As Chemistry Module 3

As Chemistry Module 3

1. Periodicity

   • Across a period
     • Across period 2, 2 electrons fill 2S sub shell and 6 fill 2P sub shell. Across period 3, 2 electrons fill 3S sub shell and 6 fill 3P sub shell.
     • Nuclear charge increases
     • Similar shielding
     • Nuclear attraction increases
     • Atomic radius decreases
     • First ionisation energy increases
     • Fall in ionisation energy from beryllium to boron
       • Marks filling of 2P sub shell
         • 2P sub shell in boron has higher than 2S in beryllium
           • So 2P is easier to remove than 2S.
             • So boron has lower first ionisation energy
   • The repeating trend in properties.
   • Down a group
     • Atomic radius increases
     • Shielding increases
     • Nuclear attraction to outer electrons decreases
     • First ionisation energy decreases
   • Melting points.
     • Across period 2 and 3
       • M.P increases from G1 to G4.
       • Sharp decrease in M.P between G4 and G5
         • This marks a change from giant to simple molecular structures
       • M.Ps are low from G5 to G0

2. Metallic bonding

   • When solid, each atom has lost its outer shell electrons to shared pool of electrons that spread through the whole structure.
     • Delocalised electrons are mobile throughout the structure
   • Cations left behind are made of nucleus and inner electrons shells of the atoms
     • In fixed position to maintain structure and shape
   • The strongest electrostatic attraction between Cations and delocalised electrons
   • Held together in a giant metallic lattice

3. Metal properties

   • Most have strong metallic bonds
   • High electrical conductivity
     • Conduct in solid and liquid state
       • In metal, delocalised electrons move through structure and carry electrical charge
   • High melting and boiling points
     • Depends on strength of metallic bonds
     • High temperatures are necessary for the large amounts of energy needed to break strong electrostatic forces of attraction.
   • Do not dissolve
     • Any interaction would result in a reaction

4. Giant covalent structures

   • Prior knowledge
     • A lot of non metals are simple covalently bonded molecules
       • When solid, these form a simple molecular lattice held together by WEAK intermolecular forces
         • So low M.P and low B.P
   • Boron, carbon and silicon
     • Atoms held together by STRONG covalent bonds
       • Forms a giant covalent lattice
       • Makes very stable structures that are hard to break down
     • Carbon (diamond form) and silicon (group 14) utilise their four unpaired electrons to form covalent bonds to other carbon and silicon atoms
       • Makes tetrahedral structure
         • 109.5 ° bond angles.
       • All 4 outer shells are used in covalent bonding, so non can be used for conducting.
   • High melting and boiling points
   • Insoluble
   • Non-conductors (graphite is an exception).

5. Redox reaction and reactivity

   • Elements in group 2 are reactive metals
     • Redox are most common reactions
     • Reactivity increases as we move down the group
       • Ionisation energies decrease down group as atomic radius increases and shielding increases
   • Reducing agent
     • Each group 2 element has 2 more outer electrons than a noble gas, which are in the S sub shell.
     • Each metal is oxidised as it loses two electrons to become a 2+ ion.
       • Another species gain these electrons and is reduced
         • So group 2 elements are reducing agents
   • Redox with oxygen
     • Forms metal oxide
       • General formula MO
       • 2MgO
         • Mg is oxidised and O₂is reduced
   • With water
     • Alkaline hydroxide
       • M(OH)₂ + H₂
   • With dilute acids
     • Forms a salt and hydrogen gas
     • Metal + acid ~> salt + hydrogen
   • Group 2 oxides
     • Reactions with water
       • Group 2 oxides react with water to release hydroxide ions and form alkaline solutions of the metal hydroxide
         • Group 2 hydroxides
           • Only slightly soluble in water
             • When solution is saturated, any further metal and OH⁻ ions form a solid precipitate
             • Solubility of hydroxides in water increases down the group, making the resulting solutions more alkaline as there will be more OH⁻ ions
               • Test for this trend:
                 • 1) Add a spatula of each group 2 oxide to water
                 • 2) shake. There will be too little water on this scale to dissolve all metal hydroxide that forms, so we get a saturated solution of each
                 • 3) Measure PH of each solution
                   • Alkalinity decreases down the group
     • Uses as bases
       • Calcium hydroxide sis added to fields to increase PH of acidic soils
       • Antacids for treating acid indigestion

6. Halogens

   • Physical properties
     • Most reactive non metallic group (7)
     • Don't occur as their elemental form in nature
       • Occur as stable halide ions dissolved in sea water or combined with sodium or potassium as solid deposits
     • At RTP halogens exist as diatomic molecules.
     • Changes from gas to liquid to solid down the group
       • In solid state, they male lattices with simple molecular structures