you add 0.05ml of 0.10M agno3 to 4ml of 2.0 nacl it forms saturated solution of agcl<---> ag+ + cl-
using a saturated solution you set up
ag | ag+( test solution after reaction) || ag+(1.0ml) | ag
how do we calc the equilibrium of ag+ of the test solution after reaction ?
do we use nernsts????
would the answer be 2.46x10^-21?!
To calculate the equilibrium concentration of Ag+ in the test solution after the reaction, you can indeed use the Nernst equation. The Nernst equation relates the concentration of ions in a solution to the measured potential difference (or electromotive force) of an electrochemical cell.
The Nernst equation is given by:
E = E° - (0.0592/n) * log10([Ag+]/[Ag+1]),
Where:
E is the measured potential difference
E° is the standard electrode potential (for the Ag/Ag+ half-cell)
n is the number of electrons transferred in the cell reaction
[Ag+] is the concentration of Ag+ in the test solution after the reaction
[Ag+1] is the concentration of Ag+ in the standard solution (1.0 mL in this case)
In this case, the half-cell reaction is AgCl ↔ Ag+ + Cl-. Since the Ag/Ag+ half-cell is after the reaction, we need to determine the concentration of Ag+ using the Nernst equation.
1. Determine E°:
Look up the standard electrode potential for the Ag/Ag+ half-cell, which is usually given as +0.800 V. This value is already standard, so E° = +0.800 V.
2. Determine n:
In this reaction, 1 mole of Ag+ is produced, meaning 1 electron is transferred. Therefore, n = 1.
3. Measure E:
Use a voltmeter to measure the potential difference (E) of the electrochemical cell.
4. Solve for [Ag+]:
Rearrange the Nernst equation to solve for [Ag+]:
[Ag+] = 10^(([Ag+1]/0.0592) * (E° - E)).
Substitute the measured values for E, [Ag+1] (concentration in the standard solution), and E° into the equation to calculate the equilibrium concentration of Ag+ in the test solution after the reaction.
Remember to use proper units throughout the calculation, such as volts (V) for E and moles per liter (M) for concentrations.