Textbook Question
Predict the anode, cathode, and overall cell reactions when an aqueous solution of each of the following salts is electrolyzed in a cell having inert electrodes. (c) LiOH
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Ch.19 - Electrochemistry
McMurry8th EditionChemistryISBN: 9781292336145
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Table of contents
- Ch.1 - Chemical Tools: Experimentation & Measurement143
- Ch.2 - Atoms, Molecules & Ions134
- Ch.3 - Mass Relationships in Chemical Reactions149
- Ch.4 - Reactions in Aqueous Solution157
- Ch.5 - Periodicity & Electronic Structure of Atoms92
- Ch.6 - Ionic Compounds: Periodic Trends and Bonding Theory73
- Ch.7 - Covalent Bonding and Electron-Dot Structures47
- Ch.8 - Covalent Compounds: Bonding Theories and Molecular Structure51
- Ch.9 - Thermochemistry: Chemical Energy104
- Ch.10 - Gases: Their Properties & Behavior138
- Ch.11 - Liquids & Phase Changes43
- Ch.12 - Solids and Solid-State Materials56
- Ch.13 - Solutions & Their Properties98
- Ch.14 - Chemical Kinetics79
- Ch.15 - Chemical Equilibrium107
- Ch.16 - Aqueous Equilibria: Acids & Bases94
- Ch.17 - Applications of Aqueous Equilibria125
- Ch.18 - Thermodynamics: Entropy, Free Energy & Equilibrium66
- Ch.19 - Electrochemistry130
- Ch.20 - Nuclear Chemistry73
- Ch.21 - Transition Elements and Coordination Chemistry122
- Ch.22 - The Main Group Elements155
- Ch.23 - Organic and Biological Chemistry43
Chapter 19, Problem 141
What constant current (in amperes) is required to produce aluminum by the Hall–Héroult process at a rate of 40.0 kg/h?
Verified step by step guidance1
Identify the chemical reaction involved in the Hall–Héroult process: \(2\text{Al}_2\text{O}_3 + 3\text{C} \rightarrow 4\text{Al} + 3\text{CO}_2\).
Determine the molar mass of aluminum (Al), which is approximately 26.98 g/mol.
Convert the production rate of aluminum from kg/h to moles per second: \(40.0 \text{ kg/h} \rightarrow \text{g/s} \rightarrow \text{mol/s}\).
Use Faraday's laws of electrolysis to relate the moles of aluminum produced to the charge required: \(n = \frac{Q}{zF}\), where \(n\) is moles of electrons, \(Q\) is charge in coulombs, \(z\) is the number of electrons transferred per mole of aluminum (3 for Al), and \(F\) is Faraday's constant (96485 C/mol).
Calculate the current \(I\) using the formula \(I = \frac{Q}{t}\), where \(t\) is time in seconds.
Key Concepts
Here are the essential concepts you must grasp in order to answer the question correctly.
Hall–Héroult Process
The Hall–Héroult process is an industrial method for extracting aluminum from its ore, bauxite. It involves the electrolysis of aluminum oxide dissolved in molten cryolite, where aluminum ions are reduced to form aluminum metal at the cathode. Understanding this process is crucial for calculating the current required, as it directly relates to the amount of aluminum produced and the efficiency of the electrolysis.
Faraday's Laws of Electrolysis
Faraday's Laws of Electrolysis describe the relationship between the amount of substance produced at an electrode during electrolysis and the quantity of electric charge passed through the electrolyte. The first law states that the mass of a substance altered at an electrode is proportional to the total electric charge passed. This principle is essential for determining the current needed to produce a specific mass of aluminum in the Hall–Héroult process.
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Faraday's Constant in Electrochemistry
Current Calculation
To calculate the current required for electrolysis, one must consider the molar mass of aluminum and the charge needed to reduce aluminum ions. The formula I = (n * F) / t relates current (I) to the number of moles (n), Faraday's constant (F), and time (t). This calculation is vital for determining the amperage necessary to achieve the desired production rate of aluminum, ensuring efficient operation of the electrolysis process.
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Related Practice
Textbook Question
Predict the anode, cathode, and overall cell reactions when an aqueous solution of each of the following salts is electrolyzed in a cell having inert electrodes. (a) Ag2SO4
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Textbook Question
In order to charge a lead storage battery (Section 19.10) 500.0 g of PbSO4(s) must be converted into PbO2(s) and Pb(s). (a) Does the reaction represent an electrolytic or galvanic cell?
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Textbook Question
How many hours are required to produce 1.00 * 103 kg of sodium by the electrolysis of molten NaCl with a constant current of 3.00 * 104 A? How many liters of Cl2 at STP will be obtained as a by-product?
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Textbook Question
Electrolysis of a metal nitrate solution M(NO3)2(aq) for 325 min with a constant current of 20.0 A gives 111 g of the metal. Identify the metal ion M2+.
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Textbook Question
In order to charge a lead storage battery (Section 19.10) 500.0 g of PbSO4(s) must be converted into PbO2(s) and Pb(s). (b) How many coulombs of electrical charge are needed?
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