Determination of Iron Chemistry Worksheet

Accurately Preparing Solutions of Known Concentration: OnlineStudent Name:
• Carefully read the entire document. Insert values, calculations, and explanations where necessary
• Please be mindful of document formatting when done.
This Lab Report component contains 85 points. The remaining 15 points for this experiment are
earned through completion of the Lab Quiz questions.
Report
Component
Points
Available
Criteria & Maximum Point Deductions
0
–5 for live experiments for incomplete notebook
preparation according to the “Laboratory Notebook
Guidelines” document
10
–7 Title page not present
–5 Missing Results Statement
–3 Results statement lacking experimental findings
–3 Incorrect Title Page format
–5 Overall document formatting is ‘sloppy’
15
–1 to –15 Based on correct insertion of missing values
as determined from video (units, Sig Figs, and
magnitude)
Lab
Calculation
Section
20
–1 to –20 Based on correct calculations (final number,
and some degree of work shown where necessary)
Post-Lab
Questions
20
–1 to –20 Based on Answer Key.
20
Include 1-2 paragraph explanation of the general
principles used in the experiment, how they relate to
each other, the results of the experiments, and
whether they were accurate.
See specific requirements later in this document.
Limited to a maximum 1-2 paragraphs
–1 to –5 Poor grammar, punctuation, sentence
structure or not writing in past tense.
Notebook
Prep
Title Page,
Results
Statement,
and Overall
Formatting of
document
Lab Data
Section/Data
Tables
Discussion &
Conclusion
Points
Earned
NA
(Remote)
Total points (out of 85):
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Accurately Preparing Solutions of Known Concentration: Online
(Remove this, and insert your title page and results statement here)
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Accurately Preparing Solutions of Known Concentration: Online
General Directions: All work should be completed and typed directly into this word document
whenever possible. Alternatively, students may insert images of hand-written work/calculations.
Save the file with your last name in the title (i.e. Student_Solutions_Lab.) Upload the completed
document to Canvas in the correct assignment tab by the due date.
Specific Directions: This lab has three Parts: Part A, Part B and Part C. Complete all data sections by
typing in data after reading the background and watching the appropriate experiment videos carefully.
After reporting necessary data in “Data Tables”, complete the “Calculations and Results Section” that
follow. Write a ~1-2 paragraph “Discussion/Conclusion” for the entire lab This should include a brief
explanation of the general principles (in your own words) used in the experiment and how they relate
to each other, the results of the experiments, and whether they were accurate according to your
calculations.
For this experiment, you must calculate various masses and volumes of reagents. In other words,
wherever the directions say “weigh the required amount” you should have a calculated value.
For Part A1, calculate the theoretical mass of solid copper(II) acetate monohydrate, Cu(CH3COO)2
• H2O (molar mass 199.64 g / mole) needed to make the concentrated solution.
For every other required calculation, use the dilution formula (M1V1 = M2V2) and assume the
molarities are exact as given. Include these numbers in your notebook and put the exact values you use
next to them as you complete the experiment.
OBJECTIVES
1. To prepare solutions of accurate concentration from solids and by dilution.
2. To gain an understanding of significant figures in deciding on which pieces of equipment to use
to prepare the most accurate solutions.
3. To use Beer’s Law to determine the concentration of a solution.
BACKGROUND
In this experiment, you will accurately prepare several standard solutions. Standard solutions are solutions
whose concentrations are accurately known. This usually requires accurately measuring the mass of a
solute and dissolving it in a liquid or diluting a known amount of a solution in a liquid. In Part A of this
experiment you will make a solution of copper (II) acetate from solid and then make a second solution by
diluting part of this original solution. In Part B you will make a potassium permanganate solution by
performing two serial dilutions. In Part C you make a potassium chromate solution using one dilution
Solution preparation always requires measuring the volume of a liquid. The picture below shows various
kinds of glassware that can be used for measuring the volume of solutions. Notice which of these devices
can provide the most accurate measurements. These should be used in this experiment.
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Accurately Preparing Solutions of Known Concentration: Online
volumetric pipet
(high accuracy)
buret
(high accuracy)
graduated cylinder
(medium accuracy)
volumetric flask
(high accuracy)
Erlenmeyer flask
(approximate only)
(not for measuring)
beaker
(approximate only)
(not for measuring )
Figure 1. Common Glassware for Preparing Solutions
When accurate liquid volume measurements are needed, volumes of liquids being transferred from one
container to another are usually measured in volumetric pipets or burets. Volumetric flasks are also highly
accurate, but they are calibrated to contain the specified volume of liquid, not to transfer it.
Beakers and Erlenmeyer flasks are containers. The calibration marks on and can be used to determine how
much liquid they can hold. They are NOT to be used for making approximate volume measurements.
Graduated cylinders are used to measure volumes when high accuracy isn’t needed, but are not
sufficiently accurate for high accuracy volume measurements.
Solutions of accurately known concentration are commonly prepared by one of the following two ways:
Starting with a solid: weigh out accurately a known amount of the solid, dissolving it in enough water to
produce an accurately measured volume of solution in a volumetric flask
Starting with a more concentrated solution: dilute an accurately measured volume of a previously
prepared standard solution to a known total volume of solution in a volumetric flask.
1. Starting with a solid: ( see Figure 2)
– Weigh the required amount of the solid into a small, clean, tared beaker (or weighing dish, if the amount
of solid needed is < 1.000 g ) - Dissolve it in a small amount of distilled water (typically up to ⅓ of the final volume of solution you are trying to prepare - Stir the solid to dissolve as much of it as you can. Some solids dissolve readily in distilled water, others require may include crushing the solid with a glass stirring rod and stirring the solution. - Pour all the solution into a clean volumetric flask of the required size to contain exactly the volume of Page 4 of 14 Accurately Preparing Solutions of Known Concentration: Online solution you will make, being careful not to spill any. - Rinse the beaker, and any remaining solid, into the volumetric flask using several small portions of distilled water from your wash bottle. - Continue rinsing you judge that the entire solid in the beaker has been transferred (rinsed) into the volumetric flask. - Be careful not to use more distilled water than necessary to fill the volumetric flask to about ½ full. - Swirl the solution in the volumetric flask to dissolve any remaining solid and to mix the solution. - Carefully add water to the calibration mark, using a disposable pipet. - Stopper the flask. While holding the stopper, invert the volumetric flask several times to mix the final solution thoroughly Figure 2. Making a standard solution from a solid 2. Starting with a more concentrated solution: ( see Figure 3) - use a volumetric pipet or buret to transfer the required volume of the more concentrated solution into the volumetric flask. - carefully add distilled water to the calibration the mark on the volumetric flask, using a disposable pipet as you get close to the calibration mark. Page 5 of 14 Accurately Preparing Solutions of Known Concentration: Online - when the final volume is achieved, swirl, stopper and mix as described above. Figure 3. Making a standard solution from a more concentrated standard solution measuring/Mohr pipet volumetric pipet Measuring pipets, volumetric pipets, and burets (not shown) can be used to accurately measure volumes of solutions. Volumetric pipets have the greatest accuracy because their diameter is smallest near the calibration mark. However, they can be used to measure only the single volume for which they have been calibrated. They are available in 1, 2, 3, 5, 10, 25, 50, 100 mL, etc. sizes. When a more concentrated solution is used to prepare a more diluted solution, the concentrations of the two solutions are related to the measured volume of the concentrated solution and the final volume of the solution you are preparing. The relationship between the two solutions is given in Equation 1. C1V1 = C2V2 or M1V1=M2V2 (1) This equation is not used for any purpose other than for dilution of a solution. It is sometimes referred to as the dilution equation. Page 6 of 14 Accurately Preparing Solutions of Known Concentration: Online From this relationship, the volume of the concentrated solution required to produce a known volume of a solution can be calculated if the two concentrations are known. Alternatively, if both volumes and the starting concentration are known, the concentration of the final solution can be calculated. Spectrophotometer and Beer’s Law A spectrophotometer is an instrument that measures the amount of light absorbed by the colored components of a solution. It operates by shining light of a specific wavelength through a solution and measuring the difference between the Intensity of light entering and existing the solution. This difference in the light Intensity is the amount of light absorbed by the colored components of the solution. The amount of light absorbed depends on the concentration of the compound in solution. More concentrated solutions appear more intensely colored to your naked eye because they absorb more light. The color of the light absorbed by the colored components of a solution is measured at a specific wavelength, and depends on the distance that the light passes through the solution. The distance that the light passes through a solution is called the cell path length. This relationship between the absorbed light, the cell path length and the colored component concentration is known as Beer’s Law (Equation 2). A = bc (2) where A represents absorbance,  (also written as a) is the molar absorptivity constant, b is the path length and c is the solution concentration. If the wavelength and path length are constant, the amount of light absorbed by the colored component of a solution is proportional to the concentration of the colored species in the solution. This allows us to measure the “absorbance” of a solution to find its concentration. Absorbance has no units. The salts that you will be using contain several ions however, many of them are colorless (K + and CH3COO-). The wavelength that is used for each solution corresponds to the maximum absorbance of the colored ion (Cu+2, MnO4- and CrO4-2). Be sure to record the colors of the solutions. Page 7 of 14 Accurately Preparing Solutions of Known Concentration: Online PROCEDURES Note: Prior to attempting the lab experiment or calculations please view the following videos: How to use a volumetric pipet: https://www.youtube.com/watch?v=Pg9WHKYL8g8&t How to prepare a solution from a solid sample: https://www.youtube.com/watch?v=1iBQXJye8wI Beer’s Law: https://www.youtube.com/watch?v=wYRkw3ThqmA A student prepared several solutions according to the procedure below then measured their absorbances using a spectrophotometer. Part A. Copper (II) acetate solution (AI) Prepare 100.0 mL of ~0.1250 M copper (II) acetate solution from solid copper (II) acetate monohydrate. Refer to background information in making a solution from a solid. 1. Tare a small (~50.0 mL) beaker and weigh the required amount of Cu(CH3COO)2•H2O into it. The exact mass is recorded in the Lab Data section below. Base subsequent calculations on this amount. 2. Add ~25 mL of distilled water to the beaker. Swirl and stir gently with a glass stirring rod in the beaker to help dissolve the solid. This solid dissolves slowly and tends to foam when solving. Minimize the bubbles that may form by minimizing the agitation to the solution. Dissolve as much of the solid as you can in the beaker before transferring it to the volumetric flask. 3. Rinse the beaker with distilled water and add this to the volumetric flask. 4. When all the solid is dissolved and in the volumetric flask, fill it to the mark, stopper it and invert several times to mix the solution. (AII) Prepare 25.00 mL of ~0.05000 M copper (II) acetate solution from your ~ 0.1250 M solution. 1. Pipet the required amount of solution (AI) into a 25.00 mL volumetric flask. 2. Fill it to the mark with distilled water, stopper it and invert several times to mix the solution. 3. Measure the absorbance of this solution at a wavelength of 610 nm, the wavelength where this compound absorbs light most strongly. ( = 5.97 dm3/mol cm) Part B. Potassium permanganate solution (BI) Prepare 100.0 mL of ~0.00100 M KMnO4 solution from the 0.02XXX M KMnO4 stock solution. Record the exact molarity of the stock solution. 1. Take about 20 mL of the stock solution in a small beaker to your bench. 2. Using the appropriately sized volumetric pipet, transfer the required amount of stock solution from Page 8 of 14 Accurately Preparing Solutions of Known Concentration: Online the beaker to the 100.0 mL volumetric flask. 3. Fill it to the mark with distilled water, stopper it and invert several times to mix the solution. (BII) Prepare 25.00 mL of ~4.000x10-4 M KMnO4 solution via dilution of Solution BI (~0.001XX M KMnO4). 1. Using the appropriately sized volumetric pipet, transfer the required amount of your solution from (BI) into a 25.00 mL volumetric flask. 2. Fill it to the mark with distilled water, stopper it and invert several times to mix. 3. Measure the absorbance of this solution at wavelength of 530 nm, the wavelength where this compound absorbs light most strongly. ( = 2.098 x 103 dm3/mol cm) Part C. Potassium chromate solution Prepare ~ 8.75 × 10–4 M K2CrO4 solution from the ~1.25× 10–3 M K2CrO4 stock solution. Record the exact molarity of the stock solution. Calculate the required amount of stock solution to make 10.00 mL of dilute solution. Adjust this value as necessary (ratio) based on the volumetric flask available. 1. Measure the required amount of stock solution from the buret directly into the volumetric flask. Record both the initial and final buret volume readings (including units) to the nearest 0.01 mL in your lab notebook Record both the initial and final buret volume readings (including units) to the nearest 0.01 mL in your lab notebook. 2. Fill it to the mark with distilled water, stopper it and invert several times to mix. 3. Measure the absorbance of this solution at wavelength of 440 nm, the wavelength where this compound absorbs light most strongly. ( = 277.2 dm3/mol cm) For Parts A–C: Calculate the theoretical molarities of all the final solutions using the exact masses, volumes, and molarities of the solutions: Show all work and watch significant figures. Volumetric pipettes and volumetric flasks are accurate to 2 decimal places: 100.00 mL, 25.00 mL, 10.00 mL, etc. Using Beer’s law and the extinction coefficients that are provided, calculate the experimental molarities from the absorbance of each solution. The cell path length (b) is 1 cm. Calculate the % error between your theoretical and experimental molarities. Page 9 of 14 Accurately Preparing Solutions of Known Concentration: Online 15 Points Total. Lab Data Section: Report Sheet (to be submitted with lab report for grading) Data Collection Information Part A ( 1 point Solution AI Mass of Cu(CH3COO)2•H2O weighed (g) 2.446 g Moles of Cu(CH3COO)2•H2O Calculated Molarity of solution AI 1 point Solution AII Volume of Stock Solution AI added via pipet Calculated Molarity of Solution AII Absorbance at 610 nm from Spectrovis: 0.297 Data Collection Part B (This blank table needs to be in your lab notebook before you come to lab.) 2 points. Solution BI Molarity of stock KMnO4 : (read from bottle in lab) 0.01994 M Volume of KMnO4 Stock Solution added via pipet Calculated Molarity of solution BI 2 points M . Solution BII Volume of Stock Solution BI added via pipet Calculated Molarity of Solution BII Absorbance of Solution BII at 530 nm from Spectrovis: 0.847 Page 10 of 14 Accurately Preparing Solutions of Known Concentration: Online Data Collection Part C (This blank table needs to be in your lab notebook before you come to lab.) 1 Point Solution C Molarity of stock K2CrO4 : (read from bottle in lab) 0.00125 M Volume of K2CrO4 stock solution added via buret (mL): Calculated molarity of Solution C Absorbance of ~ 8.75 × 10–4 M K2CrO4 at 440 nm 0.254 The calculated molarities from the tables above are the values of “M (theoretical/calculated)” in the table below. Use Beer’s Law and your absorbance values to calculate the experimentally determined molarities. Obtain Percent errors for each. (8 points) For Completed Table Prepared M Constant, Cell Path M Solution (theoretical/  (dm3/mol Length, b (experimental: calculated) (cm) Beer’s Law) cm) AII. Cu(CH3COO)2 5.97 1.0 BII. KMnO4 2.098 x 103 1.0 CI. K2CrO4 277.2 1.0 Difference in M Percent Error in M Page 11 of 14 Accurately Preparing Solutions of Known Concentration: Online 20 Points Total. Lab Calculation Section General Directions for inserting calculations: A. Either insert photos of your (neatly) written out calculation, Or B. Type out the equations that you are solving. Remember to include proper units! • For this choice, use the “Insert→Equation” function in MS Word. This is a good opportunity to practice typing equations. Show examples of the following calculations from the laboratory procedure below. Use the exact molarities of the stock solutions that were provided in the lab. 2 Points. AI. Calculate the mass of copper (II) acetate monohydrate needed to prepare 100.0 mL of ~0.1250 M solution from solid copper (II) acetate monohydrate. [Remove this statement and enter calculation here] 2 Points. Solution AII. Calculate the volume of concentrated ~ 0.1250 M copper (II) acetate monohydrate solution needed to prepare 25.00 mL of ~0.0500 M copper (II) acetate monohydrate from your stock solution. [Remove this statement and enter calculation here] 2 Points. Solution BI. Calculate the volume of concentrated ~0.02XXX M(use the exact molarity from the lab) M KMnO4 stock solution needed to prepare 100.0 mL of ~ 0.00100 M KMnO4 solution. [Remove this statement and enter calculation here] 2 Points. Solution BII. Calculate the volume of ~0.00100 M KMnO4 (Solution BI) needed to prepare 25.00 mL of ~ 4.000 x 10–4 KMnO4 solution. [Remove this statement and enter calculation here] 2 Points. Calculate the volume of ~1.25× 10–3 M K2CrO4 stock solution needed to prepare 10.00 mL of dilute solution. 8.75 × 10–4 M K2CrO4 solution. [Remove this statement and enter calculation here] 4 Points. Calculate the theoretical (i.e. calculated) molarity of solution AII. [Remove this statement and enter calculation here] 3 Points. Using Beer’s law and the extinction coefficients that are provided, calculate the experimental molarities from the absorbance of solution AII. The cell path length (b) is 1 cm. [Remove this statement and enter calculation here] 3 Points. Calculate the % error between your theoretical and experimental molarities for solution AII. [Remove this statement and enter calculation here] Page 12 of 14 Accurately Preparing Solutions of Known Concentration: Online Discussion/Conclusion Statement (20 points): [Insert Discussion/Conclusion paragraphs here. Discussion and conclusion statements should be standalone paragraphs that anyone could read and quickly understand what you determined, how precise and/or accurate your results were, sources of error, potential improvements to the experiment and applications of the experiment to a broader topic. Some discussion of the topics and how they relate is expected] Discussion/conclusion statements should be stand-alone paragraphs that anyone could read and quickly understand what you determined, how precise and/or accurate your results were, sources of error, potential improvements to the experiment and applications of the experiment to a broader topic. For example, do not simply state that the objective of the experiment was met but rather state what the actual result was. Avoid repeating the procedures from the experiment but be clear when you discuss sources of error and improvements. For example: “It was difficult to get the volume correct” vs “Due to lack of experience it was difficult to use the pipet bulb and the volumetric pipets which may have resulted in significant errors in the volume.” Do not include personal statements about the experiment – I enjoyed this lab, Billy and I finished it quickly, I didn’t understand the procedures…. For this experiment, be sure to address the following at a minimum: (5 points) -What were the solutions made and the % error for each molarity? (8 points) - Which solution was made most accurately/least accurately? Does it “make sense” based on the accuracy of the instruments used (4 points) – Were there any sources of error or improvements to the experiment? How would your sources of error affect your calculated results? (3 points) - Find a real-world example related to any topic covered in this experiment, discuss it briefly and include a citation. Page 13 of 14 Accurately Preparing Solutions of Known Concentration: Online Post Lab Questions (20 points). Answer the following questions. Beer’s Law is applied most accurately when a calibration graph is used. To do so, several solutions of known concentration (the standards) are prepared, and the absorbance of each is measured. A student prepared five standard solutions of copper (II) acetate to determine the molar extinction coefficient of copper (II) acetate at 625 nm and the concentration of an unknown solution copper (II) acetate solution. The student started with a 0.5000 M copper (II) solution and then prepared the other for solutions by dilution. 1. (12 points). Using the data in Table 1, prepare a calibration graph of absorbance (A) vs. concentration using spreadsheet software (MS Excel or Google Sheets). [Insert graph here] Standard Solution 1 2 3 4 5 Table 1 Concentration (M) 0.025 0.050 0.100 0.200 0.250 Absorbance 0.049 0.153 0.242 0.499 0.694 For details on how to plot and fit lines in Excel, see: https://www.youtube.com/watch?v=GfFFpkvvyao For access to MS Office software for Rowan students see: https://irt.rowan.edu/service-catalog/software/downloads/index.html#office • • • • Plot both the experimentally measured absorbance (y-axis) vs. concentration (x-axis). Fit the data with a linear trendline (see video above). Display the equation for the linear line and your R2 value on the graph (see video above). This trendline expresses the mathematical relationship between your absorbance values and concentration, since: y = mx+b A = bc 2. (5 points). What is the molar extinction coefficient at 625 nm for this copper (II) acetate solution? How was it determined? [Insert answer here] 3. (3 points). Can the calibration curve be used to accurately predict the absorbance of 0.0150 M copper (II) acetate solution at 625? Why or why not? [Insert answer here] Page 14 of 14 Determination of Iron by Redox Titration with Permanganate: Remote Student Name: • Carefully read the entire document. Insert values, calculations, and explanations where necessary • Please be mindful of document formatting when done. This Lab Report component contains 85 points. The remaining 15 points for this experiment are earned through completion of the Lab Quiz questions. Report Component Points Available Criteria & Maximum Point Deductions 0 –5 for live experiments for incomplete notebook preparation according to the “Laboratory Notebook Guidelines” document 10 –7 Title page not present –5 Missing Results Statement –3 Results statement lacking experimental findings –3 Incorrect Title Page format –5 Overall document formatting is ‘sloppy’ 15 –1 to –15 Based on correct insertion of missing values as determined from video (units, Sig Figs, and magnitude) Lab Calculation Section 20 –1 to –20 Based on correct calculations (final number, and some degree of work shown where necessary) Post-Lab Questions 20 –1 to –20 Based on Answer Key. 20 Include 1-2 paragraph explanation of the general principles used in the experiment, how they relate to each other, the results of the experiments, and whether they were accurate. See specific requirements later in this document. Limited to a maximum 1-2 paragraphs –1 to –5 Poor grammar, punctuation, sentence structure or not writing in past tense. Notebook Prep Title Page, Results Statement, and Overall Formatting of document Lab Data Section/Data Tables Discussion & Conclusion Points Earned NA (Remote) Total points (out of 85): Page 1 of 10 Determination of Iron by Redox Titration with Permanganate: Remote (Remove this, and insert your title page and results statement here) Page 2 of 10 Determination of Iron by Redox Titration with Permanganate: Remote General Directions: All work should be completed and typed directly into this word document whenever possible. Alternatively, students may insert images of hand-written work/calculations. Save the file with your last name in the title (i.e. Student_Redox_Lab.) Upload the completed document to Canvas in the correct assignment tab by the due date. Specific Directions: Complete all data sections by typing in data after (1) reading the background and (2) watching the appropriate experiment videos carefully. After reporting necessary data in “Data Tables”, complete the “Calculations and Results Section” that follow. Then, write a ~1-2 paragraph “Discussion/Conclusion” for the entire lab This should include a brief explanation of the general principles (in your own words) used in the experiment and how they relate to each other, the results of the experiments, and whether they were accurate according to your calculations. OBJECTIVE 1. To determine the percentage of iron in an unknown solid sample via titration. BACKGROUND Your unknown sample in this experiment will consist of iron (II) ammonium sulfate, Fe(NH4)2(SO4)2•6 H2O and other inert material. Your objective is to determine the percentage of iron (by weight) in the sample: % Fe = # g Fe × 100% # g sample (1) To do this, you will weigh out an amount of the sample and measure the amount of iron in it. You will use a redox titration to measure the amount of iron in the sample. In a titration, a carefully measured amount of one substance (the titrant, in the buret) is added to another substance (the analyte, in a beaker or flask) until exactly enough of the titrant has been added to react with all the analyte (the substance whose amount you are trying to determine). In this experiment potassium permanganate, KMnO4, will be the titrant and Fe2+ from the unknown will be the analyte. The balanced equation for the titration reaction is: 8 H+ + MnO4ˉ + 5 Fe2+ → Mn2+ + 5 Fe3+ + 4 H2O violet/pink (2) colorless By carefully measuring the amount of KMnO4 required to react with the Fe2+ in the sample, you can use the stoichiometry of the balanced reaction to calculate the amount of iron it reacted with. Page 3 of 10 Determination of Iron by Redox Titration with Permanganate: Remote Titrations As with all titrations, the most important point in the titration is the equivalence point. The equivalence point is where the carefully measured amount of the titrant from the buret (in this experiment KMnO4), is exactly enough to react with all the analyte (in this case the Fe2+). That is the only point in a titration where the titrant and analyte are present in the stoichiometric ratio indicated by the titration reaction. The end point of a titration is the point where something changes to indicate that the equivalence point has been reached. This is necessary so that you know when to measure the amount of titrant used. Potassium permanganate, KMnO4, is widely used as an oxidizing agent in redox titrations. In acid solution, MnO4ˉ ion is reduced to Mn2+ as shown in the titration reaction above. Because the MnO4ˉ ion is violet (appearing pinkish in dilute solution) and all other ions and compounds in the titration are nearly colorless, the end point in this titration is when the first permanent pink color (due to unreacted MnO4ˉ) appears in the solution. This titration, which involves oxidation of Fe2+ ion to Fe3+ by permanganate ion, is carried out in sulfuric acid solution to help prevent oxidation of Fe2+ by air. The Fe+3 that is formed has a slight color to however, it forms an essentially colorless complex with phosphoric acid (H3PO4). So, the addition of phosphoric acid near the end point of the titration markedly sharpens the color change for the end point. Page 4 of 10 Determination of Iron by Redox Titration with Permanganate: Remote PROCEDURE A video depicting data collection for the following procedural steps can be viewed here: Analysis of Iron in via Titration with Permanganate: https://www.youtube.com/watch?v=UCReOPYYDU The following experimental procedure was used to generate the data in the tables below. 1. Obtain a magnetic stirrer, stir bar, buret, and an unknown iron (II) sample. Record its number if necessary. Use this sample for all titrations. 2. Prepare the buret: Rinse your buret thoroughly with distilled water and attach it to a ring stand. Check to see that the buret drains freely (no clinging droplets) at least in the calibrated region. If significant droplets cling to the inside of the buret in the calibrated region, use only the calibrated part of the buret below the clinging droplets for your titration. Rinse the buret with a few milliliters of the KMnO4 three times, discarding the waste each time into your waste storage beaker. Drain and then fill the buret with the KMnO4 solution. Drain some KMnO4 making sure to fill the tip with solution before proceeding. Record the initial volume of the solution in your laboratory notebook. 3. Accurately weigh out a 1.2xx ± 0.2 g portion of your unknown sample directly into a clean tared 125mL (or 250-mL) Erlenmeyer flask. (Do not use a watch glass). Record the mass of the sample. 4. Add about 50 mL (measured in a 100-150 mL beaker) of 1.0 M H2SO4 to the iron sample. Add a magnetic stirring bar, and place it on top of the magnetic stirrer. Stir the sample and it should dissolve completely fairly quickly. Ensure the heat is turned off on the stir plate 5. Without much delay (to help prevent air oxidation of Fe2+), record the initial level of KMnO4 solution in the buret, and begin to titrate this iron solution with the KMnO4 solution. 6. When a light yellow color develops in the iron solution during the titration, add about 3.xx mL of 85% H3PO4 from a 10-mL graduated cylinder. (caution: corrosive acid.) 7. Continue the titration until you obtain the first pink color that persists for at least 15 to 30 seconds. Record the final level of KMnO4 solution in the buret in your notebook. Discard your titrated solution into a waste beaker and recover the magnetic stirring bar. Rinse the flask with distilled water and reuse it for the remaining titrations. 8. Repeat the titration with as many samples (up to four) as you have time for in lab. Page 5 of 10 Determination of Iron by Redox Titration with Permanganate: Remote The experimental procedure above was used to generate the data in the tables below. Data Collection Table Titration Mass of Sample (g) Initial Buret Reading (mL) Final Buret Reading (mL) 1 1.234 1.05 19.95 2 1.213 4.00 23.50 3 1.273 0.30 20.05 4 1.292 5.95 25.65 For each titration calculate: -the number of moles of potassium permanganate used in the titration (M x L of KMnO4) -the number of moles of iron present in the sample (stoichiometry in balanced equation) -the percentage by mass of iron in the solid sample (Equation 1). -the average mass % of iron in your unknown. -the standard deviation in your % Fe results. Page 6 of 10 Determination of Iron by Redox Titration with Permanganate: Remote 15 Points Total. Lab Data Section: Report Sheet (to be submitted with lab report for grading) Concentration of KMnO4: 0.01946 M (12 points for correct and complete table) Mass of Volume of KMnO4 Titration Sample (g) Used (mL) 1 1.234 2 3 4 Moles of KMnO4 Moles of Fe % Fe in Sample 1.213 1.273 1.292 Average or Final reported %Fe in sample: __________ [Insert answer here] (3 points) Standard Deviation %Fe:_____________ [Insert answer here] Page 7 of 10 Determination of Iron by Redox Titration with Permanganate: Remote 20 Points Total. Lab Calculation Section General Directions for inserting calculations: A. Either insert photos of your (neatly) written out calculation, Or B. Type out the equations that you are solving. Remember to include proper units! • For this choice, use the “Insert→Equation” function in MS Word. This is a good opportunity to practice typing equations. Show examples of the following calculations below. 4 Points. Calculate the number of moles of potassium permanganate used Trial 1 of the titration. [Remove this statement and enter calculation here] 4 Points. Calculate the number of moles of iron present in the sample used in Trial 1. [Remove this statement and enter calculation here] 4 Points. Calculate the percentage by mass of iron in the solid sample from Trial 1. [Remove this statement and enter calculation here] 4 Points. Using the data for all trials, calculate the average mass % of iron in your unknown. [Remove this statement and enter calculation here] 4 Points. Using the data for all trials, calculate the standard deviation in your % Fe results. [Remove this statement and enter calculation here] Page 8 of 10 Determination of Iron by Redox Titration with Permanganate: Remote Discussion/Conclusion Statement (20 points): [Insert Discussion/Conclusion paragraphs here. Discussion and conclusion statements should be standalone paragraphs that anyone could read and quickly understand what you determined, how precise and/or accurate your results were, sources of error, potential improvements to the experiment and applications of the experiment to a broader topic. Some discussion of the topics and how they relate is expected] Discussion/conclusion statements should be stand-alone paragraphs that anyone could read and quickly understand what you determined, how precise and/or accurate your results were, sources of error, potential improvements to the experiment and applications of the experiment to a broader topic. For example, do not simply state that the objective of the experiment was met but rather state what the actual result was. Do not repeat all the procedures from the experiment but be clear when you discuss sources of error and improvements. For example: “It was difficult to get the volume correct” vs “Due to lack of experience it was difficult to use the pipet bulb and the volumetric pipets which may have resulted in significant errors in the volume.” Do not include personal statements about the experiment – I enjoyed this lab, Billy and I finished it quickly, I didn’t understand the procedures…. (4 points) -what was the final % of iron in your sample (4 points) -discuss accuracy using percent error (4 points) deviation - discuss precision of your results (were all your values close together or not) using standard (4 points) -discuss sources of error and improvements to the experiment (4 points) - find a real-world example related to any topic covered in this experiment, discuss it briefly and include a citation. Page 9 of 10 Determination of Iron by Redox Titration with Permanganate: Remote Post Lab Questions (20 Points). Answer the following questions. A solid sample containing some Fe2+ weighs 2.360 g. It required 36.44 mL 0.0244 M KMnO4 to titrate the Fe2+ in the dissolved sample to a pink end point. The balanced redox reaction is shown below. 8 H+ + MnO4ˉ + 5 Fe2+ → Mn2+ + 5 Fe3+ + 4 H2O Calculate each of the following quantities. A. (4 points) How many moles MnO4– were used in the titration? [Remove this statement and enter calculation here] B. (4 points) How many moles Fe 2+ are there in the sample? [Remove this statement and enter calculation here] C. (4 points) How many grams of Fe are there in the sample? [Remove this statement and enter calculation here] D. (4 points) What is the percentage of Fe (by mass) in the sample? [Remove this statement and enter calculation here] 2. (4 points) Calculate the percentage of Fe in pure iron(II) ammonium sulfate hexahydrate: (NH4)2[Fe(SO4)2]• 6 H2O. [Remove this statement and enter calculation here] Page 10 of 10