Voltaic and Electrolytic Cells Review Items

Core Principles of Voltaic Cells and Electrode Dynamics

A voltaic cell is a type of electrochemical cell that facilitates a spontaneous oxidation-reduction reaction to generate electrical energy. To identify the components of a voltaic cell, students must utilize Table J, the activity series. The metal that is higher on Table J is more active and serves as the anode, whereas the metal lower on the table serves as the cathode. The fundamental mnemonics used to remember electrode processes are AN OX and RED CAT. AN OX stands for anode oxidation, meaning oxidation always occurs at the anode, leading to a loss of mass as the solid metal is converted into aqueous ions. RED CAT stands for reduction cathode, meaning reduction always occurs at the cathode, leading to an increase in mass as aqueous ions are reduced into solid metal and deposited on the electrode.

In a typical voltaic cell involving zinc and copper, the oxidation half-reaction at the zinc anode is represented as $Zn_{(s)} \rightarrow Zn^{2+}_{(aq)} + 2e^-$. Conversely, the reduction half-reaction at the copper cathode is represented as $Cu^{2+}_{(aq)} + 2e^- \rightarrow Cu_{(s)}$. The overall net ionic equation for this spontaneous process is $Zn_{(s)} + Cu^{2+}_{(aq)} \rightarrow Cu_{(s)} + Zn^{2+}_{(aq)}$. Electrons always flow through the external wire from the anode to the cathode. A crucial component of the voltaic cell is the salt bridge, which functions to allow the migration of ions between the two separate containers. This migration maintains electrical neutrality within the half-cells, preventing the buildup of charge that would otherwise stop the reaction.

Principles of Electrolytic Cells and Electroplating Processes

Electrolytic cells differ from voltaic cells because they require an outside energy source, such as a battery, to force a non-spontaneous chemical reaction to occur. In these cells, electrical energy is converted into chemical energy. Despite being non-spontaneous, the sites of chemical activity remain consistent with the general rules of electrochemistry: oxidation occurs at the anode and reduction occurs at the cathode. In an electrolytic setup, the anode is attached to the positive end of the power source, while the cathode is attached to the negative end. Electrons are forced to flow from the positive anode to the negative cathode by the external power supply.

Electroplating is a primary application of electrolytic cells, such as plating a spoon with nickel or a key with copper. In the example of plating a silver-colored spoon with nickel, the spoon acts as the cathode. When the switch is closed, the nickel anode undergoes oxidation ($Ni_{(s)} \rightarrow Ni^{2+}_{(aq)} + 2e^-$) and loses mass. Simultaneously, the nickel ions in the solution ($Ni^{2+}_{(aq)}$) are attracted to the negative cathode (the spoon), where they gain electrons and undergo reduction ($Ni^{2+}_{(aq)} + 2e^- \rightarrow Ni_{(s)}$). This results in the spoon gaining mass as the solid nickel metal is deposited onto its surface. Similarly, when plating a key with copper in a $CuSO_{4(aq)}$ solution, the reduction of copper ions ($Cu^{2+}_{(aq)} + 2e^- \rightarrow Cu_{(s)}$) occurs at the cathode electrode.

Comparative Analysis of Voltaic and Electrolytic Systems

There are several distinct differences between voltaic and electrolytic cells that are essential for full understanding. A voltaic cell converts chemical energy into electrical energy and functions as a battery, while an electrolytic cell requires a battery to convert electrical energy into chemical energy. In terms of spontaneity, voltaic reactions are always spontaneous, meaning the Gibbs free energy change is negative, whereas electrolytic reactions are non-spontaneous. The physical setup also differs: voltaic cells usually consist of two separate containers connected by a salt bridge, while electrolytic cells typically occur within a single container and do not require a salt bridge.

Polars of the electrodes also shift between the two systems. In a voltaic cell, the anode is negative and the cathode is positive. In an electrolytic cell, the anode is positive and the cathode is negative. However, the direction of electron flow remains the same in both: electrons always move from the anode to the cathode through the external circuit. Furthermore, the positive ions (cations) in the electrolytic solution are always attracted to the negative cathode, where they undergo reduction.

Redox Half-Reactions, Electron Transfer, and Stoichiometric Balancing

In any oxidation-reduction reaction, the number of electrons lost in the oxidation half-reaction must be exactly equal to the number of electrons gained in the reduction half-reaction. This principle of conservation of charge is vital when balancing redox equations. For example, in the reaction between magnesium and nitrogen, the unbalanced equation is $Mg_{(s)} + N_{2(g)} \rightarrow Mg_3N_{2(s)}$. To balance this using the smallest whole-number coefficients, one must assign three atoms of magnesium to react with one molecule of diatomic nitrogen, resulting in $3Mg_{(s)} + N_{2(g)} \rightarrow Mg_3N_{2(s)}$. Here, magnesium serves as the anode species because it is being oxidized from an oxidation state of $0$ to $+2$. The balanced half-reaction for the oxidation of magnesium is $Mg_{(s)} \rightarrow Mg^{2+} + 2e^-$.

Another example is the balanced equation $3Mg + 2Fe^{3+} \rightarrow 3Mg^{2+} + 2Fe$. In this scenario, each iron ion ($Fe^{3+}+$) gains three electrons to become neutral iron ($Fe^0$). Because there are two iron ions, a total of $6$ electrons are transferred. Thus, the coefficient for both $Fe^{3+}$ and $Fe$ is $2$. Identifying a reduction half-reaction involves looking for the gain of electrons on the reactant side, such as $Fe^{3+} + 3e^- \rightarrow Fe$. If electrons are on the product side, the reaction represents oxidation, such as $Fe \rightarrow Fe^{3+} + 3e^-$.

Oxidation Number Assignments and Chemical Periodicity

Assigning oxidation numbers is a standardized process based on specific rules. An atom in its elemental state, such as Nitrogen in $N_2$, always has an oxidation number of $0$. In a compound, the sum of the oxidation numbers must equal the overall charge of the species. For the molecule perchloric acid, $HClO_4$, the oxidation number of Hydrogen is $+1$ and Oxygen is $-2$. Since there are four oxygen atoms, the total negative contribution is $-8$. To make the molecule neutral, the oxidation number of Chlorine must be $+7$, calculated as $(+1) + (Cl) + 4(-2) = 0$.

When identifying which species is most easily oxidized, one must refer to Table J; magnesium ($Mg$) is more active and thus more easily oxidized than cobalt ($Co$), silver ($Ag$), or copper ($Cu$). Furthermore, specific trends exist for electronegative elements like oxygen. While oxygen usually has an oxidation number of $-2$, it will exhibit a positive oxidation number only when combined with an element more electronegative than itself, which is fluorine ($F$). In the case where $Cr^{3+}$ ions are reduced to $Cr^{2+}$ ions, electrons are gained, which causes the oxidation number of the chromium to decrease from $+3$ to $+2$.

Questions & Discussion

Question: What will occur when switch S is closed in the electroplating diagram showing a nickel electrode and a spoon? Response: Based on the principles of electrolytic cells, the spoon will gain mass as nickel ions are reduced onto it, and the nickel electrode ($Ni_{(s)}$) will be oxidized at the anode. The correct observation is that the spoon gains mass and the nickel metal is oxidized.

Question: Which species is oxidized in the voltaic cell involving $Mg_{(s)}$ and $Ag^{+}_{(aq)}$? Response: In the reaction $Mg_{(s)} + 2Ag^{+}_{(aq)} \rightarrow Mg^{2+}_{(aq)} + 2Ag_{(s)}$, the magnesium metal ($Mg_{(s)}$) starts with an oxidation state of $0$ and ends with $+2$. Therefore, $Mg_{(s)}$ is the species that is oxidized.

Question: In the zinc and copper electrochemical cell reaction $Zn_{(s)} + Cu^{2+}_{(aq)} \rightarrow Zn^{2+}_{(aq)} + Cu_{(s)}$, what happens as the reaction takes place? Response: As the reaction progresses, the zinc anode is oxidized into ions, meaning the mass of the $Zn_{(s)}$ electrode decreases and the concentration of $Zn^{2+}_{(aq)}$ increases. Simultaneously, the copper ions are reduced into solid metal, meaning the mass of the $Cu_{(s)}$ electrode increases while the concentration of $Cu^{2+}_{(aq)}$ decreases.