Battery or cell converts chemical energy of the redox reaction to electrical energy. In fuel cells (a galvanic cell), the chemical energy of combustion of fuels like H2, ethanol, etc., is directly converted to electrical energy. In a fuel cell, H2 and O2 react to produce electricity, where H2 gas is oxidized at the anode and oxygen is reduced at the cathode, and the reactions involved are:
Anode reaction: H2 + 2OH- → 2H2O + 2e-
Cathode reaction: O2 + 2H2O + 4e- → 4OH-
67.2 L of H2 at STP reacts in 15 minutes.
To determine the moles of hydrogen oxidized, calculate the moles of hydrogen gas involved in the reaction. At Standard Temperature and Pressure (STP), one mole of any gas occupies 22.4 liters. Given 67.2 liters of H2 reacted, the moles are calculated as:
Number of moles of H2 = \(\frac{67.2 \ \text{L}}{22.4 \ \text{L/mol}}\)
The calculation yields:
= \(\frac{67.2}{22.4} = 3.0 \ \text{moles}\)
Thus, 3.0 moles of hydrogen were oxidized, consistent with the correct answer.
This query concerns the quantity of electrons generated by the oxidation of hydrogen (H2). The pertinent anode reaction in a fuel cell is:
H2 + 2OH- → 2H2O + 2e-
This equation indicates that one mole of H2 yields two moles of electrons (2e-).
The problem states that 67.2 L of H2 is consumed at Standard Temperature and Pressure (STP). At STP, one mole of any ideal gas occupies 22.4 L. Therefore, the moles of H2 can be calculated:
Moles of H2 = 67.2 L / 22.4 L/mol = 3 moles
Given that one mole of H2 produces two moles of electrons, three moles of H2 will produce:
3 moles H2 × 2 moles e-/mole H2 = 6 moles of electrons
Consequently, 6 moles of electrons are produced during the oxidation process.
The quantity of electricity (in Coulombs) generated from the oxidation of 67.2 L of H2 is calculated through the following steps:
To determine the silver deposited, we first analyze the electrochemical reaction's stoichiometry. The reduction of silver ions (Ag+) to silver metal (Ag) is represented by:
Ag+ + e- → Ag
This shows one mole of electrons is needed per mole of silver deposited. Silver's molar mass is approximately 108 g/mol.
Step 1: Calculate the total charge (Q) from the volume of hydrogen gas involved.
The provided volume of hydrogen gas is 67.2 L at STP, equivalent to 3 moles of H2 (since 1 mole of gas occupies 22.4 L at STP). The anode reaction releases 2 moles of electrons per mole of H2. Therefore, the total moles of electrons transferred is:
Total moles of electrons = 3 × 2 = 6
Using Faraday's constant (F ≈ 96485 C/mol), the charge (Q) passed is:
Q = 6 moles × 96485 C/mol = 578910 C
Step 2: Calculate the mass of silver deposited.
The number of moles of Ag deposited directly corresponds to the moles of electrons transferred, as the stoichiometry dictates a 1:1 ratio.
Mass of Ag = moles of Ag × Molar mass of Ag = 6 moles × 108 g/mol = 648 g
Consequently, the amount of silver deposited is 648 g.
The Apollo moon missions generated electrical power using H₂-O₂ fuel cells, an application grounded in electrochemistry. These fuel cells offered a dependable and efficient power source for the mission. The H₂-O₂ fuel cell operates as a galvanic cell. Hydrogen (H₂) undergoes oxidation at the anode, while oxygen (O₂) undergoes reduction at the cathode, producing electricity via these reactions:
Anode reaction: H2 + 2OH- → 2H2O + 2e-
Cathode reaction: O2 + 2H2O + 4e- → 4OH-
A volume of 68.2 L of H₂ at standard temperature and pressure (STP) was consumed, ensuring the necessary power supply. Fuel cells were selected over alternatives such as Lead storage batteries, generator sets, or Ni-Cd cells due to their superior energy density, their capacity to produce water as a byproduct, and their overall suitability for space exploration where weight and efficiency are paramount considerations.