Lecture Notes from CHM 1341
8 July 1996

Oxidation Numbers

The following is Table 3.5 from Atkins & Beran

TABLE 3.5 Rules for assigning oxidation numbers

Work through the rules in the order given.
Stop as soon as the oxidation number has been assigned.

                                       Oxidation Number
1. The sum of the oxidation numbers
   of all the atoms in the species
   is equal to its total charge.

2. For atoms in their elemental form   0

3. For elements of Group I            +1
                   Group II           +2

       Group III (except B)           +3 for M3+

                                      +1 for M+

    Group IV (except C, Si)           +4 for M4+

                                      +2 for M2+

4. For hydrogen                       +1 in combination with nonmetals
                                      -1 in combination with metals

5. For fluorine                       -1 in all its compounds

6. For oxygen                         -2 unless combined with F

                                      -1 in peroxides (O22-)

                                    -1/2 in superoxides (O2-)

                                    -1/3 in ozonides (O3-)

---

Chlorine

Chlorine comes in an even broader range of oxidation states than does Sulfur. It can be as low as -1 (as in chloride ion) to +7 (as in perchlorate ion):

Ox. #  -I          0         +I          +III         +V          +VII
         _                      _           _            _           _
E.g.:  Cl         Cl2        ClO         ClO2         ClO3        ClO4

     chloride  chlorine  hypochlorite  chlorite     chlorate   perchlorate
                         
Acid: HCl                   HClO        HClO2        HClO3       HClO4

     hydrochloric        hypochlorous   chlorous    chloric    perchloric

That's a lot of nomenclature to remember! But there's a mnemonic (memory trick) to it. Notice that the suffix changes between III and V? And that the state below III has III's name with the prefix hypo? Well, "hypo" (actually "hupo") means "beneath" in Greek, referring here to the +I state being below the +III state in an ascending order. Likewise, state VII is the same name as V but with the prefix per which means "thoroughly" or "completely." In this case, it means "thoroughly oxidized!"

And IUPAC (International Union of Pure and Applied Chemistry, the same enlightened group which wants to take Professor Seaborg's name off Sg, Seaborgium, an element he and his group at Berkeley discovered) is working on something useful; rationalizing the names of oxoacids and oxoanions so hypochlorous acid would become chloric(I) acid etc. They have already succeeded in rationalizing transition metal oxidation state names so that Hg₂Cl₂ is no longer mercurous chloride but mercury(I) chloride and HgCl₂ is no longer mercuric chloride but mercury(II) chloride so if they succeed, Hg(ClO₄)₂ would be mercury(II) chlorate(VII), much easier!

Don't hold your breath.

While sulfur has a neat redox trick which in one reaction destroys two pollutants by reacting it's -I (sulfide) state with its +IV (sulfur dioxide) state to produce the relatively benign elemental sulfur:

     [-I]    [IV]              [0]
   4 H2S(g) + SO2(g)  ----->  5 S(s) + 2 H2O(l)

in a reaction known as comproportionation (product of an oxidation state intermediate between reactants'), chlorine does the opposite trick!

Instead of two different oxidation states of the atom redoxing one another, chlorine gas (aqueous) redoxes itself in a process known as disproportionation (reactant at an oxidation state intermediate between those of the products):

  [0]                     [-I]_        +        [+I]
  Cl2(aq) + H2O(l)  ----->  Cl (aq) + H (aq) + HOCl(aq)

Since chlorine gas is stable, it cannot pull this trick without the wonderful ionizing power of aqueous solution. What's even more interesting is that sulfur cannot pull off its comproportionation trick at any reasonable rate unless catalyzed by a bit of water! (Remind me to show the class a video of that reaction; it's impressive.)

Note the simplicity of the stoichiometric coefficients in the chlorine case as opposed to the sulfur one! That's because chlorine is both gaining and losing only one electron, but sulfur is losing 1 but gaining 4 (the latter in dropping from [IV] to [0]); so 4 hydrogen sulfides must oxidize to produce the 4 electrons sulfur dioxide needs to reduce. Writing out the half reactions would've made this plain.

As with all the oxoacids, the more highly oxidized versions are much more acidic than the less...because the extra electronegative drain of more oxygens makes it easier to ionize the acidic hydrogens. But also the fact that chlorine is more electronegative than sulfur means that its most acidic (perchlorate) acid is stronger than sulfur's (sulfuric acid). We would expect, therefore, that the acids made from phosphorous, left of both of these atoms on the periodic table, would be weaker than either, and that is most definitely true.

Chris Parr
University of Texas at Dallas
Programs in Chemistry, Room BE3.506
P.O. Box 830688  M/S BE2.6 (for snailmail)
Richardson, TX  75083-0688

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