Calculate the EMF of the cell CR|Cr3+(0.1M)||Fe2+(0.1M)|Fe
Cell reaction- 2Cr(s)+3Fe2+(aq)->2Cr3+(aq)+3Fe(s)
How fool completely wrong answer you confused more than before you idiot
If u don't know how to solve don't post the solution bloody hell
E°cell= E°cathode-E°anode
E°cell=-0.44-(-0.74) = 0.30V
Ecell= E°cell - 0.0591/n log [Cr3+]^2/[Fe2+]^3
Ecell= 0.30- 0.0591/6 log[0.1]^2/[0.1]^3
Ecell= 0.30 - 0.0591/6 log 10
Ecell= 0.30 - 0.0591= 0.2409V
Ecell=Ecathode-Eanode
=-0.4-(-0.9)
Ecell=0.5
nernst equation:
=ecell-0.059/n log product/react.
=0.5-0.059/2 log 0.1/0.1
=0.5-0.0295×0
=0.5
To calculate the EMF (Electromotive Force) of the given cell, you can use the Nernst equation. The Nernst equation relates the cell potential to the concentration of the species involved in the cell reaction. The formula for the Nernst equation is as follows:
Ecell = E°cell - (RT/nF) * ln(Q)
Where:
- Ecell is the cell potential
- E°cell is the standard cell potential
- R is the gas constant (8.314 J/(mol·K), or 0.0821 L·atm/(mol·K))
- T is the temperature in Kelvin
- n is the number of electrons involved in the cell reaction
- F is the Faraday constant (96,485 C/mol)
- Q is the reaction quotient
To determine the EMF of the cell, we need to substitute the relevant values into the Nernst equation:
1. Calculate Q:
Q = [Cr3+]/[Fe2+]
2. Determine the standard cell potential, E°cell:
The E°cell is given as not provided. To proceed, we will assume that it is zero, as we don't have information to determine its exact value.
3. Substitute the values into the Nernst equation and solve:
Ecell = 0 - (RT/6F) * ln([Cr3+]/[Fe2+])
Please provide the temperature (in Kelvin) at which you want to calculate the EMF so that we can proceed with the calculation.