Topic 5: Electrochemistry and Industrial Electrolytic Processes

Fundamentals of Electrochemistry and Electrolytic Cells

Electrochemistry is defined as the study of chemical processes that cause electrons to move. This scientific field has broad practical applications, most notably in the processes of electrolysis and electroplating. At the heart of these applications is the electrolytic cell, which is a type of electrochemical cell designed to convert electrical energy into chemical energy. This conversion occurs because the supply of electricity initiates a chemical reaction that would otherwise not happen spontaneously. This specific process, driven by an external electrical source to decompose chemical compounds, is known as electrolysis.

Electrolysis operates on general principles regarding the state of the matter involved. It is the process by which a molten ionic compound or an aqueous solution of an ionic compound is broken down by an electric current. It is important to note that covalent compounds do not conduct electricity and therefore cannot undergo electrolysis. Similarly, ionic compounds in their solid state are unable to conduct electricity because their ions are fixed in a rigid lattice and cannot move to carry a charge. For electrolysis to occur, the ionic substance must be either molten or in a solution, as these states allow the ions to move freely and act as charge carriers.

Key Components and Charge Transfer in Electrolytic Cells

Several key terms are essential for understanding the mechanics of a simple electrolytic cell. An electrode is a rod, typically made of metal or graphite, that conducts electricity into or out of the electrolyte. These are often made of inert materials. The electrolyte itself is the ionic compound, either molten or dissolved, that conducts the electricity and is subsequently decomposed during the process. The anode is the positive electrode where oxidation takes place, while the cathode is the negative electrode where reduction occurs. Ions within the electrolyte are categorized by their charge: anions are negatively charged and are attracted to the anode, whereas cations are positively charged and are attracted to the cathode.

In terms of charge transfer, electrolysis focuses on the movement of electrons rather than the direction of conventional current. Electrons flow through the external circuit from the negative terminal of the power supply to the positive terminal. Specifically, they flow alphabetically from the anode to the cathode. While electrons serve as the charge carriers in the external circuit, the ions serve as the charge carriers within the electrolyte itself. Electrons flow from the power supply's negative terminal to the cathode, giving it a negative charge. These electrons do not pass through the solution. Instead, at the cathode, cations gain these electrons to form atoms. Simultaneously, anions move to the anode and lose electrons, which then flow back through the external circuit to the positive terminal of the power supply, completing the circuit.

Thermodynamics and Redox Principles in Electrolysis

The chemical reaction occurring within an electrolytic cell is a redox reaction, which involves the simultaneous processes of reduction and oxidation. Reduction is defined as the gain of electrons from another atom, molecule, or ion, while oxidation is the loss of electrons. These processes are codependent; electrons are transferred from the substance being oxidized to the substance being reduced. Two mnemonics are commonly used to remember the locations of these reactions: ANOX (Anode = Oxidation) and REDCAT (Reduction = Cathode).

General rules govern the products formed at the electrodes. At the negative electrode (cathode), either metals or hydrogen gas are formed. The specific outcome depends on the reactivity series: if the metal in the solution is less reactive than hydrogen, such as copper or silver, the metal is produced. However, if the metal is more reactive than hydrogen, such as sodium or magnesium, hydrogen gas is produced instead. At the positive electrode (anode), non-metals (other than hydrogen) are formed from negative ions. Electrolysis is essentially an electrochemical decomposition reaction where ions move toward the electrodes and are discharged: negative ions at the positive electrode and positive ions at the negative electrode.

Electrolysis Case Study: Molten Lead(II) Bromide

The electrolysis of molten lead(II) bromide ($PbBr_2$) serves as a primary example of decomposing a binary ionic compound. Because lead(II) bromide is insoluble in water, it must be heated until molten to allow the ions to dissociate and move freely. In this setup, two graphite rods are used as inert electrodes. When the power is applied, the follow reactions occur: at the anode, negative bromide ($Br^-$) ions are attracted and undergo oxidation, losing electrons to form bromine molecules. This results in bubbling and the release of brown bromine gas. The half-reaction at the anode is $2Br^-_{(aq)} \rightarrow Br_{2(g)} + 2e^-$.

At the cathode, positive lead ($Pb^{2+}$) ions are attracted and undergo reduction by gaining electrons. This results in the formation of grey lead metal, which deposits as a coating on the electrode surface. The half-reaction at the cathode is $Pb^{2+}_{(aq)} + 2e^- \rightarrow Pb_{(s)}$. The total ionic equation for the process is $2Br^-_{(aq)} + Pb^{2+}_{(aq)} \rightarrow Br_{2(g)} + Pb_{(s)}$, and the overall reaction is represented as $PbBr_{2(l)} \rightarrow Pb_{(s)} + Br_{2(g)}$.

Electrolysis of Aqueous Solutions: Sulfuric Acid and Hydrochloric Acid

When electrolyzing dilute sulfuric acid ($H_2SO_4$) using inert platinum or carbon electrodes, the presence of water must be considered. The ions present include $H^+$ and $SO_4^{2-}$ from the acid, and $H^+$ and $OH^-$ from the water. At the cathode, hydrogen ions from both sources are reduced to form hydrogen gas: $2H^+_{(aq)} + 2e^- \rightarrow H_{2(g)}$. At the anode, hydroxide ions ($OH^-$) are oxidized more readily than sulfate ions, producing oxygen gas and water: $4OH^-_{(aq)} \rightarrow 2H_2O_{(l)} + O_{2(g)} + 4e^-$. The sulfate ion is a spectator ion and does not participate. The overall reaction is the decomposition of water: $2H_2O_{(l)} \rightarrow 2H_{2(g)} + O_{2(g)}$.

In the electrolysis of concentrated hydrochloric acid ($HCl$), the acid ionizes into $H^+$ and $Cl^-$. At the cathode, hydrogen ions undergo reduction to form hydrogen gas. At the anode, chloride ions are oxidized, producing bubbles of yellowish-green chlorine gas. This specific experiment requires a well-ventilated room because chlorine gas is highly toxic. The reactions follow the standard redox patterns for hydrogen and halides in concentrated solutions.

Electrolysis of Concentrated Aqueous Sodium Chloride (Brine)

The electrolysis of concentrated aqueous sodium chloride, or brine, involves the dissociation of $NaCl$ into $Na^+$ and $Cl^-$ and the ionization of water into $H^+$ and $OH^-$. At the cathode, hydrogen ions are reduced to form hydrogen gas: $2H_2O + 2e^- \rightarrow H_{2(g)} + 2OH^-_{(aq)}$. Although sodium ions are present, hydrogen is produced because sodium is more reactive. At the anode, chloride ions are oxidized to form chlorine gas: $2Cl^- \rightarrow Cl_{2(g)} + 2e^-$. The remaining $Na^+$ and $OH^-$ ions stay in the solution, forming sodium hydroxide ($NaOH$), which is a valuable industrial byproduct. This process typically utilizes a diaphragm or membrane cell to keep the products separate while maintaining temperature and concentration via direct current.

Predicting Products and the Effect of Concentration

Predicting the products of aqueous electrolysis requires evaluating the relative reactivity of the ions and their concentrations. Water molecules always dissociate into $H^+$ and $OH^-$, and these must be factored in. At the anode, if halide ions ($Cl^-$, $Br^-$, $I^-$) are present in high concentration, the halogen is produced. If the solution is dilute or lacks halides, hydroxide ions are discharged to form oxygen. For example, concentrated barium chloride produces chlorine gas at the anode, while dilute barium chloride produces oxygen. At the cathode, the reactivity series is the deciding factor: if the metal is above hydrogen in the reactivity series, hydrogen gas is produced; if the metal is below hydrogen (like copper), the metal itself plates the electrode.

In the case of copper(II) sulfate ($CuSO_4$) with carbon electrodes, copper metal forms at the cathode as it is lower in the reactivity series than hydrogen. Oxygen gas is released at the anode from the discharge of hydroxide ions. The sulfate ion remains as a spectator. This principle is utilized to purify impure copper, where copper ions are transferred from an impure anode to a pure cathode.

Principles and Applications of Electroplating

Electroplating is the process of coating a more reactive metal with a thin layer of a less reactive metal, such as silver, gold, or chromium. This is achieved using an electrolytic cell where the article to be plated (the substrate) is placed at the cathode (the negative electrode). The anode is made of the metal that will be used for plating. The electrolyte is a solution containing the salts of the plating metal, such as copper sulfate or silver nitrate, to facilitate the flow of ions. When a power source applies current, the metal at the anode oxidizes into ions, which then travel through the solution to the cathode, where they are reduced and deposited as a solid layer.

Common applications of electroplating include creating gold-plated jewelry to reduce costs and coating steel cans with tin to prevent rusting. Chromium plating is used on car bumpers, kettles, and bath taps to prevent corrosion, resist wear and scratching, and provide a shiny aesthetic. Silver-plating is also frequently applied to cutlery. For silver-plating a spoon, the spoon is connected to the cathode in a silver nitrate ($AgNO_3$) solution, while a pure silver rod serves as the anode.

Industrial Applications: Aluminium and Chlorine Production

The industrial manufacture of aluminium from bauxite relies on electrolysis. Bauxite, primarily consisting of aluminium oxide ($Al_2O_3$), is purified with sodium hydroxide to remove impurities like silicon dioxide and iron(III) oxide. The resulting pure aluminium oxide has a melting point exceeding $2000^{\circ}C$. To save energy, it is dissolved in molten cryolite, which significantly lowers the melting temperature. The electrolysis takes place in a cell with carbon anodes and a carbon lining that serves as the cathode. Aluminium forms at the cathode, and oxygen forms at the anode.

Another major industrial process is the production of chlorine and sodium hydroxide from concentrated brine. This process requires specific conditions, including the use of a membrane or diaphragm cell to divide the cell in two. Direct current is applied to ensure the oxidation of chloride ions at the anode (releasing chlorine gas) and the reduction of water at the cathode (releasing hydrogen gas and leaving hydroxide ions in solution). The resulting sodium hydroxide solution is then tapped off for use in various industries.

Questions & Discussion

Why is cryolite used in the manufacture of aluminium? Cryolite is used because aluminium oxide has an extremely high melting point of over $2000^{\circ}C$. Dissolving the oxide in molten cryolite lowers the melting temperature, which reduces the amount of energy required for the process and makes it significantly less expensive.

Why do you need to replace the anode regularly in the aluminium extraction process? The anode in the aluminium extraction process is made of carbon. During electrolysis, oxygen is produced at the anode. At the high temperatures of the cell, the oxygen reacts with the carbon anode to form carbon dioxide gas. This gradually consumes the anode, necessitating its regular replacement.