Bonding

• Bonding in a metal


  Electrostatic attraction between positive ions and delocalised electrons which extends throughout the giant metallic lattice
• Properties of metals


  • High melting/ boiling points
  • Strong
  • Ductile: can be drawn into wires
  • Malleable: can be beaten into sheets
  • Insoluble
  • Good conductors of electricity
  • Good conductors of heat
• Why do metals have high melting points?


  Due to strong metallic bonds which extend throughout the giant metallic lattice
• Factors affecting the melting point of a metal


  • The charge of the ion/ number of delocalised electrons (larger = stronger)
  • The ionic radius (smaller = stronger)
  • Therefore, the electrostatic attraction between the cation and delocalised electrons is stronger/ weaker
• Why are metals malleable and ductile?


  Layers of ions can slide over each other
• Why are metals good conductors of electricity and heat?

  Electrons are delocalised
• Why are metals insoluble?

  The metallic bonds are too strong to be overcome by interactions with water molecules
• Ionic bonding


  Strong electrostatic forces between positive and negative ions (oppositely charged ions) which extend throughout the giant ionic lattice
• Ionic lattice


  Repeating pattern of oppositely charged ions
• Arrangement of particles in an ionic crystal


  1. Minimum 4 ions shown in 2D square arrangement placed correctly

  2. Further 3 ions shown correctly in a cubic lattice
• Ionic formulas


  • SO4^2-
  • OH^-
  • NO3^-
  • CO3^2-
  • HCO3^-
  • NH4^+
  • CN^-
  • PO4^3-
• Why do ionic compounds have a high melting point?

  Strong attraction between oppositely charged ions requires lots of energy to overcome
• Properties of ionic compounds


  • High melting point (solids at room temperature)
  • Usually soluble in polar solvents
  • Conduct electricity when molten or aqueous
  • Brittle
• Why can't solid ionic compounds conduct electricity?


  Ions can't move in solid salt
• Why are ionic compounds brittle?

  A blow can force two ions with like charges to come into contact and repel each other
• Factors affecting the melting point of an ionic compound


  • The charge of the ion (larger = stronger)
  • The ionic radius (smaller = stronger)
  • Therefore, the electrostatic attraction between the cations and anions is stronger/ weaker
• Covalent bond


  A shared pair of electrons
• When is a bond covalent rather than ionic?


  The elements involved have similar electronegativities
• Dative/ co-ordinate bond


  A shared pair of electrons where both electrons have been supplied by one atom
• Shape and bond angle(s) of molecules

  • Linear, 180°
  • Trigonal planar, 120°
  • Bent, 118°
  • Tetrahedral, 109.5°
  • Trigonal pyramid, 107°
  • Bent, 104.5°
  • Trigonal bipyramid, 90° and 120°
  • Octahedral, 90°
• Why is methane a tetrahedral molecule?

  There are 4 bonding pairs and 0 lone pairs surrounding the central atom. All bonding pairs repel equally
• Why does ammonia have bond angles of 107 rather than 109.5°?


  There are 3 bonding pairs and 1 lone pair surrounding the central atom. Lone pairs repel more than bonding pairs which reduces the bond angle
• Why is it difficult to predict the shape of transition metal complexes?


  Many orbitals in the d subshell
• Electronegativity


  The ability of an atom to attract a pair of electrons in a covalent bond towards itself
• Why are noble gases not included in trends of electronegativities?


  They don't form covalent bonds
• Trend in electronegativities across a period


  • General increase due to: increased nuclear charge, same shielding, smaller atomic radius, therefore stronger attraction between nucleus and 2 electrons in covalent bond
• Trend in electronegativities down a group


  • General decrease due to: increased shielding, larger atomic radius, therefore weaker attraction between nucleus and 2 electrons in covalent bond - these factors outweigh the increased nuclear charge
• Polar bond


  A bond where the electrons are not shared equally due to a difference in electronegativity of the atoms involved. The more electronegative atom is δ- and the other is δ+
• Intermolecular forces from weakest to strongest


  • Van der Waals forces
  • Dipole-dipole forces
  • Hydrogen bonds
• How van der Waals forces arise


  Electron movement causes a temporary dipole in one molecule which induces an instantaneous dipole in neighbouring molecules. There is induced-temporary attraction between the δ- and δ+ ends of adjacent molecules
• Why are van der Waals' forces in noble gases weak?


  Noble gases are single atoms with electrons closer to nucleus (atomic radius decreases across a period), cannot easily be polarised
• How dipole-dipole forces arise


  Molecules with a permanent dipole moment participate in dipole-dipole forces as the δ- end of one molecule is attracted to the δ+ end of adjacent molecules
• Why do symmetrical molecules never participate in dipole-dipole forces?


  Dipoles cancel in a symmetrical molecule
• How hydrogen bonds arise


  Hydrogen bonds arise when a δ+ hydrogen is sandwiched between 2 very electronegative elements: N/O/F. N/O/F are much more electronegative than hydrogen which sets up a H-N/O/F dipole. A hydrogen bond forms between H and lone pair on N/O/F on an adjacent molecule
• Physical properties affected by intermolecular forces


  • Melting and boiling points
  • Volatility
  • Viscosity
  • Surface tension
  • Solubility
• Unusual properties of water due to hydrogen bonding


  • Relatively high melting and boiling points
  • Ice is less dense than liquid water
  • High surface tension
• Types of substances that can dissolve in water


  • Most ionic compounds (unless their lattice enthalpy is too high/ they have significant covalent character)
  • Some polar molecules (short chained)
• Non-polar molecules dissolve in what kind of solvent?

  Non-polar solvents such as hexane/ cyclohexane
• Types of crystal structure


  • Metallic
  • Ionic
  • Macromolecular
  • Molecular
• Allotropes of carbon with macromolecular structures


  • Diamond
  • Graphite
  • Graphene
  • Buckminster fullerenes