1 Answer
The electrochemical equivalent of aluminium refers to the amount of aluminium deposited or dissolved when a current of 1 ampere flows through an electrolyte for 1 second.
For aluminium, the electrochemical equivalent can be calculated using the formula:
\[
Z = \frac{M}{n \times F}
\]
Where:
Now, plugging in the values for aluminium:
\[
Z = \frac{26.98}{3 \times 96500} \approx 9.33 \times 10^{-4} \, \text{g/C}
\]
So, the electrochemical equivalent of aluminium is approximately 9.33 × 10⁻⁴ grams per coulomb. This means that if a current of 1 ampere passes through the electrolyte for 1 second, around 9.33 × 10⁻⁴ grams of aluminium would be deposited or dissolved, depending on the direction of current.
For aluminium, the electrochemical equivalent can be calculated using the formula:
\[
Z = \frac{M}{n \times F}
\]
Where:
- \( Z \) is the electrochemical equivalent (in grams per coulomb),
- \( M \) is the molar mass of the metal (for aluminium, \( M = 26.98 \, \text{g/mol} \)),
- \( n \) is the valency of the metal (for aluminium, \( n = 3 \) since it forms Al³⁺ ions),
- \( F \) is Faraday’s constant (\( F \approx 96500 \, \text{C/mol} \)).
Now, plugging in the values for aluminium:
\[
Z = \frac{26.98}{3 \times 96500} \approx 9.33 \times 10^{-4} \, \text{g/C}
\]
So, the electrochemical equivalent of aluminium is approximately 9.33 × 10⁻⁴ grams per coulomb. This means that if a current of 1 ampere passes through the electrolyte for 1 second, around 9.33 × 10⁻⁴ grams of aluminium would be deposited or dissolved, depending on the direction of current.