C4 - The Periodic Table

C4 - The Periodic Table

1. Atomic Struture

   • Protons: positively charged, charge +1, mass 1
   • Neutrons: no relative charge, charge 0, mass 1.
   • Electrons are negatively charged, charge -1, mass 0.0005
   • Atomic Number = number of protons element contains found at the bottom of element on P.T
   • Mass Number = total number of protons and neutrons element contains
   • Isotopes = same element, same atomic number, different mass number.

2. Electronic Structure

   • Compound = two or more elements, chemically combined.
   • Chemical formula = symbols of all elements in a compound and number of atoms of each element.
   • Electrons occupy shells around the nucleus
     • first shell can home two electrons.
     • second and third shells can home up to eight electrons

3. Ionic Bonding

   • If an atom loses one or more electrons it becomes a positively charged ion
     • Hydrogen and metal atoms
   • If an atom gains one or more electrons it becomes a negatively charged ion
     • Non-metal atoms
   • Metal reacts with non-metal
     • Electrons transferred from metal to non-metal
       • Positive metal ions attract negative non-metal ions and form ionic bonds.
         • Dot and cross diagrams

4. Ionic Compounds

   • Giant ionic lattice
     • Structure repeated many times
     • Positive ions strongly attracted to negative ions.
   • Ionic bonds need a lot of energy to be broken
     • High Melting Point.
     • Can conduct electricity in a molten or liquid solution - ions are free to move around

5. Periodic Table

   • horizontal rows = period
   • vertical column = group
     • Elements in a group have similar properties
   • number of electron shells occupied = period element is found
     • Number of electrons in outer shell = group number

6. Covalent Bonding

   • two non-metals
   • Shared pair of electrons
   • Each covalent bond is shown as a straight line in displayed formula
   • Weak intermolecular forces
   • Low melting point
   • Does not conduct electricity
   • Shared electrons shown on diagram in shared area between atoms

7. Group One Elements

   • React vigorously with water to form alkaline solutions
     • Stored in oil to keep them away from water and air
   • Metal + Water = Metal Hydroxide + Hydrogen
     • 2M + 2H20 = 2MOH + H2
   • reactivity increases as you go down the group.
   • In water, lithium keeps its shape, potassium and sodium melt and form silver balls
     • Potassium ignites explosively and burns with a lilac flame
       • Lithium disappears slowly, sodium disappears quickly and potassium disappears very quickly
   • Each have one electron in outer shell
   • Loss of electron = oxidation
     • M = (M+) + (e-)
   • Densities increase as you go down the group
   • Melting points increase as you go down the group

8. Flame Tests

   • Lithium burns red
   • Sodium burns orange
     • Potassium burns lilac

9. Group Seven Elements

   • Halogens - react with metal to make salts
     • Chlorine sterilises pools and tap water
     • Iodine is used as antiseptic on wounds
   • Chlorine is a green gas
     • Group Seven elements react vigorously with Group One
       • Metal Halide clouds produced
       • 2M + X2 = 2MX (where M = group one and X = group 7)
   • Bromine is an orange liquid
   • Iodine is a grey solid

10. Displacement reactions

    • A more reactive group 7 will displace a less reactive group 7 from its metal hallide.
      • X2 + 2MY = Y2 + 2MX where X = group 7 and MY = metal hallide
        • Gain of electron = reduction
          • X2 + (2e-) = 2X-)

11. Transition Elements

    • between groups 2 and 3
      • They are all metals
        • Strong, malleable and shiny when cut
    • Copper compounds are often blue
    • Iron(II) compounds are often light green
      • Iron(III) compounds are often orangey brown
        • Often used as catalysts, EG: iron is used in the haber process.
        • Thermal Decomposition = substance breaks down to form two or more substances when heated
          • Metal carbonate = metal oxide + carbon dioxide
            • Normally causes a colour change

12. Precipitation reactions

    • Copper(II) - blue metal hydroxide precipitate
      • Iron (II) - grey/green metal hydroxide precipitate
        • Iron (III) - orange/brown metal hydroxide precipitate

13. Metal

    • Good conductor of electricity
      • hard, lustrous, malleable and have tensile strength
    • Metallic bonding - strong - high melting and boiling points
      • Closely packed ions
      • sea of delocalised electrons
        • Takes a lot of energy to overcome bonds

14. Water Purification

    • Water comes from underground aquifers, resevoirs, lakes and rivers.
      • contains many dissolved substances which need to be removed
    • Water can be used in industry as a solvent, a coolant or a cheap raw material
      • Sulfate ions can be detected by barium chloride solution
        • Sample of water in test tube, barium chloride added
          • Sulfate Ions present = white precipitate
            • Na2SO4 + BaCl2 = 2NaCl + BaSO4
      • Hallide ions can be detected by silver nitrate
        • Chloride Ions = White precipitate
        • bromide ions = cream precipitate
          • NaCl + AgNO3 = NaNO3 + AgCl
        • iodide ions = pale yellow precipitate
    • Only 3% of the Earth's water is fresh
      • Pollutants in water = nitrates, lead compounds and pesticides
        • Distillation can make pure water from sea water but requires lots of energy and is more expensive
      • Large suspended particles removed by sedimentation
        • Small suspended particles removed by filtration
          • chlorination kills microbes in the water