Which statements are true? Correct any that are false. (a) In a voltaic cell, the anode is negative relative to the cathode. (b) Oxidation occurs at the anode of a voltaic or electrolytic cell. (c) Electrons flow into the cathode of an electrolytic cell. (d) In a voltaic cell, the surroundings do work on the system. (e) A metal that plates out of an electrolytic cell appears on the cathode. (f) In an electrochemical cell, the electrolyte provides a solution of mobile electrons.

In electrochemical cells, understanding the roles of the anode and cathode is crucial. The anode and cathode serve different purposes based on the type of cell in question, whether it is a voltaic (or galvanic) cell or an electrolytic cell.

In a **voltaic cell**, which generates electrical energy from spontaneous chemical reactions, the anode is the electrode where oxidation takes place. Here, electrons are released, causing the anode to be negative relative to the cathode. The cathode is where reduction occurs, meaning it gains electrons, making it positively charged.

In contrast, in an **electrolytic cell**, which uses electrical energy to drive non-spontaneous reactions, the anode is still the site of oxidation. However, it is positive because it is connected to the positive terminal of an external power source. The cathode in an electrolytic cell is negative, as it is connected to the negative terminal and is where reduction happens, leading to a gain of electrons.

Understanding whether the anode or cathode is positive or negative is essential because it determines the flow and direction of electrons and the overall reaction dynamics.

Oxidation and reduction are fundamental to the function of electrochemical cells. These processes occur simultaneously in what is known as a redox (reduction-oxidation) reaction.

**Oxidation** is the process where an atom or ion loses electrons. Consequently, this occurs at the anode in both voltaic and electrolytic cells. For example, in a zinc-copper voltaic cell, zinc undergoes oxidation as it loses electrons to form zinc ions:
\[ \text{Zn (s)} \rightarrow \text{Zn}^{2+} + 2e^- \ \text{[Oxidation]} \ \text{- At the Anode} \ \text{In Voltaic cell: Negative} \ \text{In Electrolytic cell: Positive} \]
**Reduction** is the process where an atom or ion gains electrons, occurring at the cathode. For instance, in the same voltaic cell, copper ions gain electrons to form copper metal:
\[ \text{Cu}^{2+} + 2e^- \rightarrow \text{Cu (s)} \ \text{[Reduction]} \ \text{- At the Cathode} \ \text{In Voltaic cell: Positive} \ \text{In Electrolytic cell: Negative} \]
Understanding these concepts helps in predicting and analyzing the behavior of different electrochemical reactions and their applications in real-world scenarios.

The electrolyte in an electrochemical cell has a vital role: it allows the movement of ions within the cell, ensuring that the circuit remains complete.

In both **voltaic** and **electrolytic cells**, the electrolyte facilitates ion transport between the anode and the cathode, but it does not conduct electrons; that's the job of the external circuit. Electrons flow through the wire connecting the electrodes, while ions move through the electrolyte, completing the circuit.

Electrolytes are typically ionic compounds like salts, acids, or bases dissolved in a suitable solvent (often water). For example, in a salt bridge of a voltaic cell, the electrolyte might be a salt solution like potassium nitrate (KNO₃). The function of the salt bridge is to maintain electrical neutrality by allowing ions to flow between the two half-cells, preventing the solutions from becoming charged.

In an **electrolytic cell**, the electrolyte is essential for carrying the current driven by an external power source. It ensures that the positive ions move toward the negative cathode, where they gain electrons (reduction), and negative ions move toward the positive anode, where they lose electrons (oxidation).

To sum up, the proper function of an electrolyte ensures the smooth operation of both voltaic and electrolytic cells, enabling continuous redox reactions.