3: Atoms where you live- Water Quality by Molecular and Atomic Spectroscopy

Background Information

Water quality is a hugely important part of public health and safety, but how do we monitor it? Today’s lab will explore metal ions as part of water quality, though water quality involves many other factors as well. We will specifically focus on measuring \(\ce{Fe^{2+}}\), \(\ce{Na^{+}}\), \(\ce{Ca^{2+}}\), and \(\ce{Mg^{2+}}\) in water. \(\ce{Ca^{2+}}\) and \(\ce{Mg^{2+}}\) are the primary cause of hard water. Due to their low solubility, they easily precipitate and cause solid scale to form which can clog up pipes. Other factors of water quality, such as the pH of water or the presence of pathogens, toxic metals, or compounds, will not considered in today's lab.

In parts A-B, we will use a spectrophotometer to measure the amount of \(\ce{Fe^{2+}}\) in solution through absorbance. In order for this to work, we need to have a compound with visible color, so we will react the \(\ce{Fe^{2+}}\) with phenanthroline, which gives a reddish color that changes in intensity depending on the amount of \(\ce{Fe^{2+}}\) present. This will require making a standard curve (part A) so that we can then test our unknowns (part B), and will be a nice review of what we did in lab last week.

Parts C-D will look at atomic emission spectroscopy, where we detect atoms by observing electron movement between shells when given energy. In part C, we will examine hydrogen, the simplest element with only 1 electron, and then carry principles from that over to look at more complicated metals using the Microwave Plasma Atomic Emission Spectrometer (MP-AES) in the instrument room across the hall. This instrument works by converting all atoms to a high-energy plasma state, which excites electrons in each atom to the highest energy shell possible. When the electrons come crashing back down, you get an extremely large sample per atom, allowing you to detect metals that are present even at very low concentrations.

Section A: Calibration of the Spectrophotometer

Just like in our previous lab, the spectrophotometer must be calibrated to understand the relationship between absorbed light and iron concentrations. You will begin by running four reference standards with known iron concentrations: 1 ppm, 2 ppm, 3 ppm, and 6 ppm. We also need a blank, a solution with all reagents except the analyte. Water and other reagents absorb a small amount of light, and this blank allows you to measure that amount so that you can subtract it from the light absorbed by the orange iron-phenanthroline compound.

Procedure:

1. Obtain the following solutions in separate beakers from your kits. Carefully label these solutions so that you do not get them mixed up. These are the solutions you will use to prepare your tests and Unknown iron samples.

a) About 20 mL of the 0.538 mM iron solution (\(\ce{Fe^{2+}}\)).

b) About 25 mL of 0.002 M phenanthroline (Phen).

c) About 15 mL of sodium acetate solution (\(\ce{NaOAc}\)).

d) About 15 mL of hydroxylamine (\(\ce{NH2OH}\)).

e) About 20 mL of deionized water (DI)

2. Download a local copy of the linked Excel spreadsheet, enter the concentration of the available iron standard (C₁) in ppm in cell B2.

3. For each target concentration (C₂), you will make up a 3.0 mL (V₂) solution directly into a cuvette. Label a piece of paper blank, 1 ppm, 2 ppm, 3 ppm, and 6 ppm and place cuvettes above each label. Recall the equation you generated by dimensional analysis was \(\ce{C1$*$V1 = C2$*$V2}\). You will use this equation often throughout your science coursework and should memorize it.

4. To each solution you will add:

a. 0.600 mL of 10% sodium acetate buffer solution.

b. 1.500 mL of 1,10-phenanthroline solution

c. 0.300 mL of hydroxylamine solution

d. Enough iron solution to obtain the target concentration.

e. Enough water to reach 3.000 mL total

5. Create a formula in column B6-B9 that calculates the volume of standard (V₁) you must add to get 3.0 mL (V₂, column C) of the target concentration (C₂, column A), given the standard concentration (C₁ in B2).

6. Cell G5 is set up with a formula for how much water needs to be added to get the total water volume up to 3 mL based on all the other components. Drag the formula down to fill cells G6-G9.

7. Make up your solutions according to your calculations. Use a clean tip for each reagent. Add the 1,10 phenanthroline last in two measurements of 0.750 mL. After creating each solution, mix by pipetting up and down repeatedly, moving from low concentration to high.

8. Connect a SpectroVis to a LabQuest. Make sure the units are set to “Abs”. If they aren’t, click on the red banner and select “Change Units”.

9. Click on the red banner and select “Calibrate”. Follow the instructions on screen.

10. Click on the red banner again and select “Mode”. Choose “Events with Entry”, and set the wavelength to 562 nm. For the event, give the label “[Fe]” and the unit “ppm”.

11. Use the “play” button to start taking data.

12. Place your 1 ppm sample in the reader, making sure the clear sides line up with the light and arrow icons on the SpectroVis. Hit “Keep” on the bottom of the screen, and enter “1” for the first event. Repeat for 2, 3, and 6 ppm samples.

13. Hit “stop” to complete data collection.

14. Use the “Table” icon to navigate to your data in table form, and record the absorbance data on your spreadsheet in cells H5-H9.

15. Create a scatterplot with Target concentrations (A6-A9) as the x-axis and Absorbance (H6-H9) as the y-axis. Make sure to label axis and the graph clearly. Refer to last week’s labs if you forgot how to do this.

16. Add a linear-trendline for the displayed data, showing the equation and R² value.

Section B: Determination of Iron in an Unknown Sample

Now that we have a calibration relationship between the iron-phenanthroline complex and absorbance, we can measure the absorbance of samples to determine iron (II) concentration.

Procedure:

1. Gravity filter by your sample. (Preparing a filter). Retain this filtered sample for part D as well.

2. Obtain a clean, dry cuvette.

3. Using a calibrated pipette, measure out 0.600 mL of your filtered sample, and add it to a cuvette.

4. To each solution add:

a. 0.600 mL of 10% sodium acetate buffer solution.

b. 0.300 mL of hydroxylamine solution

c. 1.500 mL of 1,10-phenanthroline solution

5. Measure 2 more 3.0 mL samples for a total of 3 measurements. Record each in your local copy of the attached spreadsheet.

6. In the column next to these absorbance values (y), calculate concentration (x) using the equation for the trendline you created in Section A (solve for x in terms of y).

7. Remember that the calibration you performed was based on the final concentrations, but your water sample only makes up 0.6 mL of 3.0 mL. Using the equation C₁*V₁ = C₂*V₂, find the concentration of iron in your sample before dilution.

8. Extend these calculation for each absorbance value. Calculate the mean and standard deviation for both absorbance and the concentration.

9. There are better ways to use the linear regression data for your trendline to report uncertainty, but the simplest way of reporting this data is as the mean concentration ± the standard deviation of the concentration.

Section C: AE Analysis of Hydrogen

The first technique used in this lab analyzed specifically for \(\ce{Fe^{2+}}\). With AE, you can analyze many ions from the same sample solution, measuring many different water quality parameters at once. Calcium and magnesium are the key ions that contribute to the hardness of water, and are the most abundant cations in solution. Sodium is included for later charge balance analysis. The iron analysis will enable you to compare your results from the AE test with those from the SpectroVis measurements. Before we run the atomic emission experiment for more complex atoms, we will look at hydrogen emission spectrum which matches the material The hydrogen emission spectrum will help us to understand what is happening in the microwave-plasma atomic emission spectrophotometer (MP-AES).

As covered in class, atoms, molecules and subatomic particles have energy. If something gains energy, then its total energy content increases. But if something loses energy then its total energy content decreases. You also know that the ground state is a molecule’s lowest possible energy state while all higher energy levels are called excited states. We have also just learned that these are quantum states, which means that only certain levels of energy are allowed. In this lab we will observe several transitions between these energy states and see how the energy gained or lost must match the difference between the energies of two levels within the molecule.

This energy that is gained or lost can be in the form of thermal energy (from collisions between particles or molecular vibrations) or can be in the form of electromagnetic radiation. We see both of these energy changes at the same time in molecular spectroscopy, resulting in the broad absorption spectrum you saw in parts A and B. Atomic spectroscopy differs from molecular spectroscopy because the lines in atomic spectroscopy are much thinner than the wide bands we see in molecular spectroscopy, owing almost entirely to electronic transitions. Because atoms don’t have bonds like molecules, there is no widening in the electronic transitions from the vibrations of the bonds. Every element has electronic transitions in the UV-Vis region, meaning that we can use this method for most elements, without any sensitizing compound (like the phenanthroline from parts A and B), and because the signals are narrow lines rather than broad bands, modern instruments can resolve many different molecules simultaneously. On the other hand, atomic spectroscopy tends to produce atoms or occasionally ions of the same oxidation state for analysis, regardless of their oxidation state in solution.

We will use electricity to excite hydrogen gas (our simplest atom) and observe its visible emission spectra (Balmer lines).

1. Obtain one spectroscope for your group and the LabQuest setup that consists of a LabQuest, SpectroVis, fiber optics cable, and USB cable. Do not crimp the optics cable! Place the black block at the end of the fiber optics cable into the cuvette position properly (the white triangles aligned with each other) or you will not get any spectra on the LabQuest screen. Change the LabQuest units to intensity.

2. With the LabQuest, observe and record the spectra for hydrogen. Record as many lines as you can see. Fill in the charts on your local copy of the linked Excel spreadsheet on the “Hydrogen Emission” sheet under “Balmer lines”.

3. Convert the wavelengths observed in nm to wavelength in meters, then to Energy in kJ / mol (E = hc/λ). Constant values are given in the Excel spreadsheet, so if you use absolute references (ie $A$1), you should be able to create one formula and drag it down.

4. Use the next table to calculate the seven lowest energy levels and the highest possible energy level (accomplished by using an arbitrarily large number) for the electron orbits for hydrogen as described by Bohr . Remember n = energy level and R is the Rydberg constant.

5. The numbers directly from the equation are quite small with units for E given in J/electron. Convert these into units of kJ/mol in the next column.

6. The Balmer lines are all from transitions between higher shells and shell 2. Calculate the energy of these transitions using the “Energy of transitions table” on the spreadsheet.

7. The energy of the Balmer lines should closely align with energy of the electronic transitions. Place the wavelength values you detected next to the energy transitions to which they are closest.

Section D: AE Analysis of Samples for Mg, Ca, Na, and Fe concentrations

Goals: (1) To utilize an analytical tool that can selectively measure ions in a mixed solution. (2) To determine the concentration of Ca²⁺, Mg²⁺, Na⁺, and Fe²⁺ in your water sample.

Discussion: The first technique used in this lab analyzed specifically for Fe^2+-. With AE, you can analyze four (or more) ions from the same sample solution, thus measuring a number of different water quality parameters at once. We will tests for Ca²⁺, Mg²⁺, Na⁺, and Fe²⁺.

Experimental Steps:

1. Take the samples across the hall to the AE.

2. Conduct analyses for \(\ce{Mg}\), \(\ce{Ca}\), \(\ce{Na}\), and \(\ce{Fe}\). Record into your local copy of the linked Excel spreadsheet, including units. Use this data in the determination of total positive charge in your solutions. You will need to perform unit conversions to get from the units generated by our calibrations (parts per million (ppm) or parts per billion (ppb) to Molar units (M), which are moles per liter. These units are similar in form to percent, literally “per hundred”. It may be convenient to think of these as similar units – just as a percent can be thought of as the mass present in 100 grams of solution, ppm can be thought of as the mass present in 1 million grams of solution and ppb can be thought of as the mass present in 1 billion grams of solution. From here you can use the molar mass of the analyte (g analyte / mole analyte) and the density of the solution (g solution / mL solution) to convert to Molar (mole analyte / L solution). You may use Excel to do the calculations, or show your work on paper.

\(\ce{$x$ mass percent (part per hundred) = \Large\frac{$x$ g analyte}{10^2 g solution}}\)

\(\ce{$x$ ppm analyte = \Large\frac{$x$ g analyte}{10^6 g solution}}\)

\(\ce{$x$ ppb analyte = \Large\frac{$x$ g analyte}{10^9 g solution}}\)

Post-lab questions

1) Based on the slope of your trendline for iron analysis, what is the molar absorptivity ( ) of the iron solution at 562 nm?

2) Show that a 0.538 mM solution of ammonium iron (II) sulfate ((NH₄)₂Fe(SO₄)₂(H₂O)₆) is 30 ppm in iron by converting the units.

3) Why is it important to calibrate a spectrometer? How does the calibration indicate sensitivity? To what part(s) of the Beer-Lambert law do(es) this relate? Why is the buffer important?

4) In your own words, describe how the molecular spectroscopy works. List an advantage and a disadvantage of this technique compared to atomic emission (AE). Which:

· has a higher signal / atom (sensitivity)

· is easier to prepare

· is faster

· is more accessible

· is more error prone

· can determine ion charge / oxidation state?

5) Another version of this lab uses ferrozine to bind iron. If the calibration has been completed for the iron-phenanthroline complex, reading a iron-ferrozine complex gives an absorbance of nearly zero. Why?