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Questions and Answers
Which of the following terms did Michael Faraday introduce to the field of electrolysis?
Which of the following terms did Michael Faraday introduce to the field of electrolysis?
- Anode, cathode, and salt bridge
- Ion, electrode, and electrolyte (correct)
- Acid, base, and pH
- Oxidation, reduction, and redox
According to Faraday's first law of electrolysis, what is the relationship between the amount of chemical change produced and the quantity of electricity passed?
According to Faraday's first law of electrolysis, what is the relationship between the amount of chemical change produced and the quantity of electricity passed?
- Inversely proportional
- Logarithmically proportional
- Exponentially proportional
- Directly proportional (correct)
If a current of 2 amps is passed through an electrolytic cell for 30 minutes, what is the total charge, in coulombs, that passes through the cell?
If a current of 2 amps is passed through an electrolytic cell for 30 minutes, what is the total charge, in coulombs, that passes through the cell?
- 60 C
- 3600 C (correct)
- 900 C
- 1800 C
What is the significance of the Faraday constant in the context of electrolysis?
What is the significance of the Faraday constant in the context of electrolysis?
According to Faraday's second law of electrolysis, what is the relationship between the quantities of substances deposited at an electrode and their chemical equivalent weights?
According to Faraday's second law of electrolysis, what is the relationship between the quantities of substances deposited at an electrode and their chemical equivalent weights?
In the electrolysis of molten NaCl, if 2 moles of electrons are passed through the electrolytic cell, how many moles of sodium metal will be produced?
In the electrolysis of molten NaCl, if 2 moles of electrons are passed through the electrolytic cell, how many moles of sodium metal will be produced?
If electrolysis is used to plate a metal item with silver, and a current of 5.0 A runs for 30.0 minutes, what additional information is needed to calculate the mass of silver deposited?
If electrolysis is used to plate a metal item with silver, and a current of 5.0 A runs for 30.0 minutes, what additional information is needed to calculate the mass of silver deposited?
During the electrolysis of water, hydrogen and oxygen gas are produced. If 0.5 moles of oxygen are generated, how many moles of electrons were transferred in the process?
During the electrolysis of water, hydrogen and oxygen gas are produced. If 0.5 moles of oxygen are generated, how many moles of electrons were transferred in the process?
Why is it important to balance ionic equations before applying Faraday’s Laws in electrolysis calculations?
Why is it important to balance ionic equations before applying Faraday’s Laws in electrolysis calculations?
In the context of electrochemistry, what does one Faraday equal?
In the context of electrochemistry, what does one Faraday equal?
If a battery operates with a current of 0.2 A for 2 hours, what is the total charge delivered by the battery?
If a battery operates with a current of 0.2 A for 2 hours, what is the total charge delivered by the battery?
What additional information, besides current and time, is needed to determine the mass of a substance produced during electrolysis?
What additional information, besides current and time, is needed to determine the mass of a substance produced during electrolysis?
Why is understanding stoichiometry important when applying Faraday's laws?
Why is understanding stoichiometry important when applying Faraday's laws?
In an electrolytic process, if it is desired to increase the mass of the substance deposited at the cathode, how should the current be adjusted, assuming the time remains constant?
In an electrolytic process, if it is desired to increase the mass of the substance deposited at the cathode, how should the current be adjusted, assuming the time remains constant?
Calculate the mass of copper deposited at the cathode when a constant current of 10A is passed through a solution of copper(II) sulfate ($CuSO_4$) for 60 minutes. (Molar mass of Cu = 63.5 g/mol)
Calculate the mass of copper deposited at the cathode when a constant current of 10A is passed through a solution of copper(II) sulfate ($CuSO_4$) for 60 minutes. (Molar mass of Cu = 63.5 g/mol)
If 2.0 grams of hydrogen fuel are supplied to a fuel cell, what is the theoretical maximum number of grams of oxygen that would be consumed, assuming complete reaction to form water?
If 2.0 grams of hydrogen fuel are supplied to a fuel cell, what is the theoretical maximum number of grams of oxygen that would be consumed, assuming complete reaction to form water?
A voltaic cell is set up using zinc and copper electrodes in their respective sulfate solutions. If the cell runs for 1 hour and produces a steady current of 0.5 A, calculate the mass of zinc consumed at the anode. (Molar mass of Zn = 65.4 g/mol)
A voltaic cell is set up using zinc and copper electrodes in their respective sulfate solutions. If the cell runs for 1 hour and produces a steady current of 0.5 A, calculate the mass of zinc consumed at the anode. (Molar mass of Zn = 65.4 g/mol)
What is the main reason why the actual yield of a product in electrolysis might be less than the theoretical yield calculated using Faraday's laws?
What is the main reason why the actual yield of a product in electrolysis might be less than the theoretical yield calculated using Faraday's laws?
In the electrolysis of $Al_2O_3$ to produce aluminum, if the process is only 80% efficient, what adjustment is required to calculate the actual amount of aluminum produced compared to the theoretical amount?
In the electrolysis of $Al_2O_3$ to produce aluminum, if the process is only 80% efficient, what adjustment is required to calculate the actual amount of aluminum produced compared to the theoretical amount?
Suppose you want to electroplate a metal spoon with silver. If you double both the current and the time used for electroplating, how will this affect the mass of silver deposited on the spoon?
Suppose you want to electroplate a metal spoon with silver. If you double both the current and the time used for electroplating, how will this affect the mass of silver deposited on the spoon?
Flashcards
Michael Faraday
Michael Faraday
A British chemist who studied electrolysis in the 1800s and explained how it worked, introducing key terms.
Electrolysis of Water
Electrolysis of Water
Decomposition of water into hydrogen and oxygen by passing an electric current through it.
Faraday's First Law
Faraday's First Law
The amount of chemical change produced is proportional to the quantity of electricity passed.
Charge Equation
Charge Equation
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Faraday Constant
Faraday Constant
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Faraday's Second Law
Faraday's Second Law
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Charge and Moles Equation
Charge and Moles Equation
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Faraday's Laws Use
Faraday's Laws Use
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AA Battery Current
AA Battery Current
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Battery Anode Reaction
Battery Anode Reaction
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Stoichiometry and Electrolysis
Stoichiometry and Electrolysis
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Study Notes
- Michael Faraday was a British chemist who studied electrolysis in the 1800s.
- Around 1820, chemists observed that passing an electric current through water decomposed it into hydrogen and oxygen gases, an early example of electrolysis.
- Faraday explained electrolysis and introduced terms like ion, electrode, electrolyte, and electrolysis.
- Faraday formulated two laws of electrolysis that are still used today to predict electrolysis outcomes.
Faraday’s First Law of Electrolysis
- States that the amount of chemical change produced by a current is proportional to the quantity of electricity passed (m α Q).
- Charge (Q) = current (I) × time (t), or Q = It.
- The unit of charge is coulombs (C), current is amps (A), and time is seconds (s).
- The charge on one mole of electrons is equal to one Faraday, which is a constant used in calculating the mass of a metal that is electroplated.
- 1 Faraday (F) = 96,500 coulombs/mol.
Faraday’s Second Law of Electrolysis
- States that the quantities of substances liberated or deposited at the electrode from a given quantity of electricity are proportional to the chemical equivalent weights of those substances.
- One Faraday equals the charge on one mole of electrons (6.02 × 1023 electrons).
- Q = n(e-) F, where Q is the quantity of electricity, n(e-) is the number of moles of electrons, and F is Faraday's constant (96,500 coulombs).
Example Calculation: Aluminium Production
- Aluminium is produced from alumina (Al2O3) by electrolysis, with the reaction: 2Al2O3(l) + 3C(s) → 4Al(l) + 3CO2(g).
- To calculate the mass of aluminium produced in 10 hours with a current of 160,000 Amps:
- Calculate the amount of charge: Q = I × t = 160,000 A × (10 hours × 60 minutes × 60 seconds) = 5.76 × 109 Coulombs.
- Calculate the moles of electrons: n = Q / F = (5.76 × 109 C) / (96,500 C/mol) = 59,689.1 mol.
- Determine the moles of aluminium produced: Al3+ (l) + 3e- → Al(s), so n(aluminium) = (1/3) × n(electrons) = (1/3) × 59,689.1 mol = 19,896.4 mol.
- Calculate the mass of aluminium produced: mass = mol × molar mass = 19,896.4 mol × 26.98 g/mol = 536.8 grams ≈ 5.4 × 102 grams (two significant figures).
Applications of Faraday's Laws and Stoichiometry
- Allows determination of the quantity of reactant consumed or product generated in an electrochemical reaction.
- Enables calculation of the current or time required to achieve a desired quantity of reactant or product by rearranging equations derived from Faraday's Laws.
AA Battery Example
- An alkaline AA battery produces a current of 0.40 A. The reaction occurring at the anode is: Zn(s) + 2OH−(aq) → ZnO(s) + H2O(l) + 2e−.
- To calculate the mass of zinc consumed if the battery operates for 50 minutes:
- Calculate the amount of charge: Q = I × t = 0.4 A × (50 × 60) s = 1200 C.
- Calculate the moles of electrons: n(e−) = Q / F = 1200 C / 96,500 C/mol = 0.0124 mol.
- Determine the moles of zinc: Zn2+(aq) + 2e− → Zn(s), so n(Zn) = (1/2) × n(e−) = (1/2) × 0.0124 mol = 0.00621 mol.
- Calculate the mass of zinc: m = n × M = 0.00621 mol × 65.4 g/mol = 0.41 g.
Fuel Cell Example
- A fuel cell has a hydrogen fuel supply of 3.80 g.
- To calculate the maximum time the cell can operate while delivering a 2.5 A current:
- Write the oxidation equation of hydrogen gas: H2(g) → 2H+(aq) + 2e−.
- Calculate the moles of hydrogen gas: n = m / M = 3.80 g / 2.0 g/mol = 1.9 mol.
- Calculate the moles of electrons: n(e−) = (2/1) × n(H2) = (2/1) × 1.9 mol = 3.8 mol.
- Calculate the time: t = (n × F) / I = (3.8 mol × 96,500 C/mol) / 2.5 A = 146,680 seconds.
- Convert to hours: t(hours) = 146,680 seconds / (60 × 60) = 40.7 hours.
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