Question

For the following reaction, identify the element that was oxidized, the element that was reduced, and the oxidizing agent. Give an explanation for each answer.

Zn + H2SO4 ---> ZnSO4 + H2

PLEASE HELP! I don't just want the answer, I need to understand this, please walk me through step by step!

Answer 1

Ok so first thing to recognize is that Elements in their free uncombined state have an oxidization number of 0.(This is also true for elements in their natural state that are diatomic O2, H2, etc).

Answer 2

Focus on the Zn on the left hand(reactant) side of the reaction. Given the information above what would its oxidization state be? (notice that it is uncombined and by itself)
Notice the Zn on the right hand(product) side of the reaction it is combined with the negatively charged polyatomic ion SO4^2-. What must the charge of the Zn must be?
Focus on the H2 on the left hand (reactant) side of the reaction. It is combined with the negatively charge polyatomic ion SO4^2-. Think about this what must the charges of the Hydrogen atoms be?
Now look at the H2 on the right hand (product) side of the reaction. It is by itself and in its natural diatomic state. Given the information in the previous response what must the charge of H2 be?

Answer 3

Also note that the Element that loses electrons is being oxidized and the element that gains electrons is being reduced.