Electrolysis
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Intros
Examples
Lessons
- Use the cell potential calculation to find conditions for electrolytic cells.
For an electrolytic cell containing 1M NiI2 (aq) :
- What reaction takes place at the anode?
- What reaction takes place at the cathode?
- What is the minimum voltage needed for this reaction to work?
- Use the cell potential calculation to find conditions for electrolytic cells.
For an electrolytic cell containing 1M HI (aq) :
Topic Notes
Introduction to Electrolysis
Electrolysis is a fundamental process in chemistry and industry, playing a crucial role in various applications. This electrochemical technique uses electrical energy to drive non-spontaneous chemical reactions, enabling the production of important materials and the purification of metals. At the heart of electrolysis is the electrolytic cell, a device consisting of electrodes (anode and cathode), an electrolyte solution, and a power source. The cell facilitates the movement of ions and the transfer of electrons, resulting in chemical changes at the electrodes. Electrolysis finds widespread use in manufacturing, metal extraction, and even in everyday items like batteries. To help visualize this process, we've included an introduction video that demonstrates electrolysis in action. This video will provide a clear, step-by-step explanation of how an electrolytic cell operates, making it easier to understand this fascinating aspect of electrochemistry. Whether you're a student or simply curious about science, this overview of electrolysis will illuminate its significance in our world.
Understanding Electrolytic Cells
An electrolytic cell is a crucial electrochemical device used in various industrial processes and scientific applications. Unlike its counterpart, the galvanic cell, which generates electricity from spontaneous chemical reactions, an electrolytic cell uses electrical energy to drive non-spontaneous chemical reactions. This fundamental difference sets electrolytic cells apart in the realm of electrochemistry.
The components of an electrolytic cell are similar to those of a galvanic cell, but their functions and arrangement differ significantly. The main components include:
- Anode: In an electrolytic cell, the anode is the positive electrode. This is where oxidation occurs, and electrons are removed from the system.
- Cathode: The cathode is the negative electrode, where reduction takes place, and electrons are added to the system.
- Electrolyte: This is the conductive solution or molten substance that allows ions to flow between the electrodes.
- Power Source: Unlike galvanic cells, electrolytic cells require an external power source, typically a battery or DC power supply, to drive the non-spontaneous chemical reactions.
The energy conversion in an electrolytic cell is a key aspect of its operation. Electrical energy from the power source is converted into chemical energy, which is stored in the products of the electrolysis reaction. This process is essentially the reverse of what occurs in a galvanic cell, where chemical energy is converted to electrical energy.
To better understand the structure and function of an electrolytic cell, consider the following diagram:
In this diagram, we can see:
- The anode (+) on the left, where oxidation occurs
- The cathode (-) on the right, where reduction takes place
- The electrolyte solution between the electrodes
- The external power source providing electrical energy
- The flow of electrons from the power source through the external circuit
- The movement of ions through the electrolyte
The operation of an electrolytic cell begins when the power source is connected, creating a potential difference between the electrodes. This causes electrons to flow from the negative terminal of the power source to the cathode. At the cathode, these electrons participate in reduction reactions, typically reducing cations in the electrolyte. Simultaneously, at the anode, oxidation occurs, often involving the removal of electrons from anions in the electrolyte or from the anode material itself.
The electrolyte plays a crucial role by allowing ions to move between the electrodes, completing the electrical circuit within the cell. This ionic movement is essential for maintaining charge balance and enabling the continuous flow of current through the system.
One of the most well-known applications of electrolytic cells is in the process of electrolysis, used for various purposes such as:
- Purification of metals
- Production of chemicals like chlorine and sodium hydroxide
- Electroplating of metals
- Water treatment and purification
The ability of electrolytic cells to drive non-spontaneous reactions makes them invaluable in many industrial processes. By controlling the applied voltage and current, it's possible to selectively produce desired chemical products or purify materials that would otherwise be difficult or impossible to obtain through spontaneous reactions.
In conclusion, electrolytic cells are fundamental to many modern industrial processes and scientific applications. Their unique ability to convert electrical energy into chemical energy, driving non-spontaneous reactions, sets them apart from galvanic cells. Understanding the components and principles of electrolytic cells is crucial for anyone studying electrochemistry or working in related fields. As technology advances, the applications of electrolytic
The Process of Electrolysis
Electrolysis is a fundamental electrochemical process that uses electrical energy to drive non-spontaneous chemical reactions. This process plays a crucial role in various industrial applications, from metal extraction to water purification. Understanding the electrolysis process step by step can provide valuable insights into its mechanisms and applications.
The electrolysis process begins with the setup of an electrolytic cell. This cell consists of two electrodes (an anode and a cathode) immersed in an electrolyte solution or molten ionic compound. The electrodes are connected to an external power source, typically a battery or DC power supply, which provides the necessary electrical energy to drive the reaction.
When an electric current is applied to the cell, it creates a potential difference between the electrodes. This applied voltage causes the movement of ions within the electrolyte. Positively charged ions (cations) are attracted to the negatively charged cathode, while negatively charged ions (anions) move towards the positively charged anode. This movement of ions is essential for the electrolysis reaction to occur.
At the electrodes, oxidation and reduction reactions take place simultaneously. Oxidation occurs at the anode, where electrons are removed from the species present. This process results in the formation of more positively charged ions or neutral molecules. Reduction, on the other hand, happens at the cathode, where electrons are added to the species present, forming more negatively charged ions or neutral molecules.
The concepts of oxidation and reduction are central to understanding electrolysis. Oxidation involves the loss of electrons, while reduction involves the gain of electrons. In the context of electrolysis, the anode is the site of oxidation, often referred to as the positive electrode. The cathode, being the site of reduction, is considered the negative electrode. These processes are summarized by the mnemonic "OIL RIG" - Oxidation Is Loss, Reduction Is Gain (of electrons).
To illustrate the electrolysis process, let's consider the example of the electrolysis of molten sodium chloride (NaCl). In this case, the electrolyte is molten NaCl, which contains Na+ and Cl- ions. When an electric current is applied, the following reactions occur:
At the cathode (reduction): Na+ + e- Na
At the anode (oxidation): 2Cl- Cl2 + 2e-
The overall electrolysis equation for molten NaCl can be written as:
2NaCl (l) 2Na (l) + Cl2 (g)
In this equation, solid sodium metal is produced at the cathode, while chlorine gas is evolved at the anode. The (l) and (g) notations indicate the liquid and gaseous states of the products, respectively. This reaction demonstrates how electrolysis can be used to decompose a compound into its constituent elements.
It's important to note that the electrolysis of aqueous NaCl would yield different products due to the presence of water molecules. In an aqueous solution, water would be preferentially reduced at the cathode, producing hydrogen gas instead of sodium metal.
The electrolysis process has numerous practical applications. In industry, it is used for the production of metals like aluminum and the purification of copper. Electrolysis is also employed in the production of chlorine and sodium hydroxide, essential chemicals in many manufacturing processes. In everyday life, electrolysis is used in electroplating, where a thin layer of metal is deposited on another surface for decorative or protective purposes.
Understanding the electrolysis process and its underlying principles is crucial for many fields, including chemistry, materials science, and engineering. By manipulating the conditions of electrolysis, such as the applied voltage, electrode materials, and electrolyte composition, scientists and engineers can control the products formed and optimize the efficiency of the process for various applications.
In conclusion, electrolysis is a powerful technique that harnesses electrical energy to drive chemical reactions. Through the movement of ions and the occurrence of oxidation and reduction reactions at the electrodes, electrolysis enables
Factors Affecting Electrolysis
Electrolysis is a crucial process in various industrial applications and scientific research. Understanding the factors that influence this electrochemical reaction is essential for optimizing its efficiency and controlling the products formed. Several key factors affect the electrolysis process, including the nature of the electrolyte, electrolyte concentration, temperature, and applied voltage.
The nature of the electrolyte plays a significant role in determining the outcome of electrolysis. Different electrolytes contain various ions, which affect the products formed at the electrodes. For instance, when electrolyzing an aqueous solution of sodium chloride, water molecules are preferentially decomposed due to their lower decomposition potential compared to sodium and chloride ions. In contrast, electrolysis of molten sodium chloride results in the production of sodium metal and chlorine gas.
Electrolyte concentration is another crucial factor affecting the electrolysis process. Higher concentrations of electrolytes generally lead to increased conductivity, allowing for a greater flow of electric current. This enhanced conductivity results in a faster rate of electrolysis and more efficient production of desired products. However, it's important to note that extremely high concentrations can sometimes hinder the process due to increased viscosity or reduced ion mobility.
Temperature is a critical factor that influences the rate of electrolysis and the nature of products formed. Elevated temperatures typically increase the rate of electrolysis by enhancing ion mobility and reducing the viscosity of the electrolyte. This increased mobility allows ions to move more freely between electrodes, facilitating faster reactions. Additionally, higher temperatures can affect the solubility of gases produced during electrolysis, potentially altering the composition of the final products.
The applied voltage is perhaps one of the most significant factors affecting electrolysis. The voltage determines the amount of electrical energy available to drive the electrochemical reactions. A higher applied voltage generally results in a faster rate of electrolysis, as it provides more energy to overcome the activation energy barrier of the reactions. However, it's crucial to apply the correct voltage to achieve the desired products and avoid unwanted side reactions.
When considering the electrolysis of molten sodium chloride, the necessary voltage to power the process is an important consideration. The minimum voltage required for this electrolysis is determined by the standard electrode potentials of the half-reactions involved. For molten sodium chloride, the reduction of sodium ions to sodium metal occurs at the cathode, while the oxidation of chloride ions to chlorine gas takes place at the anode.
The standard reduction potential for sodium is -2.71 V, while the standard reduction potential for chlorine is -1.36 V. To calculate the minimum voltage required, we need to consider the difference between these potentials: 2.71 V - (-1.36 V) = 4.07 V. This means that a minimum voltage of 4.07 V is theoretically required to drive the electrolysis of molten sodium chloride.
However, in practical applications, a higher voltage is typically applied to overcome various resistances in the system and ensure a reasonable rate of electrolysis. Factors such as electrode overpotential, solution resistance, and bubble formation can increase the required voltage. Industrial processes often use voltages in the range of 6-7 V to maintain efficient electrolysis of molten sodium chloride.
Understanding these factors and their interplay is crucial for optimizing electrolysis processes in various applications. By carefully controlling the nature and concentration of the electrolyte, adjusting the temperature, and applying the appropriate voltage, it's possible to enhance the efficiency of electrolysis and selectively produce desired products. This knowledge is invaluable in industries such as metal production, water treatment, and chemical manufacturing, where electrolysis plays a vital role in numerous processes.
Applications of Electrolysis
Electrolysis, a process that uses electrical energy to drive non-spontaneous chemical reactions, has numerous applications in industry and everyday life. This versatile technique plays a crucial role in various sectors, from metal production to chemical manufacturing. Let's explore some of the most significant applications of electrolysis and how they impact our world.
One of the most common applications of electrolysis is electroplating. This process involves coating a metal object with a thin layer of another metal to enhance its appearance, durability, or conductivity. For example, jewelry makers use electroplating to coat less expensive metals with gold or silver, creating affordable yet attractive pieces. In the automotive industry, chrome plating is applied to various parts to improve corrosion resistance and aesthetic appeal. Electronics manufacturers also utilize electroplating to create conductive layers on circuit boards and components.
Another important application of electrolysis is the purification of metals. This process, known as electrorefining, is used to produce high-purity metals from less pure raw materials. For instance, copper extracted from ores often contains impurities. By using electrolysis, manufacturers can produce 99.99% pure copper, which is essential for electrical applications. Similarly, aluminum undergoes electrolytic purification to meet the high standards required for aerospace and other advanced industries.
The production of chemicals is another significant area where electrolysis plays a vital role. One of the most important examples is the chlor-alkali process, which produces chlorine, sodium hydroxide (caustic soda), and hydrogen. These chemicals are fundamental to many industries, including water treatment, paper production, and manufacturing of plastics and textiles. The electrolysis of brine (saltwater) in this process efficiently separates these valuable components, making it a cornerstone of the chemical industry.
Electrolysis is also crucial in the extraction of reactive metals from their ores, particularly for metals that cannot be extracted by chemical reduction alone. Aluminum, for example, is produced through the Hall-Héroult process, which uses electrolysis to extract the metal from alumina (aluminum oxide). This process is energy-intensive but remains the primary method for producing this lightweight, versatile metal used in everything from beverage cans to aircraft components.
In everyday life, we encounter products of electrolysis more often than we might realize. For instance, swimming pools use electrolytic chlorine generators to produce chlorine directly from salt water, eliminating the need for handling and storing hazardous chlorine chemicals. In the realm of personal care, some water flossers use electrolysis to create oxygenated water, which can be more effective in removing plaque and bacteria.
The applications of electrolysis extend to environmental protection as well. Electrolytic water treatment systems are used to remove heavy metals and other contaminants from industrial wastewater, helping to reduce pollution and recycle water. In the field of renewable energy, electrolysis is being explored as a method to produce hydrogen fuel from water, potentially offering a clean energy storage solution to complement intermittent renewable sources like solar and wind power.
As technology advances, new applications of electrolysis continue to emerge. For example, researchers are investigating electrolytic processes for carbon dioxide reduction, which could help mitigate climate change by converting CO2 into useful chemicals or fuels. The versatility and efficiency of electrolysis ensure its continued importance in industry and innovation, driving progress in fields ranging from materials science to environmental technology.
Quantitative Aspects of Electrolysis
Faraday's laws of electrolysis, formulated by Michael Faraday in 1834, are fundamental principles that govern the quantitative aspects of electrolysis. These laws provide a crucial link between the amount of electrical charge passed through an electrolyte and the quantity of substance produced or consumed at the electrodes. Understanding and applying Faraday's laws is essential for precise calculations in electrochemistry and various industrial processes.
Faraday's First Law of Electrolysis states that the mass of a substance produced or consumed at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte. The Second Law states that when the same quantity of electricity is passed through different electrolytes, the masses of substances deposited or liberated at the electrodes are proportional to their chemical equivalent weights.
The significance of Faraday's laws lies in their ability to predict and quantify the outcomes of electrolysis reactions accurately. They form the basis for calculating the amount of product formed, the time required for a specific reaction, or the current needed to produce a desired quantity of substance. These laws have wide-ranging applications in electroplating, electrorefining, and the production of various chemicals and metals.
The electrolysis formula derived from Faraday's laws is a powerful tool for quantitative analysis. It is expressed as:
m = (Q × M) / (n × F)
Where:
- m is the mass of the substance produced or consumed (in grams)
- Q is the total electric charge passed through the solution (in coulombs)
- M is the molar mass of the substance (in g/mol)
- n is the number of electrons transferred per ion in the reaction
- F is Faraday's constant (96,485 C/mol)
To use this formula, one must first identify the half-reaction occurring at the electrode of interest and determine the number of electrons involved. Then, gather the necessary information such as the current, time, and molar mass of the substance. The charge Q can be calculated by multiplying the current (in amperes) by the time (in seconds).
Let's consider an example calculation to demonstrate the application of the electrolysis formula. Suppose we want to determine the mass of copper deposited on an electrode when a current of 2.5 amperes is passed through a copper sulfate solution for 30 minutes.
Given:
- Current (I) = 2.5 A
- Time (t) = 30 minutes = 1800 seconds
- Molar mass of copper (M) = 63.55 g/mol
- Number of electrons transferred per copper ion (n) = 2
Step 1: Calculate the charge (Q) passed through the solution
Q = I × t = 2.5 A × 1800 s = 4500 C
Step 2: Apply the electrolysis formula
m = (Q × M) / (n × F) = (4500 C × 63.55 g/mol) / (2 × 96,485 C/mol) = 1.48 g
Therefore, 1.48 grams of copper will be deposited on the electrode under these conditions.
This example illustrates how Faraday's laws and the electrolysis formula can be used to make precise predictions in electrochemical processes. By manipulating the formula, one can also calculate other variables such as the time required for a specific deposition or the current needed to produce a certain mass of substance. These calculations are invaluable in optimizing industrial electrolysis processes and in forensic applications where trace metal analysis is crucial.
Conclusion
In this article, we've explored the fascinating world of electrolysis, a crucial process in chemistry with numerous practical applications. We've discussed the key components of an electrolytic cell, including electrodes, electrolytes, and power sources. The importance of understanding electrolysis cannot be overstated, as it plays a vital role in various industries, from metal purification to water treatment. Remember to watch the introductory video for a visual demonstration of these concepts, which can greatly enhance your understanding. For those eager to delve deeper into electrolysis and its applications, we recommend exploring additional resources such as chemistry textbooks, online courses, and scientific journals. By mastering the principles of electrolysis, you'll gain valuable insights into chemical reactions and its impact on our daily lives. Whether you're a student, professional, or simply curious about science, electrolysis offers a fascinating glimpse into the power of chemical reactions and their practical applications.
Electrolysis: Definition and Explanation
What is electrolysis? Definition of electrolysis.
Step 1: Introduction to Electrolysis
Electrolysis is a process that involves the breaking up of substances using an electric current. This process is essential in various chemical reactions and industrial applications. Unlike redox reactions, where substances spontaneously react to produce electricity, electrolysis requires an external source of electrical energy to drive non-spontaneous reactions.
Step 2: Understanding the Term "Electrolysis"
The term "electrolysis" is derived from two words: "electro," referring to electricity, and "lysis," which means breaking up. In biology, lysis refers to the breaking up of cells. Similarly, in electrolysis, it refers to the breaking up of substances using electricity. This process is crucial for decomposing compounds into their constituent elements or simpler compounds.
Step 3: Comparing Electrolysis to Redox Reactions
In redox reactions, substances spontaneously react to produce electricity. However, electrolysis is the opposite. It involves non-spontaneous processes that require an external source of electricity to break up substances. This means that energy must be supplied to drive the reaction towards the products, as opposed to redox reactions where energy is released.
Step 4: Examples of Electrolysis Reactions
One common example of electrolysis is the decomposition of sodium chloride (NaCl) in an aqueous solution to produce chlorine gas (Cl2) and sodium metal (Na). Sodium chloride is a stable compound, and breaking it down into its elements requires significant energy input. Another example is the electrolysis of aluminum oxide (Al2O3) to produce aluminum metal. Aluminum is highly reactive and is not found in its elemental form in nature. Electrolysis is used to extract aluminum from its oxide form.
Step 5: The Electrolytic Cell
An electrolytic cell is used to carry out electrolysis. It consists of a power source, electrodes, and an electrolyte. The power source provides the necessary electrical energy to drive the reaction. The electrodes, usually made of inert materials like platinum or carbon, conduct the electricity. The electrolyte is the medium that allows the flow of ions. In a binary salt electrolysis, where only two ions are involved, a single-phase system is used without the need for a salt bridge.
Step 6: Operating Conditions and Cell Potential
To find the operating conditions for electrolysis, it is essential to understand the cell potential. The cell potential indicates the amount of electrical energy required to drive the reaction. For non-spontaneous reactions, the cell potential is positive, meaning energy must be supplied to the system. Understanding the cell potential helps in determining the efficiency and feasibility of the electrolysis process.
Step 7: Faraday's Laws of Electrolysis
Faraday's laws, formulated by Michael Faraday, relate the amount of substance produced or consumed at an electrode to the quantity of electric charge passed through the electrolyte. These laws are fundamental in quantifying the relationship between the electric current and the chemical reaction. By measuring the current using an ammeter, it is possible to determine the stoichiometry of the reaction and the amount of substance involved.
Step 8: Practical Applications of Electrolysis
Electrolysis has numerous practical applications, including the extraction of metals, electroplating, and the production of chemicals. For example, electrolysis is used to extract aluminum from bauxite ore, produce chlorine gas for industrial use, and electroplate objects with a thin layer of metal for corrosion resistance and aesthetic purposes.
Step 9: Conclusion
In summary, electrolysis is a vital process in chemistry and industry, involving the breaking up of substances using an electric current. It requires an external source of electrical energy to drive non-spontaneous reactions. Understanding the principles of electrolysis, including the role of the electrolytic cell, cell potential, and Faraday's laws, is essential for its practical applications.
FAQs
Here are some frequently asked questions about electrolysis:
1. What is the process of conversion in electrolysis?
In electrolysis, electrical energy is converted into chemical energy. The external power source provides electrical energy, which drives non-spontaneous chemical reactions at the electrodes, resulting in the formation of new chemical substances.
2. What is the chemical equation for electrolysis?
The chemical equation for electrolysis depends on the specific substances involved. For example, in the electrolysis of molten sodium chloride (NaCl), the equation is: 2NaCl (l) 2Na (l) + Cl2 (g). This shows the decomposition of NaCl into sodium metal and chlorine gas.
3. How is electrolysis used in everyday life?
Electrolysis has numerous everyday applications, including: - Electroplating jewelry and automotive parts - Purifying metals like copper and aluminum - Producing chlorine for water treatment - Extracting reactive metals like sodium and magnesium - Charging rechargeable batteries
4. What is the formula for calculating electrolysis?
The main formula used in electrolysis calculations is derived from Faraday's laws: m = (Q × M) / (n × F), where m is the mass of substance produced, Q is the total electric charge passed, M is the molar mass of the substance, n is the number of electrons transferred per ion, and F is Faraday's constant (96,485 C/mol).
5. What is the necessary voltage to power the electrolysis of molten sodium chloride?
The theoretical minimum voltage required for the electrolysis of molten sodium chloride is about 4.07 V. This is calculated from the difference in standard reduction potentials of sodium and chlorine. However, in practice, a higher voltage (typically 6-7 V) is used to overcome various resistances in the system and ensure efficient electrolysis.
Prerequisite Topics
Understanding the fundamental concepts that lay the groundwork for more advanced topics is crucial in any field of study, especially in chemistry. When it comes to the topic of electrolysis, having a solid grasp of prerequisite subjects is essential for comprehending its intricacies and applications. One of the most important prerequisite topics for understanding electrolysis is the introduction to chemical reactions.
Electrolysis is a complex process that involves the use of electrical energy to drive non-spontaneous chemical reactions. To fully appreciate the principles behind electrolysis, students must first have a strong foundation in basic chemical reactions. This prerequisite knowledge enables them to understand how electrical energy can be harnessed to facilitate chemical changes that wouldn't occur naturally.
The study of non-spontaneous chemical reactions is particularly relevant to electrolysis. In nature, most chemical reactions occur spontaneously, moving towards a state of lower energy. However, electrolysis deals with reactions that require an input of energy to proceed. By understanding the concepts of spontaneous and non-spontaneous reactions, students can better grasp why electrolysis is necessary and how it functions.
Moreover, the ability to write and balance chemical equations is a crucial skill when studying electrolysis. This fundamental aspect of chemistry, covered in the introduction to chemical reactions, allows students to represent the processes occurring during electrolysis accurately. It helps in visualizing the movement of ions, the transfer of electrons, and the formation of products at the electrodes.
Understanding redox reactions, which are also part of the introduction to chemical reactions, is another vital prerequisite for electrolysis. Electrolysis involves the oxidation and reduction of species at the electrodes, and a solid understanding of these concepts is essential for comprehending the electron transfer processes that occur during electrolysis.
By mastering these prerequisite topics, students build a strong foundation that enables them to tackle the complexities of electrolysis with confidence. They can more easily understand the role of the electrolyte, the function of electrodes, and the principles governing the decomposition of compounds through electrical means. This knowledge not only aids in academic understanding but also prepares students for practical applications of electrolysis in various fields, including metallurgy, electroplating, and the production of chemicals.
In conclusion, the importance of understanding prerequisite topics like the introduction to chemical reactions cannot be overstated when studying electrolysis. It provides the necessary context and foundational knowledge that allows students to fully appreciate and engage with this fascinating aspect of electrochemistry. By building on these fundamental concepts, students can develop a deeper, more comprehensive understanding of electrolysis and its wide-ranging applications in science and industry.
In this lesson, we will learn:
- To recall the definition of electrolysis and some examples of its use.
- To draw an electrolytic cell used in electrolysis.
- How to find electrolysis operating conditions using cell potential calculations.
- How to use Faraday’s law to relate charge flow to reaction stoichiometry.
Notes:
- Electrolysis is the breaking up (-lysis) of substances using electric current (electro-) – hence the name.
It can be done with dissolved or molten ionic solutions such as sodium chloride (NaCl) or aluminum oxide (Al2O3) to obtain the pure elements from their ions. This is very useful because some valuable substances like aluminum metal are not found naturally in their elemental form because they are too reactive. They must be produced from their ions.
In terms of energy, electrolysis reactions are the opposite of a redox cell reaction:- In a redox or galvanic cell, the reaction is spontaneous and chemical potential energy is used to produce electrical energy. In terms of thermodynamics, Gibbs free energy of the system drives the forward reaction to make products.
- In an electrolysis reaction, the process is not spontaneous, so electrical energy is needed for the reaction to make products. In terms of thermodynamics, the forward reaction results in a loss of Gibbs free energy of the surroundings, so it is unfavourable.
- Electrolysis is often used on binary salts – containing only two different ions – in order to obtain the constituent ions as their pure elements. This is how NaCl is electrolyzed to obtain chlorine gas (Cl2) and sodium which reacts to form NaOH, another important chemical.
Because only two oppositely charged ions are involved with binary salts, both ions will take part in the opposite processes of reduction and oxidation to return to their elemental form. When binary salts are electrolyzed, there are a few differences to the cell compared to redox (galvanic) cells:- Because the compounds being electrolyzed are water-soluble, inert electrodes are used, normally carbon or platinum.
- No salt bridge is necessary. This is because electrolyte can mix as there’s usually only one electrolyte; the ionic substance that is being electrolyzed! It is in aqueous or molten form so a current can flow easily.
See the diagram below:
- As said above, electrolysis reactions are not spontaneous!
This means that instead of the cells producing electricity, electrolysis cells need voltage to be applied. You can calculate the voltage just like you calculate redox cell potential in Calculating cell potential. Your calculated E0cell should be less than zero for an electrolytic cell.
For example, with sodium chloride, NaCl, being electrolyzed to produce sodium and chlorine:
Na+ + e-→Na E0red = -2.71 V
Cl2 + 2e-→2Cl- E0red = +1.36 V
NaCl electrolysis reduces Na+ to Na metal and oxidizes Cl- ions to Cl2 so the Cl half-equation and E0red value needs reversing. Applying this to the cell potential equation:E0cell = E0red (reduction) + E0ox (oxidation)
We can work out:E0cell = -2.71 + (-1.36) = -2.71 -1.36 = -4.07 V
This means that at least 4.07 V must be applied to the cell for the reaction to start occurring spontaneously. Higher voltage is applied usually because there is resistance in the cell and it is not at standard conditions. - Remember that electrolysis in aqueous solution means in water – the water solvent can be reduced or oxidized itself. Taking the example of NaCl again, in aqueous solution there are four possible reactions going on in the cell, each with different reduction potentials.1 Assuming neutral pH:
- The oxidation of chlorine: 2Cl-→Cl2 + 2e- where E0ox = -1.36 V
- The oxidation of water: 2H2O (l) → O2 + 4e- + 4 H+(aq) where E0ox = -1.23 V
- The reduction of water: 2H2O (l) + 2e-→H2 (g) + 2OH- (aq) where E0red = -0.83 V
- The reduction of sodium: Na+(aq) + e-→Na where E0red = -2.71 V
If we applied 4.07 V to the system, what would react? The reaction that is preferable is the one with the least required voltage! Like a spontaneous redox reaction, the half-equations with the greatest E0red and E0ox are preferred.
‘Preferred’ means this reaction is much more likely to happen. The other reaction(s) can occur but are less likely to.
One reduction and one oxidation ‘half reaction’ must happen. Think about the voltage differences (potential difference!) between the half-equations:- The reduction potential of water is less negative than sodium metal, so only water gets reduced to hydrogen gas at the cathode (where reduction occurs).
- The oxidation potential for water is less negative than chlorine but they are very similar.
The real process of NaCl electrolysis is this:2 NaCl (aq) + 2H2O(l)→2Na+ (aq) + 2OH- (aq) + H2 (g) + Cl2 (g)
Chlorine gas (Cl2) is obtained, as is hydrogen gas and NaOH which can both be isolated and sold. - You need to look at the conditions that occur in your electrolytic cell. When you have a substance to be electrolyzed, check the ions that will be produced:
- If H+ is produced, your conditions will be acidic.
- If OH- is made, it will be basic.
- If neither, assume it is neutral.
2OH- + Br-→H2O + BrO- + 2e- E0 = +0.76 V
This half-reaction could only occur in basic conditions. - As electrical energy is either used up, as in electrolytic cells, or generated in voltaic cells, we can measure the flow of electrons (the current) using an ammeter.
Faraday’s law relates this current to the stoichiometry and rate of a reaction: - = charge (in coulombs)
- = current (in amps)
- = time (seconds)
- In other words, the higher the current, the more charge per unit time passes through the cell and the faster a reaction at an electrode will occur (e.g. more Cu (s) being deposited, gaining mass, or a Ni (s) electrode being oxidised, losing mass).
- The mass change of a copper electrode during an experiment;
- The number of moles of copper this mass equates to;
- The number of moles of electrons liberated;
- The amount of charge this amount of electrons carries
- The time taken or the current reading of the cell.
- = mass change of substance at electrode (in g)
- = electric charge (Coulombs)
- = molar mass of substance at electrode
- = the Faraday constant, 96485 C mol-1
- = the number of electrons transferred per ion (e.g. 2 for Ca2+ ions).
- WORKED EXAMPLE: Using Faraday’s laws.
Calculate how long a 7A current must be applied to deposit 1.5 g of Zn metal at an electrode.
We are solving for time, so rearranging this equation for t gives us this:
Where:
This equation is important because when rearranged for , we measure the current as a rate of charge over time. Quantity of charge is proportional to quantity of mass change at an electrode.
This is where the value of the Faraday constant comes from – 96485 C mol-1. It means that one mole of electrons carries 96485 Coulombs of charge. For example, if a copper electrode was being oxidised in the half-reaction:
We know that one mole of copper liberates two moles of electrons, which has 2 × 96485 C of charge. Since moles are related to the mass of a sample substance, we now have a relationship between:
Faraday’s laws of electrochemistry can be summarized by the following:
Where:
As the equation above shows, increasing q will increase the rate of deposition (Mn+ (aq) ions becoming M (s) at the electrode) or liberation, the reverse process.
Substituting q for in the equation above gives:
This relates all the factors mentioned above.
We have the values m = 1.5; F = 96485; z = 2 (for 2e- per Zn2+ ion); I = 7; M = 65.38.
With so many variables, a good way to ‘proof’ your answer is to solve for the units. We are solving for t, so our units should end in s (seconds).
The calculation shows that it takes ten and a half minutes for this much of a mass change to take place. Alternatively, you can use the conversion factor method to go from mass to time:
Whichever method you use, we received the same answer because the units cancel!
1 Sourcefor data: ATKINS, P. W., & DE PAULA, J. (2006).Atkins' Physical chemistry. Oxford, Oxford University Press