Answered

Complete and balance the equation for this redox reaction in basic solution. The phases are optional.

For a particular redox reaction, NO is oxidized to [tex]$NO_3^{-}$[/tex] and [tex]$Ag^{+}$[/tex] is reduced to Ag.

Balanced reaction:
[tex]
NO + Ag^{+} \longrightarrow NO_3^{-} + Ag
[/tex]



Answer :

GinnyAnswer
Certainly! Let's complete and balance the given redox reaction in a basic solution step-by-step.

### Step 1: Write the unbalanced half-reactions

We have:
- Oxidation: [tex]\( \text{NO} \rightarrow \text{NO}_3^- \)[/tex]
- Reduction: [tex]\( \text{Ag}^+ \rightarrow \text{Ag} \)[/tex]

### Step 2: Balance the atoms in each half-reaction except for oxygen and hydrogen

- For the oxidation half-reaction:
[tex]\[ \text{NO} \rightarrow \text{NO}_3^- \][/tex]

- For the reduction half-reaction:
[tex]\[ \text{Ag}^+ \rightarrow \text{Ag} \][/tex]

### Step 3: Balance oxygen atoms by adding [tex]\( \text{H}_2\text{O} \)[/tex]

For the oxidation half-reaction:
[tex]\[ \text{NO} \rightarrow \text{NO}_3^- \][/tex]
We need 3 oxygen atoms on the left side:
[tex]\[ \text{NO} + 2 \text{H}_2\text{O} \rightarrow \text{NO}_3^- \][/tex]

### Step 4: Balance hydrogen atoms by adding [tex]\( \text{H}^+ \)[/tex]

For the oxidation half-reaction:
[tex]\[ \text{NO} + 2 \text{H}_2\text{O} \rightarrow \text{NO}_3^- \][/tex]
We need 4 hydrogen atoms on the right side:
[tex]\[ \text{NO} + 2 \text{H}_2\text{O} \rightarrow \text{NO}_3^- + 4 \text{H}^+ \][/tex]

### Step 5: Balance the charges by adding electrons

For the oxidation half-reaction:
[tex]\[ \text{NO} + 2 \text{H}_2\text{O} \rightarrow \text{NO}_3^- + 4 \text{H}^+ + 3 e^- \][/tex]
We added 3 [tex]\( e^- \)[/tex] to balance the charge.

For the reduction half-reaction:
[tex]\[ \text{Ag}^+ + e^- \rightarrow \text{Ag} \][/tex]

### Step 6: Equalize the number of electrons transferred in both half-reactions

Multiply the reduction half-reaction by 3:
[tex]\[ 3 (\text{Ag}^+ + e^- \rightarrow \text{Ag}) \][/tex]
This gives:
[tex]\[ 3 \text{Ag}^+ + 3 e^- \rightarrow 3 \text{Ag} \][/tex]

### Step 7: Add the two half-reactions together

[tex]\[ \text{NO} + 2 \text{H}_2\text{O} \rightarrow \text{NO}_3^- + 4 \text{H}^+ + 3 e^- + 3 \text{Ag}^+ + 3 e^- \rightarrow 3 \text{Ag} \][/tex]
Combine them to get:
[tex]\[ \text{NO} + 2 \text{H}_2\text{O} + 3 \text{Ag}^+ \rightarrow \text{NO}_3^- + 4 \text{H}^+ + 3 \text{Ag} \][/tex]

### Step 8: Convert to a basic solution by adding [tex]\( \text{OH}^- \)[/tex]

Add [tex]\( 4 \text{OH}^- \)[/tex] to both sides to neutralize [tex]\( 4 \text{H}^+ \)[/tex]:
[tex]\[ \text{NO} + 2 \text{H}_2\text{O} + 3 \text{Ag}^+ + 4 \text{OH}^- \rightarrow \text{NO}_3^- + 4 \text{H}^+ + 3 \text{Ag} + 4 \text{OH}^- \][/tex]
Simplify the water molecules:
[tex]\[ \text{NO} + 3 \text{Ag}^+ + 4 \text{OH}^- \rightarrow \text{NO}_3^- + 3 \text{Ag} + 2 \text{H}_2\text{O} \][/tex]

### Final balanced equation

The balanced redox reaction in basic solution is:
[tex]\[ \text{NO} + 3 \text{Ag}^+ + 4 \text{OH}^- \rightarrow \text{NO}_3^- + 3 \text{Ag} + 2 \text{H}_2\text{O} \][/tex]

This is the complete and balanced equation for the given reaction.