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Homework answers / question archive / Only answer laboratory report questions mentioned below the experiment 1 CHEM 25415 Instrumental Analysis Laboratory Experiments Experiment 1: Quantitative Analysis of Iron in Well Water by External Standard Method using a Visible Spectrometer Quantitative Analysis of Iron in Well Water by External Standard Method using a Visible Spectrophotometer   Objectives   The student learn how to operate the Genesys20 visible spectrophotometer by performing a colorimetric experiment for the determination of iron in well water

Only answer laboratory report questions mentioned below the experiment 1 CHEM 25415 Instrumental Analysis Laboratory Experiments Experiment 1: Quantitative Analysis of Iron in Well Water by External Standard Method using a Visible Spectrometer Quantitative Analysis of Iron in Well Water by External Standard Method using a Visible Spectrophotometer   Objectives   The student learn how to operate the Genesys20 visible spectrophotometer by performing a colorimetric experiment for the determination of iron in well water

Chemistry

Only answer laboratory report questions mentioned below the experiment 1

CHEM 25415

Instrumental Analysis Laboratory Experiments

Experiment 1: Quantitative Analysis of Iron in Well Water by External Standard Method using a Visible Spectrometer

Quantitative Analysis of Iron in Well Water by External Standard Method using a Visible Spectrophotometer

 

Objectives

 

The student learn how to operate the Genesys20 visible spectrophotometer by performing a colorimetric experiment for the determination of iron in well water. An external standard calibration graph following Beer’s Law will be plotted at the wavelength corresponding to the absorbance maximum (λmax). Multipoint regression analysis and single point method will be used to quantify the iron concentration (mg Fe/ 1.0 US gallon jug) in well water. Statistical data handling methods will be employed to evaluate the quality of the final results. The student will determine the identity of the well water by comparing their final experimental results with those posted on Slate.  

 

Introduction

 

I- Absorption of Light

 

When a chemical species such as an organic molecule absorbs a photon, the energy of the species is increased and the species is promoted to a high energy excited state (E1) as illustrated in Figure 1.

 

 

                                                 

                                 

 

E

1

 

 

 

 

 

                                                                                             

????

=

?

??

=

 

E

1

 

 

E

0

 

           

PHOTON

 

 

 

 

                                                  

                                 

 

E

0

 

Figure 1: Excitation of an Analyte by UV

-

visible Light

 

 

 

If a chemical species emits a photon, its energy is lowered and the species is said to return to its ground state (E0 - lowest energy state).

 

When incident light (Po) is absorbed by an analyte species in solution, the radiant power of the incident beam of light (P) decreases. Radiant power refers to the energy per second per unit area of the light beam. Figure 2 shows a schematic diagram of light absorption by an analyte species in solution with concentration “c”.

 

 

 

 

 

Cuvette containing absorbing analyte with concentration “C”

 

 

 

 

                                           

 

 

 

 

 

 

b = path length   

 

                                                                     

Figure 2: Absorption of

Light by an Analyte Species in Solution

 

P

o

 

P

 

 

b

 

 

 

Incident light of all wavelengths emanating from a continuous radiation source is passed through a monochromator (a grating and series of mirrors) in order to select one wavelength. Monochromatic light of a particular wavelength, with incident radiant power (Po) strikes analyte in solution with a cuvette path length (b) and the resulting radiant power is attenuated (P) emerging from the other side of the cuvette due to absorption of light by the analyte, such that P < Po. The transmittance (T) is defined as the fraction of the original light that passes through the analyte solution.

    

                                                             

 

Transmittance has a range of zero to one, and by multiplying transmittance by 100, percent transmittance (%T) ranges between 0-100%. 

 

                                                       

If light is not absorbed by the solution relative to the blank, then the %T is 100% (P = Po) and absorbance is zero (A=0). If all of the light is absorbed by analyte in solution then none is transmitted to the detector causing %T to be 0. 

 

II- Beer Lambert Law 

 

The laws of Lambert, Bouger, and Beer state that at a given wavelength (monochromatic light- λmax ), the proportion of light absorbed by a transparent medium is independent of the intensity of the incident light and is proportional to the number of absorbing analyte species through which the light passes. The combined Beer-Lambert Law may be expressed mathematically as:

 

 

where         Po - power or intensity of incident light                    P -  power or intensity of transmitted light                    T -  transmittance of the analyte solution and   A -  absorbance at λmax

       ε  -  molar absorptivity or molar extinction coefficient  (L cm-1 mol-1)             b  -  path length of cell which light passes through (cm)             c  - concentration of analyte species in solution (mol/L)  

It can be seen that absorbance is directly proportional to the analyte concentration (c) and can be expressed in moles/L with the molar absorptivity (ε) having units of “L cm-1 mol-1”. Molar absorptivity is a constant for a particular absorbing analyte in a particular solvent at a particular wavelength (λmax). The cuvette width or pathlength (b) is also constant for a particular analysis and with units of “cm”. Thus, if the analyte concentration doubles, the absorbance will also double and if the concentration is cut by half then absorbance will be cut by half. 

 

One basic assumption when applying Beer's Law is that monochromatic light is used. In experimental situations, this is not the case, but rather a band of radiation is passed (bandwidth) the width of which depends on the dispersing grating and exit slit width. The absorption spectrum of an absorbing analyte solution as seen in Figure 3 shows that different wavelengths are absorbed to different degrees; that is, the molar absorptivity changes with wavelength. At a wavelength corresponding to a fairly broad maximum on the qualitative absorption spectrum, the band of wavelengths will all be absorbed to nearly the same extent. The Genesys20 spectrophotometer bandwidth spans 8 nm compared to more sophisticated instruments that range from 0.5-4.0 nm.

 

 

 

 

Figure

3

:

 

D

etermination

of W

avelength

 

Maximum

 

 

 

The band of wavelengths (or bandwidth) emerging from the exit slit (monochromatic wavelength value set on the instrument) produce large variations in the absorbance value if measurements are recorded on the shoulder of an absorption peak as illustrated in Figure 4 – Band B. When wide bandwidths are used for quantitative analysis, the Beer’s Law calibration plots curve dramatically due to a departure from Beer’s Law. 

 

 

 

 

 

 

 

 

 

                    

 

Figure

4

:

 Deviations from Beer's

L

aw: Band A

-

 

No Deviation, Band B

-

Negative Deviation

 

 

  

Before determining the concentration of an analyte species in a sample solution using visible spectrophotometric methods, it is necessary to find a suitable wavelength band where deviation from Beer's Law will be almost negligible (i.e. broad absorbance maximum).  To achieve this goal, the analyst must first run a qualitative spectral plot of absorbance as a function of changing wavelength to determine the absorbance maximum.  The visible absorbance spectrum is unique for each absorbing analyte species and in the case of the ferrous 1,10-phenantroline complex shown in Figure 5, the absorbance maximum appears at approximately 508 nm.

 

 

                

 

 

Figure

5

:

 Qualitative Visible Spectrum of

Ferrous

1

,

10

-

phenantroline

 

 

 

 

Beer's Law plots are prepared by measuring the light absorbed at 508 nm by a series of iron external standard solutions of known concentration as shown in Figure 6. In this case, the analyte will be iron (II) complexed with 1,10-phenantroline to produce an orange colour. The absorbance of each solution will be measured and a linear calibration curve can be constructed as illustrated in Figure 6 exhibiting the equation:  

                                                               ??=????+??

 

 

 

 

 

Figure

6

:

 A Beer's Law Plot

-

Dashed Line Indicates Concentration

 

 

 

 

To determine the concentration of iron in the unknown well water sample, the sample must be diluted, acidified and complexed with excess 1,10-phenantroline. The absorbance of the resulting orange solution will be measured and the concentration determined by rearranging the linear calibration equation:

                                                       

 

When it is impossible to prepare a series of external standard calibration solutions due to time constraints, it is possible to estimate the analyte concentration in an unknown sample using a mathematical single-point approach. One only needs to prepare one single calibration standard of known concentration in addition to the unknown sample solution. The absorbance values of both solutions are measured and by using the Beer Lambert Law single point equation, the analyte concentration in the unknown sample can be calculated:

 

 

 

In order to use this method, it is assumed that the linear dynamic range is known and that the standard used for the calculation falls within that range. The major drawback of the single point approach is that one single calibration solution is used instead of a series of external standard solutions to evaluate the analyte concentration in an unknown sample. If the single standard is not prepared accurately, the calculated analyte concentration will be erroneous and accuracy is lost.

 

III-Thermo Scientific Genesys20 Visible Spectrophotometer

 

One of the most common ways to measure absorbance of visible light is to use a low cost single beam spectrophotometer shown in Figure 7 called a Genesys20. Within this instrument, the light source is an ordinary tungsten halogen lamp whose emission covers the entire visible spectrum, extending somewhat into the ultraviolet and infrared regions (325 - 1100 nm). The lamp is mounted very close to the entrance slit of the monochromator to ensure continuous energy output. An optical stop situated between the entrance slit and the turning mirror reduces the amount of stray light in the instrument. Diverging light projected onto the turning mirror is directed to the main mirror which in turn reflects the diverging light to parallel beams onto the grating.

 

 

 

Figure

7

:

 Schematic Diagram of the Optical System of a

Genesys20

Spectrophotometer

 

 

 

The Czerny-Turner type monochromator uses a 1200 line/mm reflective grating as a wavelength dispersing element. The grating disperses the collimated light into its component wavelengths back to the main mirror where it is reflected the exit slit. When white light falls on the reflection grating it is dispersed into a fan of light beams with the short wavelengths (visible (violet) - UV) at one end and the longer wavelengths (red and infrared) at the other end. Figure 8 illustrates the how the overall spectrophotometer disperses the polychromatic source radiation into component wavelengths with the reflection grating which in turn focuses the monochromatic light onto the exit slit of a fixed bandwidth.  The exit slit acts as a blocking filter allowing only a certain bandwidth of wavelengths to pass onto the sample cuvette while blocking the rest of the entire visible spectrum. The grating position determines which band of wavelengths emerge from the exit slit.

 

Any radiation not absorbed by the analyte solution (transmitted radiation) falls on the solid state photodiode array detector located in the front section of the instrument immediately following the sample compartment. The mounting plane of the detector is angled with respect to the incoming light beam in order to minimize back reflections into the sample compartment. The detector converts the incoming light signal to an electrical signal which is then amplified. The readout device registers either percent transmittance or absorbance on an LED display.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Figure

8

:

 Diagram Il

l

ustrating how the Exit Slit is used to Select the Desired Wavelength

 

Entrance

Slit

 

Stop

 

GRATING:

 

 

Micro

-

stepping

motor allows one

set

to

wavelength. 

Grating

sends

horizontally

dispersed

polychromatic

collimated

 

light

back to the main

mirror

.

 

TURNING MIRROR:

 

to

Directs

diverging

beam

main mirror.

 

MAIN MIRROR:

 

Converts

diverging beam to

parallel

light

and

directs

it

to

the

grating.

 

MAIN MIRROR:

 

Main mirror directs

selected

from

wavelength

grating

to

exit

the

slit.

 

Exit Slit

 

MONOCHROMATIC

RADIANT ENERGY

 

FALLS ON SAMPLE

 

nm

800

 

4

00

nm

 

The stop is a filter that

of

amount

reduces

the

stray

the

in

light

monochromator.

 

Source

 

 

 

 

IV- Iron in Water Analysis

 

Iron present in natural waters is usually in the form of ferric or ferrous salts and can be attributed

to the weathering of rocks and minerals, acidic mine water drainage, landfill leachates, sewage effluents and iron-related industries. The iron concentration in drinking water is normally below

1.0 ppm where most water treatment plants remove insoluble iron, the major form found in aqueous environments. The established Canadian guidelines for iron in drinking water is less than 0.3 ppm.  

In this experiment, a rapid and simple spectrophotometric method for the determination of ferrous iron Fe2+ with the ligand 1,10-phenanthroline will be studied. Pure iron does not absorb light in the visible region (360 - 900 nm) of the spectrum. In order to use the Genesys20 spectrophotometer to determine the maximum absorbance wavelength for quantification, iron must first be treated with 1,10 -phenanthroline to form an orange-red complex that does absorb  light. 1,10-phenanthroline has two pairs of unshared electrons that can be used to form coordinate covalent bonds creating a coordinate covalent complex (brightly colored with central metal atom).

 

The complexation reaction forming an orange

-

red species is described by the equation:

 

 

 

 

 

 

 

Fe

2+

 

 

+

   3 PhenH

+

  

à

 

Fe(Phen)

3

2+

  

+

   3 H

+

 

 

 

    

Iron (II) +

 

 

(

3

10

1

,

-

phenanthroline

)

 

à

 

Iron(II)

-

10

,

1

-

phenanthroline complex

 

 

                                                

 

Figure 9: Iron(II)

-

1

10

,

-

phenanthroline complex

 

 

1

,10

-

 

phenanthroline will not react with Fe

3

+

 

and

must

first

be treated with hydroxylamine to reduce

 

Fe3+ to Fe2+. For the complex to form, iron must be in the +2 oxidation state The hydroxylamine will keep iron in the +2 state. Since dissolved oxygen in water can oxidize Fe2+ to Fe3+ an excess amount of hydroxylamine reducing agent is added to ensure iron remains as Fe2+. Sodium acetate buffer solution is also added to the blank, standards and unknown solutions to maintain the pH to 3.5 in order to ensure that Fe2+ does not oxidize to Fe3+.

 

The 1,10-phenanthroline complex of iron(II) illustrated in Figure 9 is an example of a chargetransfer complex, which consists of an electron-donor group bonded to an electron acceptor. When the complex absorbs radiation, an electron from the donor is transferred to an orbital that is largely associated with the acceptor. The excited state is thus the product of a kind of internal oxidation/reduction process. In this complex the metal ion serves as the electron donor. Charge-transfer absorption by complexation is important for spectrophotometric quantitative analyses because molar absorptivities of these types of complex ions are usually large (εmax > 10,000 L/mg * cm), a circumstance that leads to high sensitivity.

 

 

 

 

 

 

 

 

 

 

 

 

Prelab Questions

  1. Calculate the theoretical concentration of Fe2+ in ppm (mg/L) if 0.734 grams of Fe(NH4)2(SO4)2 * 6H2O salt was weighed on an analytical balance and dissolved in a 1000.0 mL volumetric flask which was made to the mark.
  2. Using the Fe2+ stock solution, calculate the theoretical concentration in ppm when 10.00 mL of stock solution (question 1) is transferred to a 100.00 mL volumetric flask and made to the mark.
  3. Using the Fe2+ substock solution (question 2), calculate the theoretical concentrations in ppm for all Fe 2+ external calibration solutions when 1.000, 000, 5.000, 7.000 and 9.000 mL is transferred to five 25.00 mL volumetric flasks respectively.
  4. Describe the preparation of the blank solution. What is the role of each chemical in the blank? Explain why it is not RO water.
  5. Explain why a transfer pump rather than a volumetric pipette can be used to add 1,10phenanthroline, sodium acetate and hydroxylamine chloride solution to all the standards, unknowns and blank solutions
  6. Explain why 1,10-phenanthroline is added to the external calibration solutions as well as the unknown solutions.
  7. Explain the purpose of adding hydroxylamine and sodium acetate buffer to the external calibration solutions as well as the unknown solutions.
  8. For the unknown sample, determine the dilution factor that you are diluting the unknown to.
  9. State the analysis wavelength used for this experiment.

 

  1. Set up tables in your lab hardcover lab notebook in order to record the data collected for this experiment. See Table 1- step 2 procedure for formatting.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Procedure: 

 

NOTE: The Genesys 20 spectrophotometer must be turned at the beginning of the experiment to warm-up.

NOTEThe diluent for this experiment is RO water.  

 

A – Preparation of Fe2+Substock Solution

1. To prepare an Fe2+ substock solution, using a 10.00 mL volumetric pipette, analytically transfer 10.00 mL of Fe2+ stock solution (prepared by the technologist) into a 100.00 mL volumetric flask. Make up to the mark with RO water. 

NOTE: The stock Fe2+ solution was prepared from primary standard grade Fe(NH4)2(SO4)2 * 6H2O salt with RO water as the diluent. Record the actual mass of iron salt weighed and volume of the volumetric flask used to make the stock solution (check the stock solution bottle label) in your hardcover notebook.  Calculate the concentration of Fe2+ in stock and substock solutions using the proper significant figures. 

B – Preparation of External Standard Fe2+ Calibration Solutions  

  1. Prepare 5 calibration external standard Fe2+solutions by using either volumetric pipettes or a 5.000 mL micropipette to transfer the required volumes of Fe2+ substock solution shown in the Table 1, into 5 pre-rinsed 25.00 mL volumetric flasks. 

 

NOTE: DO NOT DILUTE YET!

 

Table 1: Fe 2+ External Standard Calibration Solution Preparation

Standard #

Fe2+ Substock Solution

Volumes pipetted (mL)

Fe2+  Calibration Standard

Concentrations ( ppm)

Absorbance

1

1.000

 

 

2

3.000

 

 

3

5.000

 

 

4

7.000

 

 

5

9.000

 

 

 

  1. Before making each calibration solution to the mark with RO water, add the following colourimetric reagents using the transfer pump apparatuses found in the fumehood to each 25.00 mL volumetric flask in step.  Follow the precise order of reagents as listed

              1.0 mL hydroxylamine chloride solution

            2.5 mL 1,10-phenanthroline solution

            4.0 mL sodium acetate solution  Dilute each calibration solution to the mark with RO water. Stopper invert and shake to ensure adequate mixing. Allow 5 minutes for the colour to develop before running the solutions using the spectrophotometer.

  1. To prepare a blank solution, add each of the colourimetric reagents, Step 3- procedure to a separate pre-rinsed 25.00 mL volumetric flask. Fill the flask to the mark with RO water. Stopper invert and shake to ensure adequate mixing.

NOTE: The blank contains all the reagents used in the analysis except the analyte.  

C - Sample Preparation of the Unknown Well Water Sample - Triplicate 

  1. Obtain an unknown well water sample from the professor. Record the sample number in your hardcover notebook.
  2. Pipette using a volumetric or micropipette 2.000 mL of well water sample into 3 separate pre-rinsed 25.00 mL volumetric flasks (triplicate solution).
  3. To each of the three 25.00 mL volumetric flasks, add the following reagents in the precise order using the transfer pump apparatus located in the fumehood:

1.0 mL hydroxylamine chloride solution

2.5 mL 1,10-phenanthroline solution

4.0 mL sodium acetate solution

 

  1. Dilute each unknown well water solution to the mark with RO water. Stopper invert and shake to ensure adequate mixing.
  2. Allow 5 minutes for the colour to develop before measuring the absorbance values for each of solutions using the spectrophotometer. The Genesys20 visible spectrophotometer must be turned on for at least ten minutes before any measurements are recorded.

D – Absorbance Measurements of the External Standards and Unknown Sample Solutions

 

  1. Obtain one cuvette and a cuvette rack. Using a Sharpie, mark a reference line on the cuvette. The standard operating procedure for the Gensys 20 spectrophotometer will be found next to the instrument.

 

  1. Analytically rinse the cuvette with the blank solution. Fill this cuvette 2/3 full with the blank solution. Use this solution to set 100 %T on the Genesys20 spectrophotometer at 508 nm.  

 

NOTE: Use a Kimwipe to wipe the cuvette prior to insertion into the sample compartment. Insert the cuvette into the sample compartment the same way each time, with the reference line aligned to the mark on the sample compartment.

 

  1. Remove the cuvette and dump the blank solution into a waste beaker. Now set the spectrophotometer to Absorbance mode.

 

 

 

  1. Analytically rinse the cuvette with standard solution #1 three times. Fill the cuvette 2/3 full with standard solution #1 and measure the absorbance. Record your results in your lab notebook. Repeat this step for each of the Fe 2+ external standard solutions.

 

  1. Once the absorbance measurements have been made for external standards #1-5, analytically rinse the cuvette with the well water sample #1 three times. Fill the cuvette 2/3 full and record the absorbance. Repeat this procedure for each replicate unknown well water sample solution.

 

  1. Rinse the cuvette with RO water and return the cuvette to the front bench. 

 

  1. Shut down the spectrophotometer using the power switch at the back of the unit. Unplug the power cord.

 

E - Disposal of External Standard and Sample Solutions

 

17. Discard all contents of the waste beaker including the substock, external standards and sample solutions into the waste container designated by the technologist. The waste will be disposed of as hazardous waste by the technologist.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Laboratory Report Questions

 

A – Fe2+  External Standard Solution Preparation

 

  1. a) Using the mass in grams of Fe salt weighed by the technologist and flask size that was used in step 1 - procedure, calculate the actual Fe2+ stock solution concentration in ppm. Show the calculation.

 

b) State the analysis wavelength used and the unknown sample number.

 

  1. Calculate the actual Fe2+ substock solution concentration in ppm prepared in step 1 - procedure. Show the calculation. 

 

  1. Calculate the actual Fe2+ external standard solution concentration (ppm) for all external standard solutions prepared in steps 2 - procedure. Show one sample calculation for standard solution # 1 only. Record the Fe2+ concentrations in ppm and corresponding absorbance values in table format

 

  1. Construct a computer generated diagram illustrating the preparation of the Fe2+ stock, substock and external calibration standard solutions. Include the volumes pipetted, volumetric flask sizes, reagents added and actual (calculated) concentrations of all solutions. 

 

  1. Record the absorbance values for the replicate unknown well water samples in a new table. 

 

  1. Construct a computer generated diagram illustrating the preparation of the well water samples. Include the unknown number, volumes pipetted, volumetric flask sizes and reagents added. 

 

B – Multipoint Method

 

  1. Using EXCEL, construct a calibration curve of absorbance versus actual Fe2+ concentration in ppm. The linear regression equation and the correlation coefficient must be shown on the graph. 

 

NOTE: When creating the Excel graph, you must follow all of the graphing rules outlined in the lecture slides and in the lab techniques course. 

 

  1. Using the linear regression equation, calculate the Fe2+ concentration in ppm in the diluted unknown triplicate well water samples. Show one sample calculation for replicate sample #1. Record your calculated results in a new table format. 

 

  1. Since a dilution was made to the well water sample prior to measurement, back calculate the the Fe2+concentration in ppm for the original well water sample. Show one sample calculation for replicate sample #1. Record your calculated results in the same table as step 8 - report. 

 

  1. Assuming that your sample came from a 1.0 US gallon jug, calculate the Fe2+ concentration in mg/ 1.0 US gallon jug  for the original well water sample. Show one sample calculation 

 

 

for replicate sample #1. Record the final concentration (mg Fe2+ / jug) in the same table as step 8 - report. 

 

  1. Calculate the mean Fe2+ concentration (mg Fe2+ / jug) in in the original well water sample absolute standard deviation, relative standard deviation and true value (µ) at a 95%  confidence interval. Show one sample calculation for each statistic and include a statement about the meaning of your µ value. Record the statistical results in a new table

 

NOTE: Before calculating any statistics, Q-test, at 95% confidence, (only if necessary- real outlier) any suspect datum (show your work, if performed). 

 

C – Single Point Method

 

  1. Compare the external standard absorbance values to those for the replicate unknown samples. Select and state the name of the external standard solution whose absorbance value is closest to that of the unknown replicate samples. Using the selected standard and the following equation, calculate the Fe2+ concentration in ppm in each diluted unknown well water sample.

                                                             

 

Show the calculation for replicate sample #1. Record calculated single point method results for each replicate sample in a new table. 

 

  1. For each replicate sample, repeat steps 10-11 - report to calculate the Fe2+concentration in (mg Fe2+ / jug) in the original well water sample. Show one sample calculation for replicate sample #1 only. Record your results in the same table as step 12 - report. 
  2. Calculate the mean Fe2+concentration (mg Fe2+/ jug) for the original well water sample, absolute standard deviation, relative standard deviation and true value (µ) at a 95% confidence interval. Before calculating any statistics, Q-test any suspect datum. Record your statistical results in a new table. 

 

NOTE: DO NOT SHOW ANY SAMPLE CALCULATIONS STATISTICS HERE. 

 

D – Summary of Results and Unknown Identification

 

  1. In a new final summary table, compare the the mean Fe2+ concentration (mg Fe2+/ jug) for the original well water sample absolute standard deviation, relative standard deviation and true value (µ) at a 95% confidence interval for the multipoint and single point methods. Results must be summarized in one table. 

 

  1. Using the multipoint method results only, create a table comparing your experimental mean   Fe2+concentration (mg Fe2+/ jug) result with all of the true values for the possible well water samples posted on SLATE. State the identity of your unknown well water sample. 

 

 

 

 

  1. Calculate the relative error for the mean Fe2+concentration (mg Fe2+/ jug) in the well water sample for both single point and multipoint methods. Show one sample calculation for the multipoint method only. Report your results in table format. 

 

 

E – Discussion (14 marks)

 

  1. Suppose that your partner zeroed the Genesys 20 spectrophotometer used tap water instead of the blank prior to making absorbance measurements for the 5 calibration standards and samples. Explain how this error would affect the accuracy of your final result. (2)

 

  1. Discuss whether or not your multipoint final results agree with your single point final results. Support your answer with numerical results from the table prepared step-15 results. (3)

 

  1. Suppose that when you pipetted standard #2 solution, you unknowingly pipetted 2.000 mL into 25.00 mL instead of 3.000 mL. When performing the calculation for the Fe concentration for this standard, you used the correct volume- 3.000 mL. (6)

 

    1. Explain how this error would affect the calibration equation and the correlation coefficient assuming that you used all five standards that you prepared.

 

    1. Explain how this mistake would affect the accuracy and precision of your final result obtained with this calibration plot. 

 

    1. Explain two changes that could be made in order to accurately determine the iron concentration in the unknown samples. 
  1. If the unknown well water samples were analyzed at 508 nm but the external standard solutions were analyzed at 550nm, fully explain how this would influence your mean  Fe2+concentration in the unknown well water sample. (3)

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