1.5 Solid Structures

• Metallic bonding - a regular giant lattice of positive metal ions surrounded by a "sea" of delocalised charge. The attraction between the ions and electrons holds the lattice together (electrostatic forces).
• Crystal coordination numbers: the number of anions around each cation in an ionic lattice and vice versa
• Nanoparticle: particles having one or more dimensions of about 100 nanometres (nm) or less.
• Nanometre = 1 x $10^{-9}$
• Monolayer - a layer of material that is only one atom or molecule thick (about $10^{-9}$m)
• Solubility: the amount of a substance (mass or moles) dissolved in a solvent - concentrations units - gdm^-3 or moldm^-3
• Saturated solution - a solution that has the maximum possible concentration of a solute at given conditions.
• Allotropes: different forms in which a chemical element exists e.g. carbon as graphite, diamond, and coal.
• Crystal: atoms or molecules joined together in a repeating pattern to create a certain shape.
• Sublime: move from solid to gas state without becoming a liquid.
• Solids have particles arranged in organised lattices held together by strong attractive forces. They cannot move but vibrate around a point (fixed positions)
• 3 strong forces that keep atoms tightly bound in a solid:

  1. Hydrogen bonds
  2. Van der vaal's forces / permanent dipole forces
  3. Ionic bonds / metallic bonds
• Sodium chloride and ceasium chloride both form giant ionic lattices but have different lattice structures due to relative sizes of the positivr ions
• Ionic radii

  • Cs+ 0.169nm
  • Na+ 0.095nm
  • Cl- 0.181nm

  Caesium obviously can accomodate more chloride atoms around it
• Coordination number of Sodium Chloride - 6:6
• Sodium Chloride

  
• Caesium Chloride coordination number - 8:8
• Caesium Chloride

  
• Ionic lattices are giant structures of regular repeating arrangement of positive and negative ions.
• Properties of ionic solids

  1. Conduct when molten or in solution
  2. Brittle
  3. High melting and boiling points
  4. Soluble
• Simple covalent compounds are usually quite small
• Boiling points and melting points of covalent compounds are low
• Covalent compounds don't conduct electricity or heat because they are non-polar, uncharged molecules
• Most simple covalent compounds are liquids or gases. Although the molecules have strong intramolecular covalent bonds between the atoms, they are held together by weak intermolecular van der Waal's forces which need little energy to break.
• Iodine is a solid at room temperature and has a crystalline structure with regular faces. However at around 30c it changes from a (silver black) solid directly to a (purple) gas, i.e. it sublimes
• Solid iodine crystal:

  
• Diamond: carbon forms 4 covalent bonds with neighbouring carbons. This means it has a tetrahedral shape with the bond angle of 109.5 degrees. Each neighbouring carbon forms 4 bonds and a giant lattice is formed. The structure is extremely regular and the bonds are very strong.

  
• Properties of diamond - because of its strong covalent bonds

  • Extremely hard - used in diamond tipped saws and drills
  • Vibrations pass through the latice so its a good thermal conductor
  • Very high melting point - sublimes at over 3800K
  • Cannot conduct electricity - no free electrons
  • Insoluble in any solvent - bonds are too strong to be broken
• Graphite: carbon only forms 3 covalent bonds with other carbons. This means it has a triginal planar shape around each carbon with the angle between bonds being 120 degrees. The carbon atom forms layers of covalent hexagons, The 4th unbonded electrons form a delocalised cloud between the layers making a weak bond. These carbons are aligned under each other in layers.

  
• The weak forces between the layers make graphite feel slippery, used as a dry lubricant and in pencils where the layers break off and leave marks.
• Layers in graphite are far apart compared to the length of the covalent bonds, making it less dense than diamond and strong and lightweight.
• Graphite has a very high melting point and sublimes at over 3900K.
• Graphite can conduct electricity because the 4th electron forms a delocalised cloud which can move between layers and carry current.
• Graphite is insoluble in any solvent because covalent bonds are too strong to be broken.
• Diamond cannot be scratched by a piece of graphite because it would suggest that graphite is harder than diamonds.
• Properties of metals

  • High melting and boiling points
  • Conduct electricity (because of delocalised electrons)
  • Ductile
  • High tensile strength
  • Malleable
  • Sonorous
  • Some metals are magnetic
  • Lusterous (because of delocalised electrons)
  • Most have high density
• As the number of delocalised electron per atom increases, the bonding becomes stronger, so the melting and boiling points increase
• If there are impurites in metals, they become more brittle
• As all the metal atoms are identical and form giant lattices with regular structures - layers of atoms. These regular layers can easily slide over eachother, the electrons act as a lubricant, and so metals can be bent, hammered and pulled into wires. Malleable and ductile.
• Because the electron cloud is mobile as a metal is puled or beaten into another shape, it is able to still bind the "ions" together.