Can two elements both be oxidized in a reaction?
2FeS + 4.5O2 +2H2O==> Fe2O3 + 2H2SO4
I see Fe is oxidized, going from Fe+2 to Fe+3
However, Isn't S going from S-2 to S+6? making sulfur ixidized too, and not reduced?
Yes. I confirmed all of that in my post BOTH times.
Yes, in this reaction, both iron (Fe) and sulfur (S) undergo oxidation. Iron is oxidized from a +2 oxidation state to a +3 oxidation state, indicating a loss of electrons. On the other hand, sulfur is oxidized from a -2 oxidation state in FeS to a +6 oxidation state in H2SO4, indicating a gain of electrons.
To determine the oxidation state of an element, you can follow these steps:
1. Determine the common oxidation states: It helps to know the common oxidation states for each element involved in the reaction. For example, iron (Fe) often has oxidation states of +2 or +3, while sulfur (S) commonly has oxidation states of -2, +4, or +6.
2. Assign oxidation states based on known rules: There are certain guidelines you can follow to assign oxidation states. For example, in a neutral compound, the sum of oxidation states should be zero. In an ion, the sum of oxidation states should be equal to the ion's charge. You can also consider the electronegativity differences between elements to assign oxidation states.
3. Check for changes in oxidation states: Compare the oxidation states of each element before and after the reaction. If there is an increase in the oxidation state, the element is oxidized, indicating a loss of electrons. Conversely, if there is a decrease in the oxidation state, the element is reduced, indicating a gain of electrons.
In the given reaction, Fe is oxidized from +2 to +3, meaning it loses an electron. Sulfur is oxidized from -2 to +6, meaning it gains electrons. Therefore, both Fe and S undergo oxidation in this reaction.