The standard solution of FeSCN2+ (prepared by combining 9.00 mL of 0.200 M Fe(NO3)3 w/1.00 mL of 0.0020 M KSCN) has an absorbance of 0.530. If a trial's absorbance is measured to be 0.150 and its initial concentration of SCN– was 0.00070 M, the equilibrium concentration of SCN– will be ?
Scroll down to the 25th question (authored by Amy) on this board, just below your post, and I addressed that same question, although the concentrations are different.
To calculate the equilibrium concentration of SCN–, we can use the Beer-Lambert Law, which states that the absorbance (A) of a solution is directly proportional to the concentration (c) of the absorbing species.
The equation for the Beer-Lambert Law is:
A = ε * b * c
Where:
A is the absorbance,
ε is the molar absorptivity (a constant),
b is the path length (the distance the light travels through the solution), and
c is the concentration.
In this case, we have the absorbance (0.150), the initial concentration of SCN– (0.00070 M), and the standard solution's absorbance (0.530). We need to calculate the equilibrium concentration of SCN–.
Let's solve this step-by-step:
Step 1: Calculate the molar absorptivity (ε):
We can rewrite the Beer-Lambert Law equation as:
A = ε * c
Since the path length (b) is constant, we can assume it cancels out. Therefore, A and c are directly proportional.
We can rearrange the equation to solve for ε:
ε = A / c
Using the standard solution's absorbance (0.530) and the provided concentration (0.0020 M), we can calculate ε as follows:
ε = 0.530 / 0.0020 = 265 M^(-1) cm^(-1)
Step 2: Use the calculated ε and the provided absorbance and concentration to find the equilibrium concentration of SCN–:
Using the rearranged Beer-Lambert Law equation A = ε * c, we can rearrange the equation to solve for c:
c = A / ε
Substituting the values, we have:
c = 0.150 / 265 M^(-1) cm^(-1)
Calculating this, we get:
c ≈ 0.00057 M
Therefore, the equilibrium concentration of SCN– will be approximately 0.00057 M.