An H₂/H⁺ half-cell (anode) and an Ag⁺/Ag half cell (cathode) are connected by a wire and a salt bridge. (b) Write balanced equations for the electrode and overall cell reactions.

Verified step by step guidance

1. Identify the half-reactions for each half-cell: the anode (H₂/H⁺) and the cathode (Ag⁺/Ag).
2. Write the oxidation half-reaction for the anode: H₂ gas is oxidized to H⁺ ions.
3. Write the reduction half-reaction for the cathode: Ag⁺ ions are reduced to Ag metal.
4. Balance the electrons in the half-reactions to ensure the same number of electrons are transferred in both half-reactions.
5. Combine the balanced half-reactions to write the overall balanced cell reaction.

Key Concepts

Electrochemical Cells

Electrochemical cells consist of two half-cells, where oxidation occurs at the anode and reduction at the cathode. In this setup, the H2/H+ half-cell serves as the anode, where hydrogen gas is oxidized to protons, while the Ag+/Ag half-cell acts as the cathode, where silver ions are reduced to solid silver. Understanding the flow of electrons and the role of each half-cell is crucial for writing balanced equations.

Oxidation and Reduction

Oxidation is the loss of electrons, while reduction is the gain of electrons. In the context of the given half-cells, the oxidation reaction involves H2 losing electrons to form H+ ions, and the reduction reaction involves Ag+ ions gaining electrons to form Ag. Recognizing these processes is essential for balancing the overall cell reaction and understanding the electron transfer that drives the electrochemical cell.

Balancing Redox Reactions

Balancing redox reactions involves ensuring that the number of atoms and the charge are equal on both sides of the equation. This can be achieved by identifying the oxidation and reduction half-reactions, balancing them separately, and then combining them. In this case, the balanced equations for the H2 oxidation and Ag+ reduction must be combined to represent the overall cell reaction accurately.