Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron flow. a. Ni2+(aq) + Mg(s)-Ni(s) + Mg2+(aq) b. 2 H+(aq) + Fe(s)-H2(g) + Fe2+(aq) c. 2 NO3-(aq) + 8 H+(aq) + 3 Cu(s)-2 NO(g) + 4 H2O(l) + 3 Cu2+(aq)

At the heart of a voltaic cell are redox reactions, which are chemical reactions involving the transfer of electrons between two species. These reactions consist of two half-reactions: oxidation, where electrons are lost, and reduction, where electrons are gained. Understanding redox reactions is essential in identifying the changes occurring at the electrodes of a voltaic cell.

In a typical redox reaction within a voltaic cell, one metal is oxidized, losing electrons, while another substance is reduced, gaining those electrons. For example, in the first given exercise reaction, magnesium (Mg) is oxidized to magnesium ions (Mg²⁺), and nickel ions (Ni²⁺) are reduced to nickel (Ni) metal. Every redox reaction involves movement of electrons from the substance being oxidized to the substance being reduced.

Electrochemical cells, including voltaic cells, are devices that convert chemical energy into electrical energy through redox reactions. Each cell is composed of two half-cells, which are essentially beakers containing a metal electrode submerged in a solution of its ions. These half-cells are connected by a conductor and a salt bridge, allowing ions to flow between them while maintaining electrical neutrality.

The external circuit provides a path for electrons to move from the anode to the cathode. This movement of electrons through the circuit can be harnessed to do work, such as powering a light bulb. The salt bridge completes the cell’s electrical circuit by permitting the flow of ions, thereby preventing the buildup of charge that would otherwise stop the reaction.

Identifying the anode and cathode is a fundamental aspect of understanding how a voltaic cell operates. The anode is where oxidation takes place—it's where the metal loses electrons and becomes positively charged ions. Conversely, the cathode is where reduction occurs—the electrons flow into this electrode, thereby reducing the ions in the solution to a neutral state or a lower oxidation state.

In the exercises, the substances that were oxidized, such as Mg, Fe, and Cu, served as anodes. These metals released electrons into the external circuit. The cathode was identified as the site where Ni²⁺, H⁺, and NO₃^- were reduced. An inert electrode, such as platinum (Pt), is often used when the cathode reaction does not involve the deposition of metal on the electrode, as in the cases of hydrogen gas formation or nitrogen oxide formation.

The direction of electron flow in a voltaic cell is from the anode to the cathode through the external circuit. It is crucial to understand that electrons flow toward the more positive potential. The flow of electrons is facilitated by the potential difference between the electrodes, which is driven by the cell’s redox reactions.

In the sketched cells for the textbook exercises, the electrons left the metal at the anode and traveled through the wire to reach the cathode. The ability to predict the direction of electron flow helps in setting up the external circuit of a voltaic cell and is key to harnessing the electric current produced by the cell.