinate) bond


  One atom donates both bonding electrons to another atom
• Molecular shapes


  • Determined by number of bond pairs and lone pairs
  • Bond angles affected by lone pairs repelling more than bond pairs
• Molecular shapes with no lone pairs


  • Linear
  • Trigonal planar
  • Tetrahedral
  • Trigonal bipyramidal
  • Octahedral
• Molecular shapes with lone pairs


  • Trigonal pyramidal
  • Bent
  • Trigonal planar (no effect)
  • Square planar (no effect)
• Electronegativity


  Ability of an atom to attract electrons in a covalent bond
• Electronegativity increases up and to the right in the periodic table (excluding noble gases)