Solve Electrolysis Equation: 11.2L of Oxygen at STP

[RPN]

hello
I saw someone had posted a question about this question but I too do not have a clear understanding of how to solve this equation. I have a very small section in my text about electroylsis and I have looked for information in others texts but I still need help. Please if anyone has an idea can you go through it step by step for me.
The question reads:
By the electrolysis of water, 11.2L of oxygen at STP was prepared
a) What charge was required?
b) If a current of 0.5A was used, how long did it take?
thank you

 

[Gokul43201]

Follow the standard steps :
1. Write the balanced eqn
2. Determine the mole ratio of electrons transferred to oxygen produced
3. From the given volume of oxygen determine the number of moles of electrons
4. Convert moles of electrons to charge in coulombs

OR

Balance and use Faraday's Law

From the charge transferred, the time taken at some current can be found

 

[ravenprp]

To solve this electrolysis equation, we need to first understand the basic principles of electrolysis. Electrolysis is the process of using an electric current to decompose a compound into its constituent elements. In this case, we are decomposing water (H2O) into its constituent elements, hydrogen (H2) and oxygen (O2).

The first step is to write out the balanced chemical equation for the electrolysis of water:

2H2O(l) → 2H2(g) + O2(g)

From this equation, we can see that for every 2 moles of water that undergo electrolysis, we will produce 1 mole of oxygen gas. Since we are given the volume of oxygen gas, we need to convert it to moles using the ideal gas law:

PV = nRT

Where:
P = pressure (in atm)
V = volume (in L)
n = moles
R = gas constant (0.0821 L·atm/mol·K)
T = temperature (in K)

At STP (Standard Temperature and Pressure), the pressure is 1 atm and the temperature is 273 K. So, using the ideal gas law, we can calculate the number of moles of oxygen gas produced:

n = PV/RT
= (1 atm)(11.2 L)/(0.0821 L·atm/mol·K)(273 K)
= 0.5 moles

Since we need 2 moles of water to produce 1 mole of oxygen gas, we can calculate the total moles of water needed:

0.5 moles O2 x (2 moles H2O/1 mole O2) = 1 mole H2O

Now, we need to convert moles of water to grams of water using the molar mass of water (18 g/mol):

1 mole H2O x (18 g/mol) = 18 g H2O

Next, we need to calculate the amount of charge (in coulombs) required to decompose 18 g of water. This can be done using Faraday's law:

Q = nF
Where:
Q = charge (in coulombs)
n = moles of electrons transferred
F = Faraday's constant (96,485 C/mol)

Since 2 moles of electrons are transferred for every 2 moles of water decomposed, we can calculate the charge required:

Q =