Determine whether or not each metal dissolves in 1 M HNO3. For those metals that do dissolve, write a balanced redox reaction showing what happens when the metal dissolves. a. Ag

Verified step by step guidance

1. Identify the nature of the reaction: Determine if the metal can be oxidized by nitric acid (HNO3).
2. Consider the standard reduction potentials: Compare the standard reduction potential of Ag+ to that of the nitrate ion (NO3-) in acidic solution.
3. Predict the reaction: If the metal has a lower reduction potential than the nitrate ion, it will dissolve, and a redox reaction will occur.
4. Write the half-reactions: Write the oxidation half-reaction for Ag and the reduction half-reaction for NO3- in acidic conditions.
5. Balance the redox reaction: Combine the half-reactions, ensuring that electrons, atoms, and charges are balanced.

Key Concepts

Redox Reactions

Redox reactions involve the transfer of electrons between two species, resulting in oxidation (loss of electrons) and reduction (gain of electrons). In the context of metals dissolving in acids, the metal typically undergoes oxidation, while the acid's ions are reduced. Understanding these processes is crucial for predicting the behavior of metals in acidic solutions.

Metal Reactivity

The reactivity of metals varies significantly, influencing their ability to dissolve in acids like HNO3. More reactive metals, such as zinc or magnesium, readily dissolve, while less reactive metals, like silver (Ag), may not. This concept is essential for determining whether a specific metal will react with nitric acid and form a solution.

Balanced Chemical Equations

A balanced chemical equation represents a chemical reaction with equal numbers of each type of atom on both sides of the equation. For redox reactions involving metal dissolution, it is important to balance the oxidation and reduction half-reactions to accurately depict the process. This ensures that the conservation of mass and charge is maintained in the reaction.