1. Redox Reactions and Oxidation States

2. The Basics...  ‘Redox’ is used to describe reactions where oxidation and reduction take place.  If one reactant is oxidised another must be reduced.  Remember OIL RIG – oxidation is loss, reduction is gain. (Of electrons).  Reducing agents therefore give electrons, reducing other reactants. (Oxidising themselves).  Oxidising agents accept electrons, oxidising other reactants. (Reducing themselves).

3. Other definitions of Oxidation  Oxidation can also mean the gain of oxygen atoms, so oxidising agents give away oxygen.  Reduction can also mean the loss of oxygen atoms, so reducing agents take oxygen from another substance.  Oxidation is the loss of hydrogen.  Reduction is the gain of hydrogen.

4. Summary of Redox... OxidationReduction Loss of electronsGain of electrons Gain of oxygenLoss of oxygen Loss of hydrogenGain of hydrogen

5. Periodicity and Oxidising Power  Oxidising power increases across a period.  Reducing power increases down a group.  Best oxidising agents = top right of table.  Best reducing agents = bottom left of table.

6. Why These Trends?  The top elements in the groups have the least number of shells, so attract electrons more easily – accepting them from other reactants and oxidising them.  The lower elements in the groups have a greater shielding effect and distance from the nucleus acting against them in this respect...it is therefore more difficult for them to attract electrons and easier to give them away, making them better reducing agents.

7. More about Trends  Atomic radii decrease across a period, because due to the added protons there is a greater attraction from the nucleus but no increased shielding effect, so electrons are more easily pulled in.

8. Best oxidising agents Best reducing agents

9. Halogens As Oxidising Agents  The halogens are in group seven – on the right side of the periodic table, so they are good oxidising agents.  They accept electrons from other atoms, becoming halide ions and reducing themselves – oxidising the other atom.  This oxidising strength reduces going down the group - think of the trends.

10. Chlorine Oxidation NumberExample +7ClO 4 -, Cl 2 O 7 +6ClO 3 +5ClO 3 - +4ClO 2 +3ClO 2 - +1ClO - 0Cl 2 Cl - Chlorine has a variety of oxidation states depending on which compound it is in – remember, as an element it will have an oxidation number of zero, just like all elements.

11. Other Halogens  Fluorine's oxidation number is always -1 in compounds.  Bromine has the same range of oxidation numbers as chlorine.  Iodine has the same as bromine and chlorine, except not +4 and +6.

12. Disproportionation Reactions  Cl 2 ClO 3 -  oxidation Cl -  reduction Example Making bleach – reacting chlorine with an alkali at 15 ° C... Cl 2 (aq) + 2OH - (aq)  Cl - (aq) + ClO - (aq) + H 2 O (l)

13. Halogens and H 2 SO 4  Chlorine... KCl + H 2 SO 4  KHSO 4 + HCl (white misty fumes)  Bromine... KBr + H 2 SO 4  KHSO 4 + HBr (white misty fumes, litmus paper red) H 2 SO 4 + 2HBr  SO 2 + Br 2 + 2H 2 O (brown/orange fumes)

14. Halogens and H 2 SO 4 continued  Iodine... KI + H 2 SO 4  KHSO 4 + HI (white misty fumes) H 2 SO 4 + 6HI  S + 3I 2 + 4H 2 O H 2 SO 4 + 8HI  H 2 S + 4I 2 + 4H 2 O (purple vapour and rotten egg smell)

15. Why...?  Sulphuric acid acts as both an acid and an oxidising agent. When acting as an oxidising agent  ○ Not strong enough to oxidise fluoride or chloride ions/fluoride and chloride ions are not strong enough reducing agents to reduce H 2 SO 4. ○ Bromide ions are oxidised to bromine. ○ Iodide ions are stronger reducing agents – they reduce H 2 SO 4 to sulphur dioxide, then to sulphur, then to hydrogen sulphide. When acting as an acid  ○ A hydrogen ion is given to the halide ion, from the acid, forming HBr, HI, HCl.