To determine the oxidizing agent in this redox reaction, we need to analyze the changes in oxidation states of the elements involved in the reaction:
[tex]\[ Cu + 2 AgNO_3 \rightarrow 2 Ag + Cu(NO_3)_2 \][/tex]
1. Identifying the oxidation states:
- Copper (Cu): In its elemental form, copper has an oxidation state of 0.
- Silver in Silver Nitrate (AgNO_3): The nitrate ion ([tex]\(NO_3^-\)[/tex]) has a charge of -1. Since the compound is neutral, the oxidation state of Ag must be +1 to balance the -1 from the nitrate ion.
- Silver (Ag): In its elemental form in the products, silver has an oxidation state of 0.
- Copper in Copper(II) Nitrate ([tex]\(Cu(NO_3)_2\)[/tex]): The nitrate ion ([tex]\(NO_3^-\)[/tex]) has a charge of -1. Since there are two nitrate ions, the total negative charge contributed is -2. Therefore, the oxidation state of copper must be +2 to balance the charges in the compound.
2. Determining the changes in oxidation states:
- Copper (Cu): The oxidation state changes from 0 in the reactants to +2 in the products.
- Silver (Ag): The oxidation state changes from +1 in the reactants (as part of AgNO_3) to 0 in the products.
3. Identifying oxidation and reduction:
- Oxidation: The substance that loses electrons and increases its oxidation state. Copper (Cu) goes from 0 to +2, so it is oxidized.
- Reduction: The substance that gains electrons and decreases its oxidation state. Silver ion (Ag^+) goes from +1 to 0, so it is reduced.
4. Identifying the oxidizing agent:
- The oxidizing agent is the substance that is reduced in the reaction. Here, the silver ion (Ag^+) from the silver nitrate (AgNO_3) is reduced.
Therefore, the substance acting as the oxidizing agent in this redox reaction is:
B. [tex]\( AgNO_3 \)[/tex]
