trends

• The Periodic Table


  • Elements are arranged according to atomic number
  • Vertical columns are called groups
  • All elements in the eight main groups contain the same outer electron configuration
  • Horizontal rows are called periods
  • All elements in a period have the same number of quantum shells containing electrons
  • The table is also divided into blocks
  • The name of the block shows the orbital in which the elements' outer electrons lie
• Ionisation energy


  • Generally increases across a period because there is an increase in nuclear charge in the same energy level
  • Decreases between Group 2 and 3 because Group 3's outer electron is partly shielded by the s electrons
  • Decreases between Groups 5 and 6 because of electron-electron repulsion between the electron pair in one p orbital
  • Decreases down a group because the outer electron has increased shielding from inner electrons
• Electronegativity


  Increases across a period because there is an increase in nuclear charge, but the bonding electrons are always shielded by the same inner electrons
• Melting and boiling temperatures


  • Generally increase from the first to the fourth element, followed by a large decrease to the fifth element and a small general decrease to the eighth element
  • This is because the bonding changes from metallic to giant covalent to simple molecular covalent
• Oxidation


  Loss of electrons
• Reduction


  Gain of electrons
• Oxidising agent


  A species that accepts electrons; it becomes reduced itself in the process
• Reducing agent


  A species that donates electrons; it becomes oxidised itself in the process
• Determining if a reaction is redox

  1. Work out the oxidation numbers of the atoms or ions

  2. If the oxidation number increases, the species is oxidised

  3. If the oxidation number decreases, the species is reduced
• Reaction of Group 1 metals with water

  1. React vigorously with cold water to form the hydroxide and hydrogen

  2. The reaction increases in vigour as you go down the group
• Reaction of Group 2 metals with water

  1. React less vigorously

  2. Magnesium reacts very slowly

  3. Calcium produces a steady stream of bubbles and a white precipitate of calcium hydroxide

  4. The hydroxide and hydrogen are formed

  5. The reaction increases as you go down the group
• Reactivity of s-block metals

  • Increases as you go down a group because when they react they lose electrons to form positive ions
  • Since ionisation energies decrease down a group, the energy needed to form positive ions decreases
• Reactivity of Group 1 vs Group 2 metals

  Group 1 metals are more reactive than Group 2 metals because they lose only one electron while Group 2 metals lose two
• Reaction of s-block metals with oxygen

  All Group 1 and Group 2 metals burn to form solid white oxides
• Reaction of s-block metal oxides with acids

  All s-block metal oxides are strong bases and neutralise acids to form salt and water
• Reaction of Group 1 oxides and barium oxide with water


  React to form a soluble hydroxide or alkali
• Most s-block elements may be identified by a flame test (Mg2+ ions give no colour)
• Characteristic colours for metal ions in flame tests

  • Li+ - red
  • Na+ - orange yellow
  • K+ - lilac
  • Ca2+ - brick red
  • Sr2+ - crimson
  • Ba2+ - apple green
• Trends for Group 2 compounds

  • All nitrates are soluble
  • All carbonates are insoluble
  • Hydroxides are more soluble as you go down the group (Mg(OH)2 insoluble, Ba(OH)2 soluble)
  • Sulfates become less soluble as you go down the group (MgSO4 soluble, BaSO4 insoluble)
• Thermal decomposition of Group 2 hydroxides and carbonates

  1. Hydroxides decompose to the oxide and steam

  2. Carbonates decompose to the oxide and carbon dioxide

  3. Thermal stability increases in both as you go down the group
• Physical properties of halogens

  • At room temperature, chlorine is a green gas, bromine a red-brown liquid and iodine a grey solid
  • As the number of electrons increases with atomic number, there is an increase in the induced dipole – induced dipole intermolecular forces holding the diatomic molecule together
  • Thus, the melting and boiling temperatures increase as you go down the group
• Reactivity and oxidising power of halogens


  Decreases down the group because the outer electrons are shielded more and are further from the nucleus, so it gets harder to attract electrons
• Reaction of halogens with metals

  Halogens react directly with most metals to form the halide
• Displacement reactions between halogens

  1. A halogen in a higher position in the group will oxidise a halide ion from lower in the group

  2. This causes colour changes
• Test for halide ions

  1. Add a few drops of nitric acid to the aqueous halide ion, then add silver nitrate solution, followed by dilute aqueous ammonia

  2. Cl- - white precipitate that dissolves

  3. Br- - cream precipitate that dissolves slightly

  4. I- - pale yellow precipitate with no change
• Chlorine gas is added to drinking water to kill dangerous bacteria and viruses such as cholera and typhoid
• Fluoride is added to water to reduce tooth decay and strengthen bones
• Formation of insoluble salts


  1. Two suitable solutions are mixed to form a soluble salt and an insoluble salt

  2. The precipitate is filtered, washed and dried
• Formation of soluble salts

  1. By neutralising an acid

  2. Any excess solid is filtered

  3. The solution is evaporated and left to cool to form crystals