Redox (I)

Redox means reduction and oxidation

Oxidation
Loss of electrons
Reduction
Gain of electrons
Reducing agent
a substance that gives electrons to anothe substance (itself becoming oxidised)
Oxidising agent
a substance that removes electrons from another substance (itself becoming reduced)

Fixed oxidation states of ions

Some elements have (mostly) fixed oxidation states, shown in the table below:

Element Oxidation number
Li, Na, K +1
Mg, Ca, Sr, Ba +2
Al +3
N (expect oxides) -3, +3, +5
O (except peroxides, superoxides and compounds with F -2
Any element 0
S -2, +1, +2, +2.5, +4, +6
F (always) -1
Cl, Br, I -1, +1, +3, +5, +7

Things to note

  • In simple ions, the oxidisation state is the same as the charge on the ion.
  • In compounds the total of all of the oxidation states is zero
  • In polyatomic ions (such as Sufate) the total of all oxidation states is equal to the charge on the ion

Common oxidising and reducing agents

Oxidising agent Product when reduced Reducing reagents Product when oxidised
Chlorine, $Cl_2$ Chloride ions, $Cl^-$ Iodide ions ($I^-$) or hydrogen iodide ($HI$) Iodine, $I^-$
Bromine, $Br_2$ Bromide ions, $Br^-$ Hydrogen sulfate, $H_2S$ Sulfur, $S$
*Manganate ($VII$) ions, $MnO4^-$ Manganese($II$) ions, $Mn^2+$ Sulfur dioxide, $SO_2$ Sulfate ions, $SO_4^{2-}$
*Dichromate ($VI$) ions, $Cr_2O_7^{2+}$ Chromium ($III$) ions, $Cr^{3+}$ Iron($II$) ions, $Fe^{2+}$ Iron ($III$) ions, $Fe^{3+}$
*Hydrogen peroxide, $H_2O_2$ Water, $H_2O$ Hydrogen peroxide, $H_2O_2$ Oxygen, $O_2$
Iron ($III$), $Fe^{3+}$ Iron ($II$) ions, $Fe^{2+}$ Tin($II$) ions, $Sn^{2+}$ Tin ($IV$) ions $Sn^{4+}$
Conc. Sulfuric acid, $H_2SO_4$ Sulfur dioxide, $SO_2$ Carbon Carbon monoxide/dioxide
Conc. Nitric acid, $HNO_3$ Nitrogen dioxide, $NO_2$ Carbon monoxide Carbon dioxide
Hydrogen ions, $H^+$ Hydrogen, $H_2$ Metals ($Mg$ etc) Metal ions, $Mg^{2+}$

*solution must be acidified