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 0.550. If a trial's absorbance is measured to be 0.350 and its initial concentration of SCN– was 0.0010 M, the equilibrium concentration of SCN– will be _____?
Can you please explain to me how to work this question out? Thanks
The concentration of SCN- in the first mixture is:
(1/10)(0.0020M) = 0.0020M)
The concentration of Fe+3 was:
(9/10)(0.200M) = 0.18 M
Since the concentration of F+3 was 90 times the concentration of SCN-, we may assume the equilibrium, Fe+3(aq) + SCN-(aq) <=> FeSCN+2(aq), if far enough to the right to assume completion: Most of the SCN- became FeSCN+2.
[SCN-](initial) = [FeSCN+2](final) = 0.0020M
FeSCN+2(aq) is resopnsible for the absorbance.
The absorbance, A, is proportional to the concentration.
A = (const.)C, or
C = A/(const.)
The proprtionality constant is:
0.550/0.0020M= 2750 M^-1
and our formula becomes,
A = 2750*C , where C = [FeSCN+2]
Use A = 0.350 to calculate C, and finally,
[SCN-](eq) = 0.0010 - C
A comprehensive look at this equilibrium system and the experiment for determining the equilibrium constant is given here:
http://www.scienceteacherprogram.org/chemistry/Tehilla98.html
Hi GK,
For the first step, isn't (1/10)(0.0020) = 0.00020 M?
Also, I understand your very clear explanation to the last step: [SCN-] = 0.0010 -C
I understand the minus C part, but how did you get 0.0010?
Thank you!
To solve this question, we'll use the Beer-Lambert Law and the concept of equilibrium.
The Beer-Lambert Law states that the absorbance of a solution is directly proportional to the concentration of the absorbing species in the solution. Mathematically, it can be represented as:
A = εbc
where A is the absorbance, ε is the molar absorptivity (a constant specific to the absorbing species and the wavelength of light used), b is the path length of the cuvette (usually 1 cm), and c is the concentration of the absorbing species.
In this question, we are given the absorbance of the standard solution, A_STD = 0.550, and the initial concentration of SCN– in the trial, c_initial = 0.0010 M.
We can start by calculating the molar absorptivity constant (ε) for SCN–. Since it is not given, we'll assume it is the same as that of the standard solution.
ε = A_STD / (bc)
= 0.550 / (1 cm * c_STD)
Next, we'll use the Beer-Lambert Law to calculate the equilibrium concentration of SCN– in the trial. We are given the absorbance of the trial, A_trial = 0.350, and we need to find c_eq, the equilibrium concentration of SCN–.
A_trial = ε * b * c_eq
Rearranging the equation, we get:
c_eq = A_trial / (ε * b)
Substituting the values, we have:
c_eq = 0.350 / (ε * 1 cm)
Now, we can substitute the previously calculated value of ε to find c_eq.
c_eq = 0.350 / (ε * 1 cm)
Finally, we can plug in the expression for ε calculated earlier:
c_eq = 0.350 / [(0.550 / (1 cm * c_STD)) * 1 cm]
Note that the path length (b) cancels out, as it is the same for both the standard and the trial. And c_STD is the initial concentration of SCN– in the standard solution, which is not given. Unfortunately, we do not have enough information to calculate the equilibrium concentration of SCN– in this trial unless the concentration of SCN– in the standard solution is provided.
In summary, without the initial concentration of SCN– in the standard solution, it is not possible to determine the equilibrium concentration of SCN– in the trial.