School of ChemistryPriestley Laboratory

One-Day Practical Coursefor Year 12 Students

Inorganic Experiments

www.chem.leeds.ac.uk

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COSHH Assessment

Section: Priestley Laboratory Experiment: Preparation and Analysis of Ammonium Iron (II)

Sulfate

Personnel Involved: Year 12 students, laboratory demonstrators, laboratory staff

Chemical Hazard Precautions Disposal

Iron (II) sulfate heptahydrate Harmful if swallowed

Irritating to eyes and skin

Wear gloves

Avoid breathing

dust

In a fumehood

dilute with

water and flush

down sink

Sulfuric acid (1 M) Corrosive Wear gloves In a fumehood

dilute with

water and flush

down sink

Ammonium sulfate Not classified as a hazardous substance In a fumehood

dilute with

water and flush

down sink

chemical waste

Ammonium iron (II) sulfate

hexahydrate

Irritating to eyes skin and

respiratory system

Wear gloves

Avoid breathing

dust

In a fumehood

dilute with

water and flush

down sink

Iron (III) chloride solution

(0.05 M)

Harmful

Corrosive

Wear gloves In a fumehood

dilute with

water and flush

down sink

Ammonia solution (2 M) Causes burns

Irritating vapour

Toxic to aquatic organisms

Wear gloves

Avoid breathing

vapour

In a fumehood

dilute with

water and flush

down sink

Potassium manganate(VII)

solution (0.02 M)

Oxidising agent

Toxic to aquatic organisms

Wear gloves In a fumehood

dilute with

water and flush

down sink

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Chemical Hazard Precautions Disposal

1,10-Phenanthroline

solution (0.006 M)

Toxic if swallowed

Toxic to aquatic organisms

Wear gloves

Avoid breathing

vapour

In a fumehood

dilute with

water and

flush down

sink

Potassium hexacyanoferrate

(II) solution

Harmful

Contact with acid liberates

toxic gases

Wear gloves Metal waste in

waste

fumehood

Potassium hexacyanoferrate

(III) solution

Harmful

Contact with acid liberates

toxic gases

Wear gloves Metal waste in

waste

fumehood

Hydroxylamine

hydrochloride

Harmful in contact with

skin and if swallowed

Irritating to eyes and skin

Potential carcinogen

May cause sensitisation

Toxic to aquatic organisms

Wear gloves

Avoid breathing

vapour

Acetone waste

Sodium ethanoate trihydrate Not classified as a hazardous substance In a fumehood

dilute with

water and

flush down

sink

Other Hazards None

Emergency action If chemicals come into contact with eyes or skin,immediately wash thoroughly with copious water.

Report incident to laboratory demonstrator.

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Schedule10:30 – 13:00 Preparation of Mohr’s salt and test tube reactions13:00 – 14:00 Lunch14:00 – 15:30 Analysis of Mohr’s salt

Preparation and Analysis of Mohr’s Salt: Ammonium Iron (II) Sulfate Hexahydrate

Reaction Overview:

The element iron forms many compounds containing Fe2+ or Fe3+, also written as Fe(II) or Fe(III).

Mohr’s Salt (ammonium iron(II) sulfate hexahydrate) is often employed as an analytical standard,

and has been used in a variety of other applications from nanomaterials to general redox reactions.

In this experiment you will mix solutions of iron(II) sulfate and ammonium sulfate to obtain pale

blue-green crystals of Mohr’s salt.

Mohrs salt is an example of a ‘double salt’, i.e. it contains the same ions as in FeSO4.7H2O and

(NH4)2(SO4). It is simple to prepare because it is less soluble than either of its constituent salts. A

double salt in solution displays the chemistries of its component ions because no new chemical

bonds are formed.

The salt is named after the German chemist Karl Friedrich Mohr, who made many important

advances in the methodology of titration in the 19th century. It is preferred over iron(II) sulfate for

titrations as it is less prone to oxidation by air. The oxidation of iron(II) to iron(III) occurs more

rapidly at higher pH; Mohr’s Salt lowers the pH of solutions slightly.

Objectives

The aims of this experiment are:

1. To prepare Mohr’s salt, ammonium iron(II) sulfate hexahydrate,

2. To investigate some aqueous reactions of iron(II) and iron(III) compounds.

3. To analyse your compound by a redox titration with potassium manganate(VII).

4. To analyse your compound by spectrophotometry.

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Part 1: Preparation of Mohr’s salt

The method described here involves the simple mixing of solutions of FeSO4.7H2O and (NH4)2(SO4),

followed by an evaporation, a crystallisation, and finally a filtration.

For this part of the experiment, you should work in a fumehood.

Place a clean 100 mL beaker onto a top-pan balance, and set the display to read zero; this is known

as taring the balance. Weigh directly into the beaker iron(II) sulfate heptahydrate (FeSO4.7H2O;

12.5 g) from the reagent bottle. Remove the beaker from the balance and carefully add dilute

(1 M) sulfuric acid (5 mL), and deionised water (10 mL). Heat the beaker on a stirrer hotplate,

stirring with a magnetic follower until all of the solid has dissolved. You should obtain a clear, pale

green solution.

Place a second clean 100 mL beaker onto a top-pan balance and tare the balance. Weigh directly

into the beaker ammonium sulfate [(NH4)2(SO4); 6 g]. Remove the beaker from the balance and

add deionised water (8 mL). Heat the beaker on a stirrer hotplate, stirring with a magnetic

follower until all of the solid has dissolved. You should obtain a colourless solution.

Carefully pour the contents of the beaker containing the solution of iron(II) sulfate into the beaker

containing the ammonium sulfate solution. The combined volume should be approximately 30 mL.

Heat the beaker to boiling until the initially opaque green solution becomes clear.

Note: Do not let the beaker boil dry.

The final solution should be a blue-green colour, and contains your product. You will isolate your

product by crystallisation, one of the most important purification techniques in synthetic

chemistry. Cool your beaker first to room temperature, and then in an ice-water bath.

Note: It is a good idea to gently clamp your beaker in place in the cooling bath to ensure it does

not tip over, losing your product.

1. Calculate the amount, in moles, of 12.5 g of FeSO4.7H2O. The relative atomic masses

are: H (1), N (14), O (16), S (32), and Fe (56).

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2. Determine the theoretical (maximum) yield (in g) of (NH4)2Fe(SO4)2.6H2O.

You should complete Part 2 of the experiment, ‘Test-Tube Reactions of Iron(II) and Iron(III) using

the (NH4)2Fe(SO4)2.6H2O provided, before performing the filtration below.

--------------------------------------------------------------------

Collect your product by vacuum filtration using a Buchner filtration flask and a small Buchner or

Hirsch funnel (See Figure 1, overleaf). Pour your suspension quickly onto the filter. Add cold

deionised water (10 mL) to the residue in the beaker, swirl the solid around and pour quickly onto

the filter. Repeat this until you have washed all of the solid out of the beaker.Finally wash the solid

on the filter quickly with 10 mL of cold deionised water from a beaker.

Transfer the product onto a 15 cm filter paper, spread the crystals out, cover with a second filter

paper and press gently to dry the crystals. Remove the top paper and leave the product to dry in

air.

Transfer your dry product to a clean, labelled and tared watch glass and record the weight of your

product on a top-pan balance.

3. Calculate the percentage yield of your ammonium iron(II) sulfate hexahydrate.

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Figure 1: Vacuum Filtration

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Part 2: Test-Tube Reactions of Iron(II) and Iron(III)

A metal complex can be formed when a molecule containing an atom with a lone-pair of electrons

is able to form a donor bond (also called a coordinate bond) to the metal ion. The molecule is called

a ligand, and typical simple ligands are H2O, NH3 and Cl–.

Larger molecules may form bonds from two or more atoms at the same time. Such a molecule is

called a chelate, and is bidentate, tridentate or polydentate if it forms 2, 3 or many bonds

respectively. 1,10-Phenanthroline is an example of a bidentate ligand:

A metal ion in aqueous solution can be represented by Mn+(aq), where (aq) denotes the solvation of

the metal ion by the ligand molecules of water. This is also in fact a complex with typically six water

molecules around the metal ion. The formation of a different complex ion involves a ligand

substitution reaction. A complex ion may also react to form an insoluble compound.

In the following scheme the metal ion M2+(aq) in the presence of OH– ions and the ligands L and L′

will form the insoluble hydroxide M(OH)2 or either of the complex ions (MLn)2+ or (ML'n)2+ depending

on the relative values of the solubility product (Ksp) and the two stability constants K and K′

respectively:

When K is very large, complex formation is virtually complete and is insensitive to ligand

concentration. For low K, complex formation is sensitive to changes in ligand concentration, and

hence is reversible (in equilibrium).

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Transition elements have d-orbitals of low energy which result in more diversity in forming chemical

bonds and a greater variety of stable oxidation states than main group metals. The formation of a

complex is often accompanied by a change in colour which our eyes can see clearly.

Procedure

In all of the test-tube experiments in this section, the graduation marks on the plastic dropping

pipettes provided (top mark = 3 mL) are sufficiently accurate when measuring out the reagents for

the test reactions.

You are provided with an approximately 0.05 M solution of iron(III) chloride.

Make up a 0.05 M solution of ammonium iron(II) sulfate hexahydrate by dissolving 0.4 g of the

sample provided in 20 mL of deionised water.

Reaction of Fe(II) and Fe(III) with ammonia solution.

Pipette 2 mL of the 0.05 M FeCl3 solution into a test-tube and add 20 drops of 2 M ammonia

solution. Record your observations below.

Pipette 2 mL of the 0.05 M ammonium iron(II) sulfate solution into a test-tube and add 20 drops of

2 M ammonia solution. Record your observations below.

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4. Write simple ionic equations for the reactions of ammonia solution with iron(II) and iron(III)

ions in aqueous solution.

Remember that when dissolved in water, ammonia reacts to produce ammonium and

hydroxide ions.

Reaction of Fe(II) and Fe(III) with potassium hexacyanoferrate(III) solution and potassium

hexacyanoferrate(II) solution.

These tests rely on the formation of the insoluble Prussian blue complex: Fe7(CN)18.xH2O (x = 14-

16).

Pipette 2 mL of the 0.05 M ammonium iron(II) sulfate solution into a test-tube and add 1 mL of

potassium hexacyanoferrate(III) solution.Record your observations below.

Pipette 2 mL of the 0.05 M ammonium iron(II) sulfate solution into a test-tube and add 1 mL of

potassium hexacyanoferrate(II) solution. Record your observations below.

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Pipette 2 mL of the 0.05 M iron(III) chloride solution into a test-tube and add 1 mL potassium

hexacyanoferrate(III) solution. Record your observations below.

Pipette 2 mL of the 0.05 M iron(III) chloride solution into a test-tube and add 1 mL potassium

hexacyanoferrate(II) solution. Record your observations below.

5. A diseased biological tissue sample needs to be tested for iron. Using your observations

from the tests above, explain how you could use the formation of Prussian blue complex to

determine whether the sample contained mainly Fe(II) or Fe(III).

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Reaction of iron(II) with potassium manganate(VII).

This observation will form the basis of the quantitative analysis carried out in Part 3.

Pipette 2 mL of the 0.05 M ammonium iron(II) sulfate solution into a test-tube and add 10 drops of

1 M sulfuric acid. Add 10 drops of 0.02 M potassium manganate(VII) solution. Record your

observations below.

6. Explain what happens when a solution of iron(II) reacts with an acidified solution of

potassium manganate(VII).

The following half equation may help:

Reaction of iron(II) and 1,10-phenanthroline.

This observation will form the basis of the quantitative analysis carried out in Part 4.

In a 10 mL measuring cylinder, dilute 1mL of your 0.05 M solution of ammonium iron(II) sulfate

hexahydrate to 10 mL with deionised water. Add your diluted solution to a depth of 1cm in a test-

tube. Then add up to three times that volume of the approximately 0.006 M solution of 1,10-

phenanthroline. Record your observations on the next page.

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7. Give an equation to describe what you observe when 1,10-phenanthroline (o-phen) is

added to the iron(II) solution.

Note: Iron(II) reacts to give the complex [Fe(o-phen)3]2+.

Return to Part 1 to filter off your product, which you will use this afternoon.

In Parts 3 and 4 you will determine the percentage by mass of iron in your sample of ammonium

iron(II) sulfate: (NH4)2Fe(SO4)2.6H2O. You can then compare the results obtained from the two

methods.

END OF MORNING SESSION

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Part 3: Iron(II) Analysis by Titration with Manganate(VII)

The concentration of iron(II) is determined here by volumetric analysis involving a manganate(VII)

redox titration. The manganate(VII) oxidises iron(II) to iron(III). Note that no indicator is required for

this titration. Due to the intense colour of the manganate(VII) ion, the titration is self indicating as

the first drop of excess reagent gives rise to a permanent pink colouration of the titration solution.

The relevant half equations for the redox titration are:

By eliminating electrons we obtain the equivalence:

(MnO4)– 5Fe2+

Procedure

In Parts 1 and 2 you used a top-pan balance that can only weigh to the nearest 0.01 g. For a

chemical analysis you need to use an analytical balance with an accuracy of 0.0001 g. Chemicals

should never be transferred directly to a vessel on an analytical balance. Spillage of chemicals will

damage the sensitive (and expensive!) instrument.

Using the top-pan balance, weigh out into a tared weighing bottle around 1.0 ± 0.1 g of your

ammonium iron(II) sulfate hexahydrate. Carefully clean up any spillage.

Take the weighing bottle to the analytical balance. Tare the balance and obtain the accurate

combined mass of the weighing bottle and your sample for analysis. Record your results in the

table on the next page.

Transfer your sample from the weighing bottle to the into a freshly rinsed 250 mL conical flask. (It

does not matter if a few fragments of the sample remain in the bottle).

Weigh accurately (using the analytical balance) the now empty weighing bottle and determine the

accurate mass of your transferred analysis sample. Record your results in the table on the next

page.

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Weighings Titration 1 Titration 2 Titration 3

Mass of weighing bottle + sample

Mass of empty weighing bottle

By difference: mass transferred

Add 50 mL of 1 M sulfuric acid to your ammonium iron(II) sulfate hexahydrate in the conical flask

and swirl to dissolve. Position the flask on a white tile underneath the burette of standard

potassium manganate(VII) solution. Record the accurate molarity of the standard manganate(VII)

solution.

Exact molarity of potassium manganate(VII) solution: ______ M

With gentle swirling, titrate the solution of iron(II) until a pale pink ‘end-point’ is observed. Record

your burette readings below. Repeat the measurement at least twice on separate weighed

samples.

Burette readings Titration 1 Titration 2 Titration 3

Initial burette reading

(mL)

Final burette reading

(mL)

Volume of titre (mL)

Mass/Titre

In this method of titration, the volume of titre should differ, but the calculated value of mass/titre

should be consistent. You may consider the results consistent if they agree with each other to

within 2%.

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Calculation of the mass of iron in your sample.

8. Calculate the number of moles of manganate(VII) in your titre

(molarity x titre/1000).

9. From the moles of manganate(VII) calculated above, calculate the number of moles of iron

in your mass of ammonium iron(II) sulfate hexahydrate and convert the number of moles of

iron into the corresponding mass of iron (mass = moles x atomic weight of iron). (Atomic

weight of iron = 55.847).

Remember (MnO4)– 5Fe2+

10. Calculate the experimental percentage by mass of iron in your sample of ammonium iron(II)

sulfate hexahydrate.

The atomic weights are: H (1.0079), N(14.0067), O (15.9994), S (32.064), Fe (55.847)

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11. Calculate the theoretical value for the percentage by mass of iron in ammonium iron(II)

sulfate hexahydrate, and compare it with the experimental value.

Part 4: Iron Analysis by Spectrophotometry

In absorption spectrophotometry we measure the absorbance at certain wavelengths of light. λmax

is the wavelength of the maximum absorbance.

At a given wavelength the absorbance measured, A, is often directly proportional to concentration:

this is known as the "Beer Lambert Law":

A = ε c l

where ε is the molar extinction coefficient (litre mol-1 cm-1)

c is the concentration of the absorbing species (mol litre-1)

l is the path length of the cell (cm)

A is the absorbance (dimensionless).

The molar extinction coefficient, ε, provides a measure of the sensitivity of absorbance.

350 400 450 500 550 600 650 700 750

Wavelength in nm

A

0.8

0.6

0.4

0.2

0

blue green red

350 400 450 500 550 600 650 700 750

Wavelength in nm

A

0.8

0.6

0.4

0.2

0

blue green red

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It is possible to produce a calibration graph using solutions of known concentration, and this can

then be used to find the concentration in any number of other ‘unknown’ solutions.

Procedure

Preparation of a ‘stock’ solution of iron(II) ammonium sulfate of ‘unknown’ iron concentration.

Measure 1 M dilute sulfuric acid (50 mL) into a clean freshly rinsed, one litre volumetric flask.

Weigh accurately (use an analytical balance) about 0.140 g of your iron(II) ammonium sulfate

using a clean weighing bottle. Carefully transfer all of the solid through a funnel into the one litre

volumetric flask. Record your results in the table below.

Weigh accurately (use an analytical balance) the empty weighing bottle in order to calculate, by

difference, the accurate mass transferred to the flask. Wash deionised water through the funnel

into the flask. Gently shake the flask until all the solid has dissolved, and finally make up to the

mark with more deionised water. Record your results in the table below.

Mass of weighing bottle + sample (g)

Mass of weighing bottle (g)

Mass of sample transferred (g)

Determination of the wavelength of maximum absorbance, λmax, for Fe(o-phen)3]2+.

Preparation of solutions of reagents.

In a 50 mL beaker, dissolve hydroxylamine hydrochloride, (1.0 g) in water (10 mL).

In a 250 mL beaker, dissolve sodium ethanoate trihydrate (17 g), in water (100 mL).

Concentration

A

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You require four 100 mL volumetric flasks. Into three of the flasks, pipette respectively 25, 10, 5 mL

of your ‘stock’ solution of iron(II) ammonium sulfate. The last flask will be used as a blank.

To each flask (including the ‘blank’ flask) in turn add:

1 mL of your previously prepared hydroxylamine hydrochloride solution (using a 10 mL measuring

cylinder)

10 mL of your sodium ethanoate solution (use a 25 mL measuring cylinder)

10 mL of the 1,10-phenanthroline solution provided (use a 25 mL measuring cylinder)

Allow the mixtures to stand for 5 minutes before diluting carefully to the mark in each flask with

deionised water and shake to mix the solutions well.

Obtain a visible spectrum of the solution of the [Fe(o-phen)3]2+ complex with highest

concentration.

Take all of the solutions to the UV/vis spectrophotometer. You will be shown how to use the

scanning UV/vis spectrophotometer. Using the solution with the highest concentration of the

[Fe(o-phen)3]2+ complex, rinse the 1 cm cell provided with a little of the solution, and then fill the

cell to within 1cm of the top. Insert the cell into the sample holder of the instrument as instructed.

Insert a second cell containing your blank solution into the reference beam. The

spectrophotometer will produce an absorption scan over the range 350 - 850 nm (i.e. the visible

region). From this spectrum of the [Fe(o-phen)3]2+ complex solution, determine the λmax value.

This should be close to 510 nm.

Measured λmax value of solution

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Absorbance measurements.

Take all four solutions to a second instrument which is a non-scanning (fixed wavelength)

spectrophotometer. You will be shown how to use the instrument. Select the wavelength to be

that of λmax for the complex as determined above. Using your blank solution as reference, measure

the absorbance of each of the three red solutions of the

[Fe(o-phen)3]2+ complex. Take an average of at least two readings for each solution.

Volume of ‘stock’ solution 5 mL 10 mL 25 mL

1st determination

2nd determination

Average value

Calculation of the iron concentration.

From A = ε c l (and with a path length of l = 1 cm), a graph of A versus c will have a gradient of ε =

A/c. The molar extinction coefficient, ε, for the [Fe(o-phen)3]2+ complex is known to be 1.08 x

104 litre mol–1 cm–1.

12. Draw a graph of absorbance (y axis), against ‘volume of unknown stock’ solution (x axis)

(i.e. 5, 10 or 25 mL), for your three red solutions of the [Fe(o-phen)3]2+ complex. Because

you used a blank to set the instrument reference, you can use the origin as a point on your

graph, and draw a straight line from the origin to best fit your data. From your graph find

the absorbance, A, at a concentration of 25 mL.

Note: This may not correspond exactly to the experimental value for this solution

A at 25 mL = ______

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13. From A / (25x/100) = 1.08 x 104 litre mol-1 cm-1, you can obtain a value for x, the

concentration x in mol litre-1 of iron in your ‘stock’ solution of iron(II) ammonium sulfate.

14. You now know both the accurate mass of iron(II) ammonium sulfate which is dissolved in

the litre of ‘stock’ solution, and also the concentration of iron in mol litre–1. Calculate the

mass of iron dissolved in the litre, and hence the percentage by mass of iron in your iron(II)

ammonium sulfate sample (atomic weight 55.847).

15. Compare the percentage mass of iron in your sample obtained by spectrophotometry and

titration. Which of the methods do you think is more accurate and why?

Percentage mass of iron by spectrophotometry =

Percentage mass of iron by manganate(VII) titration (Part 3) =

END OF AFTERNOON SESSION