Electrolysis is a captivating chemical process where electrical energy drives a non-spontaneous chemical reaction. For aqueous solutions, several factors influence the outcome of the electrolysis, resulting in various products based on the solution's nature and the conditions present.
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Deduction of Products from the Electrolysis of Aqueous Solutions
Determining the products formed during electrolysis depends on the cations and anions present in the solution, and other factors such as the electrode type and solution concentration.
Cations
- Hydrogen ions (H+): If the cation in the solution is more reactive than hydrogen, hydrogen ions will be reduced at the cathode to produce hydrogen gas.
- Equation: 2H+ + 2e- -> H2 (g)
- For instance, in the electrolysis of dilute sulphuric acid, hydrogen gas evolves at the cathode due to the reduction of hydrogen ions.
- Metal ions: Less reactive metals than hydrogen are deposited at the cathode. Here, the metal ions gain electrons and get reduced to their respective metals.
- Example: In a copper(II) sulphate solution, copper ions gain electrons and are deposited as copper metal at the cathode.
Anions
- Hydroxide ions (OH-): If the solution contains anions more reactive than hydroxide ions, oxygen gas will be the product at the anode due to the oxidation of hydroxide ions.
- Example: During sodium chloride solution electrolysis, chlorine gas, not oxygen, is produced at the anode because chloride ions are less reactive than hydroxide ions.
- Halide ions (e.g., Cl-, Br-, I-): Halogens are released at the anode if they're less reactive than hydroxide ions.
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Understanding the Electrolysis of Water
Water's electrolysis is essential in grasping the broader aspects of electrolysis because it sets the foundation for many aqueous electrolytic reactions.
Electrodes
- Cathode: Hydrogen ions in water get reduced to form hydrogen gas.
- 2H+ + 2e- -> H2 (g)
- Anode: Hydroxide ions undergo oxidation, producing oxygen gas.
- 4OH- -> O2 (g) + 2H2O + 4e-
Significance
Producing hydrogen via water's electrolysis is a vital method for obtaining clean fuel. The significance of this process grows as the global shift towards sustainable energy sources intensifies.
Water electrolysis- H2SO4 is often added to improve conductivity as water alone is a poor conductor of electricity so adding an electrolyte helps ions to move more freely. H2SO4 doesn't directly participate in the electrolysis.
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Electrolysis of Other Aqueous Solutions
Solutes can affect the products formed during water's electrolysis. Sodium chloride and copper(II) sulphate solutions are two prevalent examples.
Sodium Chloride Solution
- Cathode: Hydrogen gas is formed due to the reduction of hydrogen ions.
- Anode: Chlorine gas is produced due to the oxidation of chloride ions.
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Copper(II) Sulphate Solution
- Cathode: Copper ions accept electrons to form copper metal.
- Anode: Here, water molecules are oxidised, producing oxygen gas, and the solution turns progressively acidic.
Effects of Concentration and Electrode Nature on Electrolysis
Concentration
Concentration is pivotal in determining electrolysis products.
- High concentration of ions: When there's a high concentration of a specific ion, it's more likely to get discharged at the electrodes. For instance, in a concentrated sodium chloride solution, chloride ions are preferentially discharged, producing chlorine gas at the anode.
- Low concentration of ions: In dilute solutions, water is more likely to undergo electrolysis. As a result, hydrogen and oxygen gases are typically formed.
Electrode Nature
Electrodes, where the redox reactions occur, can considerably affect the outcome of the electrolysis.
- Inert electrodes (e.g., graphite or platinum): These don't participate in the electrolytic reactions. Thus, the ions in the solution dictate the products. These electrodes are usually employed when studying the electrolysis of a specific compound without interference.
- Reactive electrodes (e.g., copper or zinc): These can engage in the reactions. A copper anode in a copper(II) sulphate solution will dissolve over time, introducing more copper(II) ions to the solution.
In summary: Both the concentration of the solution and the type of electrode play critical roles in determining the outcomes of electrolysis. Predicting these outcomes requires an in-depth understanding of reactivity series, electrode potentials, and redox reactions principles. As students delve into electrochemistry, they'll uncover these concepts' relevance in real-world applications, from energy storage solutions to metallurgy.
FAQ
Overpotential refers to the extra voltage (over and above the theoretical voltage) required during electrolysis to drive the reaction at a noticeable rate. The nature of the electrode material can influence this overpotential. Inert electrodes like platinum or graphite have higher overpotentials for the evolution of oxygen or hydrogen as compared to more reactive metals. The material of the electrode can influence the ease with which a reaction takes place on its surface. A metal electrode that can form an oxide layer, for instance, may show a higher overpotential for oxygen evolution because the oxide layer hinders the process.
In a molten state, the ionic compound is purely composed of its constituent ions without the presence of water. Therefore, water doesn't compete in the redox processes during electrolysis. For example, in the molten state electrolysis of sodium chloride, sodium ions are reduced at the cathode to produce sodium metal, while chloride ions are oxidised at the anode to produce chlorine gas. However, in an aqueous solution of sodium chloride, as discussed earlier, water interferes, leading to the evolution of hydrogen gas at the cathode instead of sodium metal deposition.
Electrolysis, particularly of aqueous solutions, requires a significant amount of energy mainly in the form of electricity. When metals are obtained from their ores using electrolysis, the energy costs can be high due to the required voltage and the duration of the process. Reduction methods, on the other hand, such as the reduction of metal oxides with carbon in a blast furnace, often utilise chemical reactions that release energy. These exothermic reactions can aid in the overall process, making it more energy-efficient. Moreover, the raw materials like coke used in blast furnaces for reduction are often cheaper than the electricity costs associated with electrolysis.
Pure water is a poor conductor of electricity due to its very low concentration of ions. For electrolysis to occur spontaneously, there must be sufficient ions present to carry the current. Additionally, the reduction potential for hydrogen ions (from water) to form hydrogen gas and the oxidation potential for water to form oxygen gas are not favourable enough to occur spontaneously. To drive the non-spontaneous electrolysis of pure water, an external voltage (greater than the thermodynamic threshold) is required. In practical scenarios, to aid the electrolysis process, an electrolyte is added to increase the conductivity and reduce the voltage required.
Sodium is a highly reactive metal, and in an aqueous environment like that of a sodium chloride solution, water is also present as a potential reactant. During the electrolysis of an aqueous sodium chloride solution, there are two potential cations that can be reduced at the cathode: sodium ions (Na+) and hydrogen ions (H+ from water). However, sodium is more reactive than hydrogen, meaning it has a greater tendency to lose electrons (get oxidised) than to gain them (get reduced). As a result, hydrogen ions are preferentially reduced at the cathode, producing hydrogen gas, while the sodium ions remain in the solution.
Practice Questions
The electrolysis of a dilute aqueous solution of sodium chloride involves the discharge of hydrogen ions at the cathode and chloride ions at the anode. At the cathode, hydrogen ions (H+) from water are reduced to form hydrogen gas through the half-equation: 2H+ + 2e- -> H2 (g). At the anode, chloride ions (Cl-) are oxidised to form chlorine gas with the half-equation: 2Cl- -> Cl2 (g) + 2e-. If a concentrated solution of sodium chloride was used, chloride ions would still be preferentially discharged at the anode due to their abundance, producing chlorine gas. However, the increased concentration of sodium ions does not affect the products formed at the cathode, as sodium is more reactive than hydrogen and won't get deposited.
When copper(II) sulphate undergoes electrolysis using copper electrodes, the nature of the electrode significantly affects the outcome. At the cathode, copper ions (Cu2+) from the solution accept two electrons to form copper metal through the half-equation: Cu2+ + 2e- -> Cu. This results in a deposition of copper metal on the cathode. At the anode, which is made of copper metal, copper atoms lose two electrons to form copper ions (Cu2+) with the half-equation: Cu -> Cu2+ + 2e-. This leads to the anode dissolving over time. Therefore, in this scenario, the reactive copper electrodes play an active role in the electrolytic process, unlike inert electrodes which do not participate in the reactions.
