Balanced and Net ionic equations (1 Viewer)

[.kumori]

how do I solve the balanced and net ionic equation of:
a) iron + Silver (l) Nitrate
b) Tin + Lead (ll) Phosphate
Im kinda confused, bc I asked google already and it said that when iron reacts with nitrate it could have an oxidation of 2+ or 3+, and when tin reacts with phosphate, it could have an oxidation of 2+ or 4+

 

[wizzkids]

Im kinda confused, bc I asked google already and it said that when iron reacts with nitrate it could have an oxidation of 2+ or 3+, and when tin reacts with phosphate, it could have an oxidation of 2+ or 4+

Whoa, there are so many things wrong with your approach, you need to go right back to Chemistry basics.
In redox reaction (a) , iron does not react with nitrate NO₃^-
In redox reaction (b) , tin does not react with phosphate PO₄^3-.
These anions (negatively charged ions) are "spectator ions" ; they don't participate in the redox reactions.
OK so what is the real nature of the chemical change then?
It is electron exchange between one metal that is oxidised (loses electrons) and another metal which is reduced (gains electrons).
The metal that is oxidised is also called a reducing agent (because it reduces its counterpart) and the metal that is reduced is also called an oxidising agent (because it oxidises its counterpart). It is the relative power of the oxidising agent that determines what oxidation state its counterpart will be left in.
The power of an oxidising agent is measured by its Standard Oxidation Potential, E^o, which has the same magnitude as the Standard Reduction Potential (but opposite sign).
So, now can you try the exercise again?