Multiple Oxidation States Truman Chemistry Dept.

Oxidation States can change during a reaction Example: Oxidation state of oxygen CH4 + 2O2  CO2 + 2H2O In O2 the oxidation state of O is zero In CO2 and H2O the oxidation state of O is -2 So the main question here is how can we predict how the oxidation state will change? There are rules to predict this….

Rules All elements in a diatomic molecule have an oxidation state of zero (they are neutral) ie. Cl2 If an element is monatomic (by itself) the charge is only as specified (ie. Cl-). Free elements have a charge of zero (ie. Fe) All neutral compounds must have a net (total) oxidation charge of zero and the element with the highest electronegativity (furthest right on table) is the negative element.

Rules continued… Hydrogen always has a +1 charge except in hydrides where it is negative one (ie. LiH or CaH2) Oxygen always has a -2 charge except in peroxides (H2O2) where it is -1 Group 1 elements are normally +1 Group 2 elements are normally +2 Group 17 elements are often -1

Polyatomic ion rules… Polyatomic ions have a net charge that is equal to the sum of all the charges.

Lets try some: 1. The oxidation number of nitrogen in N2 is (1) +1    (2) 0    (3) +3      (4) -3 What is the oxidation number of hydrogen in CaH2?  (1) +1     (2) +2      (3) -1     (4) -2 What is the oxidation number of carbon in NaHCO3?  (1) -2     (2) +2      (3) -4   (4) +4

A few more? What is the oxidation number of chlorine in HClO4? (1) +1      (2) +5    (3) +3     (4) +7 What is the oxidation number of sulfur in H2SO4? (1) 0     (2) -2     (3) +6   (4) +4 What is the oxidation number of chromium in K2Cr2O7? (1) +12     (2) +2   (3) +3    (4) +6

Lets look at some reactions: Al + O2  Al2O3 What happened to the oxidation states? Half reactions? S + NO3-  SO2 + NO