A:For this question, assume the vitamin sample solution put into the spectrophotometer had an absorbance reading of 0.248, and the equation of the best-fit line from the calibration curve was y = 420 (x) + 0.00001 . What is the concentration of Fe2+ in the vitamin solution that was put into the spectrophotometer.
B:Assume that the known Fe solution that was put into the spectrophotometer was determined to have an Fe concentration of 0.0035 mg/mL . Follow the calculations on page 6 of Experiment 22 to calculate the mg Fe in the orginal Fe sample used (calc. step 1).
With no more details than those given I'm out on a limb here. I assume that the best fit line fits the equation A = ebc where e is the molar absorptivity constant = 420, y is Absorbance, I asume c you want in moles/L and b is the path length which I assume is 1 cm.
y = 420 (x) + 0.00001
1.
0.248 = 420 X + 0.00001. I assume X = concentration and usually that is measured in moles/L and I assume you want to calculate concentration in moles/L. So figure away and solve for X = C. However, since part B gives concentration in mg/mL I may not have assumed correctly.
When we go to part B of your question I have no idea what page 6 of experiment 22 is. Sorry about that.
You just don't give enough definitive information to help me help you.
A: To calculate the concentration of Fe2+ in the vitamin solution, we need to use the equation of the best-fit line from the calibration curve. The equation is given as:
y = 420x + 0.00001
Here, y represents the absorbance reading and x represents the concentration of Fe2+.
Given that the absorbance reading of the vitamin sample solution was 0.248, we can substitute this value into the equation and solve for x:
0.248 = 420x + 0.00001
Simplifying the equation:
420x = 0.248 - 0.00001
420x = 0.24799
x = 0.24799 / 420
x ≈ 0.0005905
Therefore, the concentration of Fe2+ in the vitamin solution is approximately 0.0005905.
B: To calculate the amount of Fe in the original Fe sample used, we will follow the calculations on page 6 of Experiment 22, assuming the known Fe solution had an Fe concentration of 0.0035 mg/mL.
The calculation steps are as follows:
1. Determine the volume of Fe solution used:
Let's assume we used V mL of the known Fe solution.
2. Calculate the amount of Fe in the solution used:
Amount of Fe = Concentration of Fe × Volume of Fe solution used
= 0.0035 mg/mL × V mL
= 0.0035V mg
Therefore, the amount of Fe in the original Fe sample used can be calculated by multiplying the concentration of Fe (0.0035 mg/mL) by the volume of Fe solution used (V mL).
A: To calculate the concentration of Fe2+ in the vitamin solution, we need to use the equation of the best-fit line from the calibration curve. This equation relates the absorbance (y) to the concentration of Fe2+ (x).
In the given equation, y = 420 (x) + 0.00001, we can see that the coefficient of x is 420. This means that for every unit increase in absorbance, the concentration of Fe2+ increases by 420 units. The intercept term (0.00001) can usually be neglected in calculations.
Given that the absorbance reading of the vitamin sample solution was 0.248, we can now calculate the concentration of Fe2+. We divide the absorbance by the coefficient of x:
concentration of Fe2+ = 0.248 / 420
concentration of Fe2+ ≈ 0.00059
Therefore, the concentration of Fe2+ in the vitamin solution that was put into the spectrophotometer is approximately 0.00059.
B: To calculate the mg Fe in the original Fe sample, we need to follow the calculations on page 6 of Experiment 22. Unfortunately, I cannot access specific pages or external documents. However, I can provide you with general steps to calculate the mg Fe in the original Fe sample.
1. Start by calculating the volume of the known Fe solution used.
- Let's assume you used a volume of V mL of the known Fe solution.
2. Determine the concentration of Fe in the known Fe solution.
- From the given information, the concentration of Fe in the known solution is given as 0.0035 mg/mL.
3. Calculate the amount of Fe in the known Fe solution using the equation:
amount of Fe (mg) = concentration of Fe (mg/mL) × volume (mL)
4. Now, you have the amount of Fe in the known solution. Next, you need to calculate the mg Fe in the original Fe sample used (calc. step 1). Since I don't have the specific calculation steps, I suggest referring to the instructions or materials provided in Experiment 22 to proceed with the calculation.
Remember to use the correct units and double-check your calculations to ensure accuracy.