To identify the species that is reduced in the reaction \( \text{Fe} + \text{AgNO}_3 \rightarrow \text{Fe(NO}_3\text{)}_3 + \text{Ag} \), we can follow these steps:
1. Assign oxidation numbers to each element in the reaction:
- Fe (elemental iron):
Oxidation number of Fe in its elemental form is 0.
- AgNO\(_3\) (silver nitrate):
- In AgNO\(_3\), nitrate (NO\(_3^-\)) is a polyatomic ion with an overall charge of -1.
- Oxygen generally has an oxidation number of -2. Since there are 3 oxygen atoms in NO\(_3^-\), the total oxidation number for oxygen is -2 x 3 = -6.
- Nitrogen must balance the -6 from oxygen to give the overall charge of -1 for nitrate, so nitrogen has an oxidation number of +5.
- Silver (Ag) in AgNO\(_3\) has an oxidation number of +1, to balance the -1 charge from nitrate.
- Fe(NO\(_3\))\(_3\) (iron(III) nitrate):
- Here, nitrate remains as NO\(_3^-\) with an overall charge of -1.
- Since there are three nitrate ions, their combined charge is -1 x 3 = -3.
- To balance these charges, the oxidation number of Fe must be +3.
- Ag (elemental silver):
Oxidation number of Ag in its elemental form is 0.
2. Determine the changes in oxidation numbers:
- For Fe:
- In the reactants, it has an oxidation number of 0.
- In the products, it changes to +3 in Fe(NO\(_3\))\(_3\).
- This change indicates that Fe is oxidized (loss of electrons).
- For Ag:
- In the reactants, Ag has an oxidation number of +1 (in AgNO\(_3\)).
- In the products, Ag changes to an oxidation number of 0 (elemental Ag).
- This change indicates that Ag is reduced (gain of electrons).
3. Identify the species that is reduced:
- The species that undergoes a reduction is the one that gains electrons, which, in this case, is the silver ion \( \text{Ag}^+ \). It changes from an oxidation number of +1 in AgNO\(_3\) to 0 in elemental Ag.
Therefore, the correct answer is:
[tex]\[ \boxed{D. \text{Ag}^+} \][/tex]
