Using Volumetric Glassware, Part 1

https://www.youtube.com/watch?v=4y_M1N4-EjM

Using Volumetric Glassware, Part 2

https://www.youtube.com/watch?v=j2vPLmdxIqs

**CHEM 150L Iron Ore Calculations Tutorial (F20a)**

https://mediaspace.minnstate.edu/media/1_166nk1qm

Quantitative Analysis of Fe in Iron Ore

INTRODUCTION: Color is one of the easier properties of a solution to monitor. The color that is observed by the eye is dependent upon the wavelengths of light that are either absorbed or transmitted by a material, known as the spectrum of a solution. In addition to the wavelengths that are absorbed, it is also important to understand how efficiently a solution absorbs light. This can be quantified by a value called the extinction coefficient (also called molar absorptivity) and Beer’s Law:

A c = ε where:

A = absorbance

ε = extinction coefficient

= path length

c = concentration

The measured absorbance of a solution depends upon not only the extinction coefficient, but also on the concentration of the solution and the amount of solution the light must pass through (the path length). For a particular wavelength and cuvette size, ε and will be constants, so a plot of A vs. c should yield a straight line with slope = ε and intercept close to zero.

In this experiment, you will use a colorimetric method to analyze the % Fe by mass in a sample of iron ore. Samples of known concentration of the red-orange colored iron-phenanthroline complex ion will be prepared, and the absorbance of these solutions will be measured. These solutions are referred to as standard solutions. Once the relationship between the concentration of this complex ion and its absorbance is established, measurement of the absorbance values of similarly prepared iron ore solutions can be used to calculate the concentration of Fe in the prepared solutions and ultimately the % Fe in the iron ore sample.

SAFETY NOTES: Care should always be taken when working with acids. If any acid is spilled on your skin, wash immediately. Any acid spills on the lab bench should be cleaned up immediately. Always use a rubber bulb to fill pipets. Dispose of all colored solutions in the appropriate waste container.

**PART I: Preparation of Standard Fe2+ Solutions **

1. Watch the prelab video at https://www.youtube.com/watch?v=QJzPae48ZDo.

2. A stock Fe2+ solution has been prepared in advance by weighing the appropriate amount of iron(II) sulfate heptahydrate, FeSO4·7H2O, and dissolving it in water to make a solution that is approximately 9 × 10-4 M in Fe2+(aq).

3. Watch the video of the preparation of the standard solutions at https://mediaspace.minnstate.edu/media/1_q4ocbjnu. Record the precise concentration of the iron standard solution in your lab notebook from the video.

4. Four 50.0-mL volumetric flasks are obtained, and pipets are used to deliver 2.00 mL, 4.00 mL, 6.00 mL, and 8.00 mL of the stock Fe2+ solution into the different flasks.

5. A 150-mL beaker is obtained, and graduated cylinders are used to add the following three reagents to the beaker, followed by mixing. a. 10 mL of 10% hydroxylamine hydrochloride (keeps iron in the 2+ state). b. 20 mL of 1 M sodium acetate (buffer, prevents the pH from changing drastically). c. 50 mL of 0.30% o-phenanthroline (produces blood-red colored complex ion with Fe2+).

6. A graduated cylinder is used to add 16 mL of the three-reagent mixture to each volumetric flask containing the iron solution and also to a fifth flask containing no iron solution to serve as a blank (needed to calibrate the spectrophotometer).

7. Each flask is gently swirled, then diluted to the calibration mark with deionized water.

8. Each flask is stoppered, and the diluted solutions are mixed well by inverting 10 times and shaking each time while upside-down.

9. The solutions are allowed to react for 30 min for the colors to develop.

10. Calculate the concentration of Fe2+ in each of the diluted standard solutions using the dilution equation: MV M V ii f f = where Mi is the undiluted concentration (listed on the stock bottle), Vi is the amount of Fe2+ solution added by pipet, Mf is the diluted concentration, and Vf is the diluted final volume (flask volume). Round your concentrations to the appropriate number of significant figures.

**PART II: Generation of the Visible Spectrum and Calibration Curve **

1. Based on the color of the iron-phenanthroline complex ion and the results of the Exploring Color experiment, predict the approximate wavelength of maximum absorbance in the spectrum.

2. Watch the video of measuring the spectra of the standard solutions at https://mediaspace.minnstate.edu/media/1_6xxwcuv8.

3. The spectrometer is calibrated using the blank solution containing no iron.

4. After the solutions have reacted for at least 30 minutes, the absorbance spectrum is recorded for all 4 solutions.

a. The least concentrated sample is used first, and the cuvette is rinsed multiple times with each successive sample before recording the spectrum. The commands “Stop” → “Collect” → “Store Latest” are used to keep the spectra of all solutions on the screen.

b. The same cuvette and spectrophotometer is throughout this experiment.

c. The wavelength of maximum absorbance, λmax, for the most concentrated solution is determined. Record in your lab notebook both this wavelength and the absorbance for each of the standard solutions at this wavelength. (Refer to the data sheet on D2L.)

5. Using Excel, prepare a plot of Absorbance vs. Concentration for the four standard solutions. This means we want absorbance on the y-axis and concentration on the x-axis, so type your concentration data in column A and absorbance data in column B of the Excel spreadsheet.

a. Insert a linear trendline, and include the best-fit equation (modify the x and y variables to be meaningful labels) and R2 value on the graph.

b. You may stop at this point for Week 1.

► Question for thought: Why is it usually considered better practice to start with the lowest concentration and go in order to the highest concentration when taking a series of absorbance measurements?

**PART III: Preparation and Analysis of Iron Ore Sample**

1. A numbered vial containing an iron ore sample is obtained. This is your unknown. Record the unknown number in your lab notebook. (You have been assigned a random sample; refer to the data sheet on D2L.)

2. The full vial with iron ore is weighed to ±0.001 g. Record its mass.

3. The ore sample is carefully transferred to a clean, dry, 50-mL beaker.

a. This is done by placing the open vial up inside the upside-down beaker, then inverting the vial and beaker together.

4. The empty vial is reweighed. Record its mass, and use subtraction to determine and record the amount of ore transferred to the beaker.

5. Watch the video of weighing and dissolving the iron ore at https://www.youtube.com/watch?v=CD8GR30DLAs. Do NOT use the mass values stated in the video! (Refer to the data sheet on D2L.)

6. In the fume hood, 3 droppers full (several mL) of concentrated HCl is added to the iron ore sample.

7. The mixture is gently heated on a hot plate for 30-60 s (depending on the hot plate temperature) 3 until the dark solid has completely dissolved. It is ensured that the solution does not boil.

a. If white solid remains, it will dissolve when the water is added.

8. When the dark solid has completely dissolved, ~10 mL of deionized water is quickly added from a graduated cylinder.

a. Note: This is not the normal safe protocol. Typically, when diluting a concentrated acid, the acid should be added to water to aid in dissipating the heat generated from the dilution. Water should not be added to acid as it may cause the acid to splatter. Because we are using a small volume of acid in this procedure and adding a relatively large volume of water, there is not much risk of the acid splashing if the water is added smoothly and quickly.

b. Because hot HCl gives off dangerous fumes, the beaker is not removed from the fume hood until the sample has been diluted.

9. The iron ore solution is transferred to a 100.0-mL volumetric flask, taking care to rinse the beaker well in order to get all of the iron into the flask. Watch the video of the dilution process at https://mediaspace.minnstate.edu/media/1_jfkvj06i.

10. The solution is diluted to the calibration mark with deionized water and mixed thoroughly. This solution contains all of the iron that was in the original ore sample.

11. 1.00 mL of the iron ore solution is pipetted into each of three 50.0-mL volumetric flasks. (We will be doing three replicate measurements.)

12. A 150-mL beaker is obtained, and graduated cylinders are used to add the following three reagents to the beaker, followed by mixing.

a. 8 mL of 10% hydroxylamine hydrochloride.

b. 16 mL of 1 M sodium acetate.

c. 40 mL of 0.30% o-phenanthroline.

13. A graduated cylinder is used to add 16 mL of the three-reagent mixture to each 50.0-mL volumetric flask containing the iron ore solution and also to a fourth flask containing no iron ore solution to serve as a blank.

14. Each flask is gently swirled, then diluted to the calibration mark with deionized water.

15. Each flask is stoppered, and the diluted solutions are mixed well by inverting 10 times and shaking each time while upside-down.

16. The solutions are allowed to react for 30 min for the colors to develop.

a. Note that each of these solutions contains only one hundredth of the amount of Fe2+ in the original sample solution.

17. Watch the video of measuring the spectra of the iron ore samples at https://mediaspace.minnstate.edu/media/1_fx4v4ax3.

18. The spectrometer is calibrated using the blank solution containing no iron.

19. After the solutions have reacted for at least 30 minutes, the absorbance spectrum is recorded for all 3 solutions.

20. Record the absorbance of all three samples at the same λmax chosen for the standard solutions. (Refer to the data sheet on D2L.)

**PART IV: Calculations**

1. Use the best-fit equation determined from the standard solutions (Part II) to calculate the concentration of each of the three prepared samples of iron ore.

a. A visual inspection of the graph should give an estimate. To obtain more precise values, plug in the absorbance value for each sample into the Beer’s law trendline and calculate the concentration corresponding to that absorbance.

b. It is usually considered a better practice to calculate the concentration of each sample, from which you can determine the % Fe in the ore, and then take the average of the results, rather than to take the average absorbance at the beginning of the process and only do the concentration calculation once. This will make the error calculation easier.

2. Use the dilution equation to determine the concentration of Fe2+ in your 100.0 mL solution.

a. You will be computing Mi from Vi (the amount pipetted into each smaller flask), Mf (computed from the trendline), and Vf (the volume of the smaller flask).

3. Compute the number of moles of Fe2+ in the 100.0 mL solution using the concentration you just calculated.

a. This is equivalent to the number of moles of Fe2+ in your original ore sample.

4. Use the molar mass of Fe to compute the number of grams of Fe2+ in your original ore sample.

5. Divide the number of grams of Fe2+ by the mass of the ore sample (and multiply by 100) to determine the % Fe by mass in your ore sample.

6. Compute the average % Fe by mass and the absolute error (as the range divided by 2). Express your final result as % Fe by mass ± absolute error, rounded appropriately (refer to Experiment 1).

To aid in understanding the calculations, think about:

How many moles of Fe2+ were present in the 50.0 mL of sample that was prepared?

How many moles of Fe2+ were present in the 1.00 mL of solution that was diluted to 50.0 mL?

How many moles of Fe2+ were present in the original 100.0 mL solution prepared from the sample of iron ore?

How many grams of iron does this represent?

In your hand-in, you must show sample calculations associated with making the standard solutions and sample calculations for all steps associated with determining the final result for the % Fe in the ore sample.

# Iron Ore Data Set 33

**Quantitative Analysis of Iron Ore** | Name: | |

Hand-In, Chem 150L, Fall 2020 Due Monday, Nov. 30, 11:59 pm | Partner: | |

1. (2 points) For the iron-phenanthroline complex ion whose absorbance was measured, what color of light does the wavelength of maximum absorbance correspond to? How does this relate to the observed color of the solution? Explain using concepts developed in the “Exploring Color” experiment.

{Type answer here}

2. (4 points) Beer’s Law states that the absorbance of a substance in solution is proportional to its concentration in the solution. Absorbance Concentration describes a straight line whose y-intercept is ideally zero.

a) Include your properly-formatted absorbance vs. concentration graph for the standard solutions below.

{Paste absorbance vs. concentration graph here}

b) How well does your data fit Beer’s Law?

{Type answer here}

c) How does this affect your confidence in your results reported below for the iron ore samples? Explain.

{Type answer here}

3. (10 points)

a) What is your iron ore sample number? _____

b) (3 points) List the % Fe^{2+} by mass found from each of the three samples prepared from the iron ore solution, as well as the average % Fe^{2+} ± absolute error (rounded appropriately) for your iron ore sample.

{List values here}

c) (7 points) Show all calculations done in this experiment below, and briefly state/explain what you are doing in each. You must show one sample calculation for each unique type of calculation. This includes a sample calculation for the concentration of one of the standard solutions made during the first part of this experiment. It also includes one complete set of calculations needed to determine the % Fe^{2+} in the ore (for one of the samples) during the second part of the experiment. In addition, show the calculation of the absolute error. Use Equation Editor. Show units. Be careful with significant figures. “Use Equation Editor” means formatting your calculations with numerators on top and denominators on the bottom where appropriate, not just using the Equation function to write in-line equations like you would in a word processing program.

{Show calculations here}

4. (4 points) What specific steps in procedure or choice of equipment were taken to minimize the error in this experiment? There were many. List and briefly discuss as many as you can.

{Type answer here}

See the attached rubric for more detailed information about grading.

| **Unsatisfactory** | **Borderline** | **Satisfactory** | **Excellent** | **Pts.** |

**Q#1****Explain color**
| Illogical answer or no answer. **0 points**
| Missing one answer. **1 point**
| Minor error. **1.5 points**
| Color of max absor-bance identified, 1 pt. Reasonable explanation of observed color, 1 pt. **2 points**
| 2 pts |

**Q#2****Discuss graph**
| No graph and no answer. **0 points**
| Graph present but poor discussion of quality of data. **2 points**
| Minor errors in graph formatting, -0.5 pt each. **3 points**
| Properly formatted graph, 2 pts. Reasonable discussion of the quality of data and implications for results, 2 pts. **4 points**
| 4 pts |

**Q#3b.****% Fe**
| **0 points** | **1 point** | **2 points** | 3 values listed and correct, 1 pt. Avg ± error listed and correct, 1 pt. Appropriate rounding, 1 pt. **3 points**
| 3 pts |

**Q#3c.****Calcs.**
| Equation editor not used. **0-1 points**
| **2-3 points** | -0.5 pt for minor errors. **4-5 points**
| Calculations shown for standards (1 pt), % Fe^{2+} (2 pts), and error (1 pt). Units (1 pt). Sig. figs. (1 pt). Explanations (1 pt). **6-7 points**
| 7 pts |

**Q#4****Discus-sion of minimi-zation of error**
| No answer given. **0 points**
| Only 1 or 2 examples with reasonable discussion. **2 points**
| Only 2 examples or no discussion. **3 points**
| At least 3 good examples (1 pt each) with reasonable discussion (1 pt). **4 points**
| 4 pts |

| | | | | 20 pts |