A current of 0.15A is passed through an aqueous solution of K2PtCl4. How long will it take to deposit 1g of Pt (M=195.1)
To calculate the time it takes to deposit 1g of Pt, we need to start by calculating the number of electrons needed to deposit 1g of Pt.
Using the molar mass of Pt (195.1 g/mol), we can convert 1g of Pt to moles:
1g Pt X (1 mol Pt/195.1 g Pt) = 0.00512 mol Pt
We know that the reduction of Pt 4+ to Pt requires 4 electrons:
Pt 4+ + 4e- → Pt
So, the number of electrons needed to deposit 0.00512 mol of Pt is:
0.00512 mol Pt X 4 electrons/mol Pt = 0.0205 mol electrons
Now we use Faraday's constant, which gives us the number of coulombs of charge in one mole of electrons.
One molt of electrons = Faraday's constant (F) = 96,500 C/mol e-
The total charge required to deposit 0.0205 mol electrons is:
0.0205 mol e- X 96,500 C/mol e- = 1,977.25 C
Finally, we can calculate the time (t) it takes to deposit 1g of Pt using the equation:
t = Q/I
where Q is the total charge (1977.25 C) and I is the current (0.15 A).
t = 1977.25 C / 0.15 A = 13,181.67 seconds
Therefore, it would take approximately 3.66 hours to deposit 1g of Pt from an aqueous solution of K2PtCl4 with a current of 0.15A.
To determine the time it will take to deposit 1g of Pt from an aqueous solution of K2PtCl4, we need to use Faraday's law of electrolysis. Faraday's law states that the amount of substance deposited during electrolysis is directly proportional to the charge passed through the solution.
The formula to calculate the time, t, required for electrodeposition is:
t = ((m * M) / (z * F * I))
Where:
t = time (in seconds)
m = mass of substance deposited (in grams)
M = molar mass of the substance (in g/mol)
z = number of moles of electrons transferred per mole of substance during the electrochemical reaction
F = Faraday's constant (96,485 C/mol)
I = current (in amperes)
In this case, we want to deposit 1g of Pt, which has a molar mass (M) of 195.1 g/mol. The number of moles of electrons transferred (z) can be determined from the balanced equation of the electrochemical reaction. For K2PtCl4, it can be assumed that z = 2, as two moles of electrons are transferred per mole of K2PtCl4.
Given that the current (I) is 0.15A, we can substitute these values into the formula:
t = ((1 * 195.1) / (2 * 96485 * 0.15))
Calculating this gives us the time required to deposit 1g of Pt.
To determine the time it takes to deposit 1g of Pt, we need to use Faraday's law of electrolysis, which relates the amount of substance deposited to the current and the time.
Faraday's law states that the amount of substance (in moles) deposited at an electrode is directly proportional to the charge passed through the circuit. The charge (Q) is equal to the product of current (I) and time (t), and can be calculated using the formula:
Q = I * t
To convert the moles of Pt to grams, we need to know the molar mass of Pt, which is given as 195.1 g/mol.
Now, let's calculate the amount of charge (Q) required to deposit 1g of Pt:
Step 1: Calculate the moles of Pt using the molar mass:
Moles of Pt = Mass of Pt / Molar mass of Pt
Moles of Pt = 1 g / 195.1 g/mol
Step 2: Calculate the charge (Q) using Faraday's law:
Q = Moles of Pt * Faraday's constant
Q = Moles of Pt * (96,485 C/mol)
Step 3: Calculate the time (t) using the formula:
t = Q / I
Now, let's substitute the values into the equation:
t = (Moles of Pt * (96,485 C/mol)) / I
Given that the current is 0.15A:
t = (Moles of Pt * (96,485 C/mol)) / 0.15 A
Substituting the value of Moles of Pt, we get:
t = ((1 g / 195.1 g/mol) * (96,485 C/mol)) / 0.15 A
Now, calculate the value of t using a calculator or by performing the math manually:
t = 3,123,780 C/A
Therefore, it will take approximately 3,123,780 seconds (or 869 hours or 36 days) to deposit 1g of Pt.