A student measures k for a certain blue dye to be 19.4 M^-1 at 560 nm. What is the molarity of the dye in a solution with an absorbance of 0.118 at this wavelength?
A = k*c
0.118 = 19.4*c
Solve for c.
.006
To find the molarity of the dye in the solution, we can use the Beer-Lambert Law, which relates the absorbance of a solution to its molar concentration.
The Beer-Lambert Law is given by the equation:
A = ε * l * c
Where:
A is the absorbance of the solution
ε (epsilon) is the molar absorptivity or the molar extinction coefficient of the dye at a particular wavelength
l is the path length of the cuvette (usually in cm)
c is the molar concentration of the dye in the solution
In this case, we know the absorbance (A = 0.118), the molar absorptivity at 560 nm (ε = 19.4 M^-1), and we need to find the molar concentration (c) of the dye.
Rearranging the equation, we get:
c = A / (ε * l)
Since we don't have the path length, we will assume it to be 1 cm, which is standard.
Now, substituting the known values into the equation:
c = 0.118 / (19.4 M^-1 * 1 cm)
Simplifying the equation:
c = 0.118 / 19.4 M^-1
c ≈ 0.00608 M
Therefore, the molarity of the dye in the solution is approximately 0.00608 M.