good use of the “researching chemistry reference … unit covers the key areas of gravimetric...

The general aim of this Unit is to develop skills of scientific inquiry, investigation, analytical

thinking, independent working, and knowledge and understanding of researching chemistry.

The Unit covers the key areas of gravimetric analysis, volumetric analysis, practical skills and

techniques and stoichiometric calculations.

You have already come across these areas of analytical chemistry during S4 and S5. At

Advanced Higher level you will develop these in much greater depth.

It is essential when working through the experiments and calculations in this booklet to make

good use of the “Researching Chemistry reference booklet”.

When you have completed the experiments and

questions in this booklet you should have detailed

knowledge of the techniques shown in the table below.

There is a Researching Chemistry Unit assessment on these skills and the calculations

involved. Also there will be questions in the examination paper relating to these techniques,

including the calculations. Additionally the skills and techniques will be invaluable to you

when you undertake your Chemical Investigation.

Technique Description

a Weighing by difference and gravimetric analysis

b Preparing a standard solution

c Using a control

d Carrying out a complexometric titration

e Carrying out a back titration

f Using a colorimeter and carrying out dilution to prepare a calibration graph

g Distilling

h Refluxing

i Using vacuum filtration

j Recrystallising

k Determining % yield experimentally

l Using thin-layer chromatography

m Using melting point apparatus and mixed melting point determination

n Using a separating funnel and solvent extraction

Stoichiometry is the study of quantitative relationships involved in chemical reactions. The

ability to balance and interpret equations enabling calculations to be carried out involving any

of the skills/techniques you will learn is an important part of chemistry at this level.

Prior to getting started on Advanced Higher level problems we will revisit some questions

encountered at higher level.

1. Calculate the number of moles of the named chemical in each of the following.

a. 1.32g of ammonium sulfate, (NH4)2SO4

b. 200cm3 of 3 mol l-1 sodium hydroxide, NaOH

2. 250 cm3 of a solution of magnesium chloride, MgCl2, is known to contain 0.5 moles of

chloride ions. What mass of magnesium does the solution contain?

3. In an experiment to prepare aspirin, 10 g of 2 – hydroxybenzoic acid, HOC6H4COOH,

was treated with excess ethanoyl chloride. The equation for the reaction is

CH3COCl + HOC6H4COOH CH3COOC6H4COOH + HCl

Calculate the mass of aspirin formed in this reaction.

4. Calculate the percentage by mass of nitrogen in urea CO(NH2)2

5. Titanium is manufactured by heating titanium(IV) chloride with sodium.

TiCl4 + 4Na Ti + 4NaCl

If this reaction is only 75% efficient, calculate the mass of sodium is required to

produce 100 kg of titanium?

6. 1.20g of magnesium was used to make hydrated magnesium sulfate crystals,

MgSO4.7H2O

The mass of crystals produced was 9.84 g. Calculate the percentage yield.

7. Zinc reacts with dilute hydrochloric acid according to the equation.

Zn + 2HCl ZnCl2 + H2

1.625 g of zinc is mixed with 250 cm3 of 2mol l-1 hydrochloric acid. Calculate which

reactant is in excess

8. 25.0 cm3 of hydrogen peroxide, H2O2, was titrated with 0.040 mol l-1 acidified

potassium permanganate. The end - point was reached when 16.7 cm3 of the

permanganate solution had been added.

5H2O2 + 2MnO4- + 6H+ 2Mn2+ + 5O2 + 8H2O

Calculate the concentration of the hydrogen peroxide in moles per litre.

9. Seaweeds are a rich source of iodine in the form of iodide ions. The mass of iodine

in seaweed can be found using the procedure outlined below.

50.00 g of seaweed is dried in an oven and ground into a fine powder. Hydrogen

peroxide solution is then added to oxidise the iodide ions to iodine molecules.

Using starch solution as an indicator, the iodine solution is then titrated with sodium

thiosulfate solution to find the mass of iodine in the sample. The balanced equation

for the reaction is shown.

2Na2S2O3(aq) + I2(aq) 2NaI(aq) + Na2S4O6(aq)

In this analysis of seaweed, 14.9 cm3 of 0.00500 mol l–1 sodium thiosulfate solution was

required to reach the end-point.

Calculate the mass of iodine present in one gram of the seaweed sample.

Preparation of a primary standard solution of 0.1 mol l–1 oxalic acid.

A standard solution is one of accurately known concentration and can be prepared

directly from a primary standard which, in this case, is hydrated oxalic acid,

(COOH)2.2H2O (GFM = 126.1 g).

To prepare 250 cm3 of 0.1 mol l–1 oxalic acid solution, the mass of hydrated oxalic

acid required must be calculated.

Requirements

balance (accurate to 0.01 g), oxalic acid AnalaR, (COOH)2.2H2O, weighing bottle,

deionised water, 250 cm3 beaker, 250 cm3 standard flask, wash bottle, dropper, glass

stirring rod, filter funnel

Hazcon

Wear eye protection and if any chemical splashes on the skin, wash it off

immediately. Oxalic acid is harmful if ingested and irritates the eyes and skin. Wear

gloves.

Calculation

Calculate the mass of oxalic acid needed to produce 250 cm3 of 0.1 mol l-1 oxalic acid

solution.

Procedure

1. Transfer approximately 3.2 g of oxalic acid crystals to the weighing bottle and

weigh accurately.

2. Pour the oxalic acid crystals into a clean beaker containing about 50 cm3 of

deionised water and reweigh accurately the weighing bottle and any remaining

crystals.

3. Stir the solution until all the oxalic acid dissolves and then transfer it to a

250 cm3 standard flask.

4. Rinse the beaker several times with deionised water and add all the rinsings to

the flask.

5. Make up the solution to the graduation mark with deionised water.

6. Stopper the flask and invert it several times to ensure the contents are

completely mixed. Retain the solution for the next experiment.

Initial Mass of weighing bottle + oxalic acid = __________

Final mass of weighing bottle + oxalic acid = __________

Mass of oxalic acid = __________

a. Use the exact mass of oxalic acid transferred to the beaker in step 2 to

calculate the concentration of the oxalic acid solution you prepared.

b. What is a standard solution?

c. (i) One requirement of a primary standard is that it should have a high gram

formula mass.

What is the reason for this?

(ii) State two other requirements a primary standard must have.

(iii) Why is sodium hydroxide NOT a primary standard?

Standardisation of approximately 0.1 mol l–1 sodium hydroxide.

Sodium hydroxide is not a primary standard and so a standard solution of it cannot

be prepared directly from the solid. However, a solution of approximate

concentration can be prepared and its exact concentration determined by titrating it

against an acid of accurately known concentration using a suitable indicator.

In this experiment, a sodium hydroxide solution is standardised against the 0.1 mol l–1

oxalic acid solution prepared in Experiment 1.

The stoichiometric equation for the titration reaction is:

(COOH)2 + 2NaOH 2H2O + (COONa)2

Requirements

10 cm3 pipette standardised oxalic acid solution (approx. 0.1 mol l –1)

50 cm3 burette sodium hydroxide solution (approx. 0.1 mol l–1)

100 cm3 beakers, phenolphthalein indicator, 100 cm3 conical flasks, deionised water,

wash bottle, pipette filler, white tile, filter funnel.

Hazcon


immediately.

0.1 mol l–1 oxalic acid irritates the eyes and skin.

0.1 mol l–1 sodium hydroxide is corrosive to the eyes and skin.

Phenolphthalein indicator solution is highly flammable and irritating to the eyes

because of its ethanol content.

Procedure

1. Rinse the 10 cm3 pipette with a little of the oxalic acid solution and pipette

10 cm3 of it into a conical flask.

2. Add two or three drops of phenolphthalein indicator to the oxalic acid solution in

the flask.

3. Rinse the 50 cm3 burette, including the tip, with the sodium hydroxide solution

and fill it with the same solution.

4. Titrate the oxalic acid solution with the sodium hydroxide solution from the

burette until the end-point is reached. This is indicated by the appearance of a

pink colour.

5. Repeat the titrations until two concordant results are obtained, (+/- 0.1 cm3).

6. Retain the standardised sodium hydroxide solution for the next experiment.

a. Use the stoichiometric equation and your mean titre volume to calculate the

concentration of the sodium hydroxide solution.

b. In a titration we try to find the “equivalence point”. In reality we can only

determine the “end-point”.

Explain the difference between the equivalence point and the end-point in a

titration.

Titration Trial 1 2 3 4

Burette

readings/cm3

Initial

Final

Titre volume/cm3

Mean titre volume/cm3

c. In any titration there are uncertainties associated with the apparatus used.

For this titration, assuming CLASS B equipment, calculate

(i) the percentage uncertainty in the volume of the oxalic acid you pipetted

into the conical flask.

(ii) the percentage uncertainty in your mean titre volume. {Do not forget

about the uncertainty in the end-point}

d. It is good practice to carry out a titration until concordant titre values are

obtained. Explain why this procedure does not make titration results more

accurate.

e. Andy and Brenda were carrying out a titration. Brenda pipetted 25 cm3 of the

solution being analysed into a conical flask. Andy noticed that the conical flask

was wet as he had not properly dried it after cleaning it with some deionised

water. He told Brenda she must discard the solution, clean the flask and then

make sure it was completely dry before proceeding. Brenda said it did not matter

that the flask was wet and carried on with the titration.

Explain whether you agree with Brenda or Andy.

The next few pages will look at how to tackle calculations based on chemical analysis.

This type of question is very popular in the exam.

Three types of titration

Acid/base

In these reactions the hydrogen ions of the acid react with the hydroxide ions of the base.

Stoichiometry ---- Every H+ needs one OH- from the base.

H+ + OH- H2O

Therefore

one mole of HCl will react with one mole of NaOH

one mole of H2SO4 will react with two moles of NaOH

In an organic acid it is only the functional group hydrogen atoms that are acidic. (COOH)

Ethanoic acid, CH3COOH, only has ONE acidic hydrogen. Oxalic acid, HOOCCOOH, has

TWO.

Carbonates will neutralise TWO H+

CO32- + 2H+ H2CO3

Primary Standards ---- Include Sodium carbonate, potassium hydrogen phthalate and

oxalic acid.

Note that most HYDROXIDES are NOT primary standards.

Indicators ---- Include Phenolphthalein and methyl orange.

REDOX

Oxidation and reduction – remember OXIDISING agents are LOW LEFT and REDUCING

agents are HIGH RIGHT in the electrochemical series in the data booklet.

Strong oxidising agents are permanganate, MnO4-, dichromate, Cr2O7

2-

Reducing agent frequently used is the thiosulphate ion, S2O32-

Stoichiometry ---- To find the stoichiometry – simply cross multiply the electrons

in the ion-electron equations.

E.g. Fe2+ Fe3+ + e

MnO4- + 8H+ + 5e Mn2+ + 4H2O

The stoichiometry is 5Fe2+ : MnO4-

Primary Standards ---- Potassium iodate, potassium dichromate.

Indicators ---- Very common REDOX titration is reaction of iodine (I2) with sodium

thiosulphate ,(S2O32-)

The indicator used in this titration is STARCH

Be careful - IODINE turns black with starch but IODIDE does not.

I2 + 2e- 2I-

COMPLEXOMETRIC

The reaction of ethylenediaminetetraacetic aicd, EDTA, with divalent metal ions like Ni2+,

Ca2+, Mg2+ to form an octahedral complex with a coordination number of six.

Stoichiometry ---- EDTA reacts in a 1:1 ratio with these ions

Indicators ---- These vary with the metal – Murexide and Eriochrome black T are examples.

EDTA titrations are pH sensitive. Buffer solutions are often added to control the pH.

{You will find out more about EDTA in the transition metals notes}

1. Potassium dichromate, K2Cr2O7, is a primary standard and is often used in acid

conditions in redox titrations.

This oxidising agent can be used to determine the mass of iron in an iron tablet.

An iron tablet weighing 0.940 g was dissolved in dilute sulphuric acid made up to

250 cm3 with water. 25.0 cm3 of this solution was titrated with 0.00160 mol l-1

K2Cr2O7 requiring 32.5 cm3 of the K2Cr2O7.

The equation for the reaction is

6Fe2+ + Cr2O72- + 14H+ 2Cr3+ + 6Fe3+ + 7H2O

Calculate the percentage mass of Fe2+ in the tablet.

2. 0.236g of impure benzoic (C6H5COOH) acid required 19.25 cm3 of 0.10 mol l-1 sodium

hydroxide for complete neutralisation.

Calculate

a. The number moles of sodium hydroxide used in titration.

b. The mass of benzoic acid which neutralised the sodium hydroxide.

c. The % purity of the benzoic acid.

3. To make a standard calcium ion solution 0.25 g of calcium carbonate, CaCO3, was

dissolved in a little dilute hydrochloric acid and made up to 250 cm3 in a standard

flask.

a. Calculate the concentration of the calcium ion in this solution.

b. 20 cm3 of the standard calcium solution were pipetted into a conical flask. A few

drops of calcium ion indicator and a pH 10 buffer were added. EDTA was added

from a burette and the end point was reached at a volume of 25.70 cm3.

Calculate the concentration of the EDTA.

Ca

calcium EDTA complex

Determination of the ethanoic acid content of white vinegar

Vinegar is a dilute solution of ethanoic acid, CH3COOH. The aim of this experiment is to

determine the concentration of ethanoic acid in a given sample of white vinegar by titration

against the sodium hydroxide solution standardised in Experiment 2.

The stoichiometric equation for the titration reaction is:

CH3COOH + NaOH H2O + CH3COONa

ethanoic acid sodium hydroxide water sodium ethanoate

Requirements

50 cm3 burette, white vinegar, 25 cm3 pipette,

standardised sodium hydroxide solution (approx. 0.1 mol l–1), 100 cm3 beakers,

100 cm3 conical flasks, phenolphthalein indicator,250 cm3 standard flask, deionised water,

wash bottle, pipette filler, dropper, white tile, filter funnel.

Hazcon

Wear eye protection and if any chemical splashes on the skin, wash it off immediately.

Vinegar irritates the eyes and skin. 0.1 mol l–1 sodium hydroxide is corrosive to the eyes and

skin. Phenolphthalein indicator solution is highly flammable and irritating to the eyes

because of its ethanol content.

Procedure

1. Rinse the 25 cm3 pipette with a little of the vinegar.

2. Dilute the sample of vinegar by pipetting 25 cm3 of it into a clean 250 cm3 standard

flask and making it up to the graduation mark with deionised water.

3. Stopper the standard flask and invert it several times to ensure the contents are

thoroughly mixed.

4. Rinse the 25 cm3 pipette with a little of the diluted vinegar(or use a second pipette)

and pipette 25 cm3 of it into a conical flask.

5. Add two or three drops of phenolphthalein indicator to the diluted vinegar in the

conical flask.

6. Rinse the 50 cm3 burette, including the tip, with the sodium hydroxide solution and fill

it with the same solution.

7. Titrate the diluted vinegar solution with the sodium hydroxide solution from the

burette until the end-point is reached. This is indicated by the appearance of a pink

colour. Repeat the titrations until two concordant results are obtained.

a. Use the stoichiometric equation and your mean titre volume to calculate the

concentration of ethanoic acid in the undiluted vinegar.

b. Refer to the vinegar bottle and note the volume. Use this information and your answer

to question (a) above to calculate the mass of ethanoic acid in the vinegar when the

bottle is full.


Burette

readings/cm3

Initial

Final

Titre volume/cm3


Determination of nickel in a nickel(II) salt using EDTA

The molecule ethylenediaminetetraacetic acid (EDTA) has the following structure.

alternatively

It is often written as H4Y where Y represent all the molecule except the four COOH

hydrogen atoms.

In alkaline conditions the molecule forms the Y4- ion. This ion reacts in a one to one

stoichiometric ratio with a range of divalent metal ions (Mg2+, Ni2+ etc).

The EDTA ion “wraps around” the metal ion forming a metal-edta octahedral complex which

is represented by the structures shown.

Since EDTA forms stable complexes with most metal ions, it

is widely used to determine metals in what are known as

complexometric titrations.

The reaction of nickel(II) ions with EDTA can be represented as

Y4– + Ni2+ NiY2–

The end-point of an EDTA complexometric titration can be detected by means of a metal

ion indicator – an organic dye which changes colour when it binds with metal ions. For it to

be suitable in an EDTA titration, the indicator must bind less strongly with metal ions

than does EDTA. Murexide is one such indicator.

Requirements

50 cm3 burette standardised 0.10 mol l–1 EDTA solution

10 cm3 pipette hydrated nickel(II) sulfate (NiSO4.6H2O)

100 cm3 standard flask, 1 mol l–1 ammonium chloride, 250 cm3 conical flasks

murexide indicator, weighing bottle, 0.88 aqueous ammonia, balance (accurate to 0.01 g),

deionised water, 100 cm3 beakers, 25 cm3 measuring cylinder, wash bottle, pipette filler,

white tile, filter funnel, glass stirring rod.

Hazcon


Hydrated nickel(II) sulfate is harmful by ingestion and inhalation. Wear gloves.

EDTA is only toxic if ingested in large quantities.

0.88 aqueous ammonia is toxic if inhaled in high concentrations or if swallowed. The solution

and vapour irritate the eyes. The solution burns the skin. Wear goggles and gloves and

handle it in a fume cupboard. 1 mol l–1 ammonium chloride is harmful and irritates the eyes.

Murexide is harmful by ingestion and if inhaled as a dust.

Procedure

1. Transfer approximately 2.6 g of hydrated nickel(II) sulfate to a weighing bottle and

weigh the bottle and contents.

2. Add about 25 cm3 of deionised water to a 100 cm3 beaker and transfer the bulk of the

nickel salt to the water.

3. Reweigh the bottle with any remaining salt.

4. Stir the mixture until the solid dissolves and transfer the resulting solution to a


5. Rinse the beaker several times with a little deionised water and add the rinsings to the

standard flask.

6. Make up the solution to the graduation mark with deionised water. Stopper the flask

and invert it several times to ensure the contents are thoroughly mixed.

7. Rinse the burette, including the tip, with 0.10 mol l–1 EDTA and fill it with the same

solution.

8. Rinse the 10 cm3 pipette with a little of the nickel salt solution and pipette 10 cm3 of

it into a conical flask. Dilute the solution to about 50 cm3 with deionised water.

9. Add murexide indicator (approximately 0.05 g) to the diluted nickel salt solution

together with approximately 5 cm3 of ammonium chloride solution.

10. Titrate the mixture with the EDTA solution and after the addition of about 8 cm3

make the solution alkaline by adding approximately 5 cm3 of 0.88 aqueous ammonia

(concentrated ammonia solution).

11. Continue the titration to the end-point, which is shown by the first appearance of a

blue-violet colour. Detection of the end-point can be difficult so keep this titrated

solution to help you detect end-points in subsequent titrations.

12. Repeat the titrations until two concordant results are obtained.

Initial Mass of weighing bottle + nickel salt = __________

Final mass of weighing bottle + nickel salt = __________

Mass of nickel salt = __________

The aim of this experiment is to determine by EDTA titration the percentage nickel in

hydrated nickel(II) sulfate, NiSO4.6H2O and to compare it with the theoretical value.

The following theory will remind you how to calculate the theoretical % of an element in a

compound.


Burette

readings/cm3

Initial

Final

Titre volume/cm3


a. Calculate the theoretical percentage by mass of nickel in NiSO4.6H2O

b. Calculate the percentage by mass of nickel in the sample of hydrated nickel(II) sulfate

using the stoichiometric equation and your experimental results.

c. Account for the difference in your answers to (a) and (b).

d. The end-point of this titration was observed due to the metal ion indicator murexide

(symbol Mu) causing a colour change.

The equations below occur at various stages during the titration

Ni2+(aq) + Mu(aq) NiMu2+(aq) (murexide added at the start)

(initial colour -blue)

Ni2+(aq) + EDTA4-(aq) NiEDTA2-(aq) (adding EDTA to nickel solution )

NiMu2+(aq) + EDTA4-(aq) NiEDTA2-(aq) + Mu(aq) (at the end-point )

(final colour- blue/violet)

Use the information in the equations and the fact that to be suitable in an EDTA

titration, the indicator must bind less strongly with metal ions than does EDTA

to explain how the indicator works.

e. Explain why it is critical not to add too much indictor in an EDTA titration .

It is good practice, especially when using an unfamiliar procedure, to carry out a

control experiment. The purpose of a control is to validate the technique.

A control experiment consists of carrying out the analysis with a pure sample of the

analyte of known mass or concentration.

For example in the analysis of an aspirin tablet the control would involve carrying

out the determination of aspirin using a pure sample of the compound. If the mass

of aspirin you determine matches the mass you started with then this establishes

the validity of the procedure and the results.

Other examples of controls are shown in the table.

Vitamin C (ascorbic acid) is an important component of our diet. Although it occurs

naturally in many fruits and vegetables, many people take vitamin C tablets to

supplement their intake. The vitamin C content of a tablet can be determined by

carrying out a redox titration with a standard solution of iodine using starch

solution as indicator:

The equation for the reaction is shown below:

C

C

HO

HO C

O

CH

CH

OH

CH2OH

O + I2

C

C

O

O C

O

CH

CH

OH

CH2OH

O + 2H+ + 2I

vitamin C(ascorbic acid)

Substance analysed Control

Vitamin C in fruit juice Pure vitamin C

Iron in Iron tablets Pure iron compound. E.g. Iron sulfate

Caffeine in tea Pure caffeine

Copper in brass Mixture of pure copper and pure zinc

Requirements

250 cm3 standard flask, 1 g effervescent vitamin C tablet, 100 cm3 conical flasks,

sample of pure ascorbic acid, 25 cm3 pipette, standardised 0.025 mol l–1 iodine solution,

50 cm3 burette, starch solution, weighing bottle,deionised water,

balance (accurate to 0.01 g), pipette filler, filter funnel, 100 cm3 beakers, dropper,

white tile, wash bottle.

Hazcon


0.025 mol l–1 iodine solution irritates the eyes and causes severe internal irritation if

swallowed. Wear gloves and treat any spills on the skin with sodium thiosulfate solution.

Control experiment using pure ascorbic acid

1. Add about 1.0 g of pure ascorbic acid to the weighing bottle and weigh the

bottle and contents.

2. Transfer the pure ascorbic acid to a beaker and reweigh the weighing bottle

3. Add some deionised water (approximately 50 cm3) to the beaker and stir the

mixture until the ascorbic acid dissolves.

4. Transfer the solution to a 250 cm3 standard flask.

5. Rinse the beaker with a little deionised water and add the rinsings to the

standard flask. Repeat this procedure several times and add the rinsings to the

flask. Make up the solution to the graduation mark with deionised water.


completely mixed.

7. Rinse the burette, including the tip, with 0.025 mol l–1 iodine solution and fill it

with the same solution.

8. Rinse the 25 cm3 pipette with the ascorbic acid solution and pipette

25 cm3 of it into a 100 cm3 conical flask.

9. Add a few drops of starch indicator to the solution and titrate to the end-

point, which is indicated by the colour changing to blue.


Determination of vitamin C (ascorbic acid) in a commercial tablet

1. Add a 1 g effervescent vitamin C tablet to a beaker.

2. Repeat steps 2 to 10 of the above procedure.

Control

Initial Mass of weighing bottle + vitamin C = __________

Final mass of weighing bottle + vitamin C = __________

Mass of vitamin C = __________

Tablet

Mass of vitamin C tablet (tared balance) = __________


Burette

readings/cm3

Initial

Final

Titre volume/cm3



Burette

readings/cm3

Initial

Final

Titre volume/cm3


a. Use your concordant result, the concentration of iodine and the gram formula mass of

vitamin C to calculate the mass of vitamin in the initial control sample.

Compare this result with the mass of vitamin C weighed out and account for any

difference.

b. Calculate the mass of vitamin C in the tablet and compare the result with the

manufacturer’s specification.

c. (i) Does the control result conform this as a valid technique? Explain

c. (ii) The percentage efficiency of the technique can be calculated as follows

% Efficiency = {experimental mass/actual mass} x 100

Use the results of your control experiment to determine the % efficiency of

the technique and use this figure to decide if the result for the tablet is

acceptable.

d. Starch is used as an indicator in this titration even although iodine could be considered

self-indicating.

(i) What does self-indicating mean?

(ii) Why does an iodine titration require the addition of starch?

(iii) Why must the starch solution be freshly prepared?

Sometimes it is not possible to use standard titration methods. For example the reaction

between determined substance and titrant can be too slow, or there can be a problem with

end point determination or the substance may not be soluble.

In such situations we can often use a technique called back titration. In back titration we

use two reagents - one, that reacts with the original sample (lets call it A), and second (lets

call it B), that reacts with the first reagent. How do we proceed? We add precisely

measured excess amount of reagent A to the sample and once the reaction ends we titrate

excess reagent A left with reagent B. Knowing the initial amount of reagent A and amount

that was left after the reaction (from titration) we can easily calculate how much reagent A

reacted with the sample and from this we can obtain data about the sample.

A sample of magnesium carbonate was analysed for purity using a back titration.

Brenda added 0.25g of impure magnesium carbonate to 40 cm3 of 0.16mol l-1 HCl.

8.1 cm3 of 0.11 mol l-1 NaOH was required to neutralise the excess HCl.

The equations for the reactions are

MgCO3 + 2HCl MgCl2 + CO2 + H2O Eqn 1

HCl + NaOH NaCl + H2O Eqn 2

Calculate the percentage purity of the magnesium carbonate.

Calculate initial moles of HCl CxV = 0.16 x 40/1000 = 0.0064 mol

Calculate moles of NaOH used in titration CxV = 0.11 x 8.1/1000 = 0.000891 mol

Use Eqn 2 to find the excess moles of HCl {HCl:NaOH = 1:1} = 0.000891 mol

Subtract excess moles of HCl from the initial moles of HCl to give moles of HCl which

reacted with the magnesium carbonate 0.0064 - 0.000891 = 0.005509 mol

Use Eqn 1 to find moles of MgCO3 reacted {2HCl: MgCO3 = 2:1}

0.005509/2 = 0.002755 mol

Use n x gfm(MgCO3) for mass of MgCO3 reacted 0.002755 x 84.3 = 0.232 g

Divide the mass of pure MgCO3 with that of the original mass and multiply by 100 to find

the percentage purity of MgCO3 0.232/0.25 x 100 = 92.8%

Determination of aspirin

Aspirin has the following structural formula:

Since it is insoluble in water, aspirin has to be determined by a back titration technique.

This involves treating a sample of accurately known mass with a definite amount of sodium

hydroxide, ie the volume and concentration of the alkali must be accurately known. The

alkali first catalyses the hydrolysis of the aspirin to ethanoic and salicylic acids and then

neutralises these acids. The overall equation for the reaction is:

An excess of alkali has to be used and the amount remaining after reaction is determined by

titrating it against a standard solution of sulfuric acid. This titration is the back titration.

Requirements

250 cm3 standard flasks, aspirin tablets, conical flasks, (100 cm3 and 250 cm3)

25 cm3 pipette, standardised 0.050 mol l–1 sulfuric acid, 50 cm3 burette,

standardised 1.0 mol l–1 sodium hydroxide, weighing bottle, phenolphthalein,

balance (accurate to 0.1 g), deionised water, hot plate (or Bunsen burner and tripod),

50 cm3 measuring cylinder, 100 cm3 beakers, pipette filler, filter funnel, white tile,

wash bottle, dropper.

Hazcon


0.50 mol l–1 sulfuric acid irritates the eyes and skin.

1.0 mol l–1 sodium hydroxide is corrosive to the eyes and skin. Gloves and goggles should be

worn.

Phenolphthalein indicator solution is highly flammable and irritating to the eyes because of

its ethanol content.

Aspirin irritates the eyes and skin.

Procedure

1. Place the conical flask on the balance. Tare the balance. Add aspirin to

the flask (around 1.5 g)

2. Rinse the 25 cm3 pipette with 1.0 mol l–1 sodium hydroxide and pipette

25 cm3 of this solution into the flask containing the aspirin tablets.

3. To the mixture in the flask, add approximately 25 cm3 of deionised water.

4. Place the flask on the hot plate and simmer the mixture very gently for

about 30 minutes.

5. Allow the reaction mixture to cool before transferring it to the


6. Rinse the conical flask with a little deionised water and add the rinsings

to the standard flask. Repeat this procedure several times and add the

rinsings to the flask. Make up the solution to the graduation mark with

deionised water.


completely mixed.

8. Rinse the burette, including the tip, with 0.050 mol l–1 sulfuric acid and

fill it with the same solution.

9. Rinse the 25 cm3 pipette with the ‘standard flask’ solution and pipette 25

cm3 of it into a 100 cm3 conical flask.

10. Add a few drops of phenolphthalein indicator to the solution and titrate

to the end-point.


Number of aspirin tablets used __________________

a. Use your concordant result and the gram formula mass of aspirin to calculate the mass

of aspirin per tablet. Compare this result with the mass of aspirin quoted by the

manufacturer and from this calculate the % efficiency of the technique.


Burette

readings/cm3

Initial

Final

Titre volume/cm3


1. Titration with solutions of potassium bromate (KBrO3) can be used to determine the

concentration of arsenic (III) ions.

The balanced equation is:

3H3AsO3 + BrO3- Br- + 3H3AsO4

What is the concentration of As(III) in a solution if 22.35 cm3 of 0.100 mol l -1 KBrO3

is needed to titrate 50.00 cm3 of the As(III) solution?

2. Alcohol(ethanol) levels in blood can be determined by a redox titration with potassium

dichromate according to the balanced equation:

C2H5OH(aq) + 2Cr2O72-(aq) + 16H+(aq) 2CO2(g) + 4Cr3+(aq) + 11H2O(l)

a. What is the blood alcohol level in mol l-1 if 8.76 ml of 0.050 mol l-1 K2Cr2O7 is required

for titration of a 10.026 cm3 sample of blood?

b. Suggest why an indicator is not required for this titration.

3. To determine the concentration of chloride ions in seawater it is titrated with

silver(I) nitrate solution.

25 cm3 of raw seawater was diluted to 250 cm3 in a volumetric flask.

A 25 cm3 sample of the diluted seawater was pipetted into a conical flask and a few

drops of potassium chromate(VI) indicator solution was added.

On titration with 0.100 mol l-1 silver nitrate solution, 13.8 cm3 was required to

react with the chloride ions in the diluted sample.

The equation for this reaction is

Ag+(aq) + Cl-(aq) Ag+Cl- (s)

a. Silver(I) nitrate is a primary standard. What is meant by “primary standard”?

b. What type of reaction is this?

c. Suggest the name of a chemical which could be used in a control experiment.

d. Calculate the concentration of chloride ions in the undiluted seawater.

4. Brass is an alloy consisting mainly of copper and zinc. To determine the percentage of

copper in a sample of brass, 2.63 g of the brass was dissolved in concentrated nitric

acid and the solution diluted to 250 cm3 in a standard flask. Excess potassium iodide

solution was added to 25.0 cm3 of this solution, iodine being produced according to the

equation:

2Cu2+(aq) + 4I–(aq) 2CuI(s) + I2(aq)

The iodine formed was titrated with 0.10 mol l–1 sodium thiosulfate solution,

Na2S2O3(aq), the volume required for complete reaction being 24.8 cm3.

I2(aq) + 2S2O32–(aq) 2I–(aq) + S4O6

2–(aq)

colourless colourless

a. How could the end-point for the titration be made more obvious?

b. Explain why the potassium iodide solution could be measured out in a measuring

cylinder instead of a pipette.

b. How many moles of sodium thiosulfate were required in the titration?

c. Calculate the percentage by mass of copper in the sample of brass.

5. In an experiment to determine the percentage by mass of ammonium sulfate,

(NH4)2SO4, in a fertiliser, 3.80 g of the fertiliser was dissolved in water and made up

to 250 cm3 in a standard flask..

To 25.0 cm3 portions of this solution, an excess of methanal was added.

4NH4+(aq) + 6HCHO(aq) C6H12N4(aq) + 4H+(aq) + 6H2O(l)

The H+(aq) ions produced were titrated with 0.100 mol l-1 sodium hydroxide solution.

The average volume required to neutralise these H+(aq) ions was 28.0 cm3.

a. Why was the methanal added in excess to the fertiliser solution?

b. Calculate the number of moles of hydrogen ions produced by the methanal in the

25 cm3 of fertiliser solution.

c. Calculate then number of moles of ammonium ions in the 250 cm3 standard flask.

d. Calculate the number of moles of ammonium sulfate in the 250 cm3 standard flask.

e. Calculate the percentage of ammonium sulfate in the fertiliser.

6. Some bleaches use hypochlorite ions (OCl-) as the bleaching agent. The concentration

of the hypochlorite can be found by adding a diluted sample of the bleach to an excess

of ethanoic acid and potassium iodide. This releases iodine, the concentration of which

can be determined using a standard thiosulfate solution with starch indicator.

OCl-(aq) + 2H+(aq) + 2I-(aq) I2(aq) + Cl-(aq) + H2O(l)

2S2O32-(aq) + I2(aq) 2I-(aq) + S4O6

2-(aq)

25.0 cm3 of bleach was diluted to 250 cm3 in a standard flask. 10.0 cm3 samples of

diluted bleach were pipetted into conical flasks containing an excess of ethanoic acid

and potassium iodide. Each sample was titrated against sodium thiosulfate solution

(concentration 0.075 mol l-1). The results are shown in the table below.

a. Why was the titration carried out three times?

b. Why was ethanoic acid added to the bleach/iodide mixture?

c. Calculate the concentration of hypochlorite ions in the undiluted bleach.

Burette reading 1 2 3

Initial/cm3 0.00 14.00 28.00

Final/cm3 13.70 27.55 41.45

Titration

Gravimetric analysis means the analysis of chemicals by weighing. Calculations based on

gravimetric analysis are generally easier than volumetric analysis and almost always involve

calculating a number of moles from an experimentally determined mass.

There are two major types of gravimetric analysis one involves heating a weighed mass of

reactant and weighing one of the solid products.

Copper(II) sulfate can exist in a white anhydrous form, CuSO4,

and the more common blue hydrated form CuSO4.X H20

The value of X (the number of moles of water molecules) can be

determined by heating the hydrated form to remove the

water molecules.

6.24 g of hydrated copper(II) sulfate was

heated until the blue colour was totally replaced

by white. The product was heated to constant

mass. The final mass of 3.99 g. Calculate the

value of X in the formula CuSO4.X H20

The final mass is less than the original mass. This is due to the molecules of water being

removed from the hydrated compound

Mass of water = 6.24 – 3.99 = 2.25 g

We now determine the number of moles of copper(II) sulphate and water in the hydrated

compound

CuSO4 H2O

Mass 3.99 2.25

GFM 159.6 18

Moles 0.025 0.125

Dividing both numbers by the smallest (0.025) will give the

simplest whole number mole ratio

0/025/0.025 = 1 0.125/0.025 = 5

So the value of X is 5 and the formula of the compound is CuSO4.5 H20

mass

n GFM

The second method in gravimetric analysis is PRECIPITATION. The analyte is dissolved in

water and converted into an insoluble product by the addition of a suitable reagent. The

resulting precipitate is then filtered, washed, dried and finally weighed.

A sample of sodium sulphite, Na2SO3, contains sodium sulphate, Na2SO4, as the only

impurity.

6.35 g of the sample was dissolved in deionised water and an excess of barium chloride was

added. Both barium sulfite, BaSO3, and barium sulfate, BaSO4 precipitated but on addition

of hydrochloric acid the barium sulfite dissolved. The precipitate of barium sulfate was

filtered, washed and dried. The mass was found to be 1.67 g

Na2SO4(aq) + BaCl2(aq) BaSO4(s) + 2NaCl(aq)

Calculate the % mass of barium sulfite in the sample

Moles of barium sulfate = mass/gfm = 1.67/233 = 0.00717

Moles of sodium sulfate = 0.00717 (1:1 ratio in equation)

Mass of sodium sulfate = n x gfm = 0.00717 x 142 = 1.02 g

Mass of sodium sulfite = 6.35 – 1.02 = 5.33 g

% sodium sulfite = 5.33/6.35 x 100 = 83.9%

Remember that solubility data is available in the data booklet. Precipitates are insoluble

compounds and a check of the data book should help identify them.

Two common precipitation reactions are

Ag+(aq) + Cl-(aq) Ag+Cl- (s)

and

Ba2+(aq) + SO42-(aq) Ba2+SO4

2-(s)

Gravimetric determination of nickel using dimethylglyoxime

Gravimetric analysis can be used to determine the nickel content of a nickel(II) salt.

This can be achieved by reacting the nickel(II) ions with dimethylglyoxime

(butanedione dioxime) in the presence of a slight excess of ammonia:

The complex, nickel(II) dimethylglyoximate, is filtered from the reaction mixture,

dried and weighed.

Requirements

500 cm3 beaker hydrated nickel(II) chloride (NiCl2.6H2O),sintered glass crucible ,

2 mol l–1 ammonia, Buchner flask and adapter, 0.1 mol l–1 dimethylglyoxime in ethanol,

water pump, 2 mol l–1 hydrochloric acid, desiccators,

balance (preferably accurate to 0.001 g), weighing bottle, hot plate, steam bath,

measuring cylinders (10 cm3 and 100 cm3), thermometer, stirring rod, dropper,oven

Hazcon


immediately.

Hydrated nickel(II) chloride is harmful by inhalation and by ingestion.

Wear gloves.

Dimethylglyoxime in ethanol is irritating to the eyes and is highly flammable.

2 mol l–1 ammonia irritates the eyes.

Procedure

1. Transfer approximately 0.5 g of hydrated nickel(II) chloride to a weighing bottle and

weigh the bottle and contents.

2. Add about 20 cm3 of deionised water to a 500 cm3 beaker and transfer the bulk of the

nickel salt to the water.

3. Reweigh the bottle with any remaining salt.

4. Stir the mixture until the solid dissolves and add about 20 cm3 of 2 mol l–1 hydrochloric

acid. Dilute the mixture with deionised water to about 200 cm3.

5. Heat the solution to 70–80°C on a hot plate and add approximately 50 cm3 of 0.1 mol l–1

dimethylglyoxime in ethanol.

6. Add 2 mol l–1 ammonia solution dropwise and with constant stirring until a permanent

red precipitate is obtained. Add a further 5 cm3 of the ammonia solution to provide a

slight excess. In all, you should have added about 30 cm3 of ammonia solution.

7. Heat the beaker and contents on a steam bath for about 30 minutes and when the

precipitate has settled test the clear liquid for complete precipitation by adding a few

drops of the dimethylglyoxime and ammonia solutions. (If more red precipitate appears

then add about 5 cm3 of 0.1 mol l–1 dimethylglyoxime solution followed by about 3 cm3

of 2 mol l–1 ammonia solution.)

8. Remove the beaker from the steam bath and allow it to cool to room temperature.

9. Dry the sintered glass crucible in an oven at 120°C, allow it to cool in a desiccator and

then weigh it.

10. Set up the filtration apparatus: sintered glass crucible, Buchner flask and adapter.

Filter off the precipitate at the water pump and wash the precipitate with a several

portions of deionised water.

11. Dry the crucible and precipitate in the oven at 120°C for about 1 hour and then

transfer them to a desiccator.

12. Once they have cooled to room temperature, reweigh the crucible and contents.

13. Heat the crucible and contents to constant mass, ie reheat for about 15

minutes in the oven at 120°C, cool in the desiccator and reweigh until two

successive readings are within 0.01 g of each other as the balance available is

only accurate to 0.01 g.

Formula of hydrated nickel(II) chloride _________________________

Initial Mass of weighing bottle + nickel chloride = __________

Final mass of weighing bottle + nickel chloride = __________

Mass of nickel chloride = __________

Mass of empty crucible = __________

1st mass of crucible and precipitate = __________

2nd mass of crucible and precipitate = __________

3rd mass of crucible and precipitate = __________

4th mass of crucible and precipitate = __________

a. Every measured result has an uncertainty associated with it. Read the pages in your

reference booklet which discuss combining uncertainties and then calculate the

uncertainties in the mass of nickel chloride and the mass of the precipitate. Your

answer should be in the form ------ {Mass +/- uncertainty)

b. In step 7 of the procedure why is dimethylglyoxime added to the clear liquid?

c. What does heating to constant mass mean?

d. Why is a desiccator used in this experiment?

e. Calculate the theoretical percentage by mass of nickel in NiCl2.6H2O

f. Use the results of your experiment to calculate the percentage by mass of nickel in

the sample of the hydrated nickel(II) chloride. Do your results match the theoretical

value.

1. Before 1947, ‘silver’ coins were made from an alloy of silver,copper and nickel. To

determine the metal composition, a coin weighing 10.00 g was dissolved in nitric

acid and the resulting solution diluted to 1000 cm3 in a standard flask.

A 100 cm3 portion was treated in the following way.

Hydrochloric acid (0.20 mol l–1) was added to this solution until precipitation of

silver(I) chloride was complete. The precipitate was recovered by filtration. It

was washed and dried and found to weigh 0.60 g.

.

a. (i) Calculate the percentage by mass of silver in the coin.

(ii) How could you tell when precipitation was complete?

b. The filtrate was treated to reduce the copper(II) ions to copper(I) ions. Ammonium

thiocyanate solution was added to precipitate the copper as copper(I) thiocyanate:

Cu+(aq) + CNS–(aq) CuCNS(s)

After filtration, drying and weighing, the precipitate was found to weigh 0.31 g.

Calculate the percentage by mass of copper in the coin.

2. Crystals of hydrated sodium carbonate left exposed to the atmosphere gradually

lose some of their water of crystallisation. The formula of the crystals may be given

by Na2CO3. xH2O, where x has a numerical value between 0 and 10.

16.0 g of the crystals was dissolved in water and made up to 250 cm3 of solution in a

standard flask. To determine the value of x in the formula, 25 cm3 of the sodium

carbonate solution was titrated with 1.0 mol l–1 hydrochloric acid. 15.0 cm3 of the acid

was required for neutralisation.

{Hint: Carbonate ions react with hydrochloric acid in a 1:2 ratio}

a. Calculate the mass of sodium carbonate (Na2CO3) in 16.0 g of the crystals.

b. Find the value of x in the formula Na2CO3.xH2O.

3. A 5.02 g sample of a silver alloy was analysed as follows. The sample was completely

dissolved in excess dilute nitric acid and then treated with an excess of sodium

chloride solution, causing a white precipitate of silver(I) chloride.

The precipitate had a mass of 2.37 g.

Calculate the percentage mass of silver in the alloy.

The visible light we observe is a only small part of the electromagnetic spectrum. (We will

discuss this in course notes – (refer to the sections on electronic structure and transition

metals in the inorganic unit).

The wavelengths of light visible to the human eye range from around 400 nm (blue light) to

700 nm (red light). Note that nm indicate nanometres – 1 nm = 1 x 10-9 metres.

A coloured solution will absorb some, but not all, of the white light that shines on it.

The light that is not absorbed will pass through the

solution (we say the light is transmitted) and combine

to give the colours we see.

For example, if a solution absorbs the

blue part of white light then the light

that is transmitted appears yellow.

Conversely, if yellow light is absorbed

then the solution will have a blue colour.

We say that blue and yellow are each

other’s complementary colour: each is

the colour that white light becomes when

the other is removed.

Complementary colours are shown diagonally

opposite each other in the colour wheel. 475 nm

Wavelengths are approximate.

510 nm

650 nm

While the colour of a solution depends on the colour of

light it absorbs, the intensity of its colour depends on

the concentration of the solution: the more

concentrated the solution, the darker its colour,

ie the more light it absorbs. We can get some

idea of the amount of light a coloured solution

absorbs by using a colorimeter

Colorimetric determination of manganese in steel

Colorimetry is an analytical technique used to determine the concentrations of coloured

substances in solution. It relies on the fact that a coloured substance absorbs light of a

colour complementary to its own and the amount of light it absorbs (absorbance) is

proportional to its concentration.

Colorimetry is particularly suited to the determination of manganese in steel because the

manganese can be converted into permanganate ions, which are coloured.

The conversion is achieved in two stages. Using nitric acid, the managanese is first oxidised

to manganese(II) ions, which are then oxidised to permanganate ions by the more powerful

oxidising agent, potassium periodate.

Requirements

standard flasks (50 cm3 and 100 cm3), steel paper clips, 50 cm3 burette, colorimeter,

standardised 0.0010 mol l–1 acidified potassium permanganate, 2 mol l–1 nitric acid,

optically matched cuvettes, 85% phosphoric acid,balance (accurate to 0.001 g),

acidified potassium periodate, glass beakers (50 cm3 and 250 cm3),

(5 g potassium periodate per 100 cm3, Bunsen burner, heating mat and tripod,

2 mol l–1 nitric acid), measuring cylinders (50 cm3 and 10 cm3),

potassium persulfate, clock glass, propanone, filter funnel, deionised water, tweezers,

anti-bumping granules, wash bottle, dropper, wire cutters.

Hazcon

Wear eye protection and if any chemical splashes on your skin wash it off immediately.

The acidified 0.0010 mol l–1 potassium permanganate is harmful if ingested and irritates the

eyes and skin. Wear gloves. nBoth 2 mol l–1 nitric acid and its vapour are corrosive and toxic,

causing severe burns to the eyes, digestive and respiratory systems. Wear gloves.

85% phosphoric acid is corrosive: it burns and irritates the eyes and skin. It is a systemic

irritant if inhaled and if swallowed causes serious internal injury. Wear gloves.

Acidified potassium periodate solution is harmful if swallowed and is an irritant to the eyes,

skin and respiratory system. It is also corrosive. Wear gloves.

Potassium persulfate is harmful if swallowed or inhaled as a dust. It irritates the eyes, skin

and respiratory system, causing dermatitis and possible allergic reactions. Wear gloves.

Propanone is volatile and highly flammable, and is harmful if swallowed. The vapour irritates

the eyes, skin and lungs, and is narcotic in high concentrations. Wear gloves.

Procedure

Part A – Calibration graph

Rinse the burette, including the tip, with 0.0010 mol l–1 acidified potassium permanganate

and fill it with the same solution.

1. Run 2 cm3 of the permanganate solution into a 50 cm3 standard flask and make up to

the graduation mark with deionised water.

2. Stopper the flask and invert it several times to ensure the contents are completely

mixed.

3. Rinse a cuvette with some of the solution and fill it.

4. Rinse and fill a second cuvette with deionised water

5. Zero the colorimeter (fitted with a green filter) with the deionised water.

6. Measure the absorbance of the coloured solution in the cuvette.

7. Repeat steps 2 to 6 with 4, 6, 8, 10, 12 and 14 cm3 of the permanganate stock solution

in the burette.

Part B – Conversion of manganese to permanganate

1. Degrease a steel paper clip by swirling it with a little propanone in a beaker. Using

tweezers remove the paper clip and leave it to dry for a minute or so on a paper towel.

2. Cut the paper clip into small pieces.

3. Place a weighing boat on the balance and then tare the balance. Weigh accurately

about 0.2 g of the paper clip pieces and transfer them to a 250 cm3 glass beaker.

4. Add approximately 40 cm3 of 2 mol l–1 nitric acid to the beaker and cover it with a

clock glass.

5. Heat the mixture cautiously, in a fume cupboard, until the reaction starts. Continue

heating gently to maintain the reaction, but remove the source of heat if the reaction

becomes too vigorous.

6. Once the steel has reacted, allow the solution to cool a little. Add a couple of anti-

bumping granules and then boil the solution until no more brown fumes are given off.

7. Once this solution has cooled considerably – no more than ‘hand hot’ – add about 5 cm3

of 85% phosphoric acid, approximately 0.2 g of potassium persulfate and a couple of

fresh anti-bumping granules. Boil the mixture for about 5 minutes.

8. To this solution, add approximately 15 cm3 of acidified potassium periodate solution

plus a couple of fresh anti-bumping granules and then gently boil the mixture. The

solution will start to turn pink. Continue gently boiling until the intensity of the pink

colour remains constant. This should take about 5 minutes.

9. Allow the pink solution to cool to room temperature and then transfer it to a 100 cm3

standard flask, leaving the anti-bumping granules in the beaker.

10. Rinse the beaker several times with a little deionised water and add the rinsings (but

not the anti-bumping granules) to the flask.

11. Make up the solution to the graduation mark with deionised water.

12. Stopper the flask and invert it several times to ensure the contents are completely

mixed.

13. Using a colorimeter fitted with the appropriate green filter, measure the absorbance

of the solution.

Mass of paper clips (tared balance) = _________________

Absorbance of paper clip solution = ____________________

a. Why is colorimetry a suitable technique to measure the mass of manganese in a paper

clip?

b. What was the wavelength of the filter used in the colorimeter.

c. Why was this filter chosen for use in the analysis?

Volume of 0.001 mol l-1

permanganate solution/cm3

Concentration of permanganate

solution/mol l-1

Absorbance

2

4

6

8

10

12

14

d. Why was the colorimeter zeroed with deionised water before any of the permanganate

solutions were analysed?

d. Prepare a calibration graph of using the results of the standard permanganate

solutions (You may wish to use Microsoft excel to do this)

e. Find the concentration of permanganate in the paper clip solution. If you used excel,

obtain the equation of the line and use this to find the concentration of the solution.

If you drew the graph manually draw line on the graph to indicate how you arrived at

the concentration value. Note, it is very important your graph is of a suitable size and

scale.

f. The formula of the permanganate ion is MnO4-. The ratio of manganese to

permanganate is therefore 1:1.

Use this information and your result from part (e) to calculate the mass of manganese

in the paperclip and hence the % mass of manganese in the paperclip.

1. An experiment was carried out to determine the % manganese in a sample of stainless

steel. A series of standard permanganate solutions were prepared from a stock

solution of 0.001 mol l-1 solution of permanganate and used to produce the

calibration graph shown below.

a. One of the standard solutions used to prepare the calibration graph had a

concentration of 1 x 10-4 mol l-1. Describe, giving appropriate volumes and apparatus,

how this solution could be prepared from the 0.001 mol l-1 solution of permanganate.

b. The results of the experiment are shown below.

Mass of steel used = 0.19 g

Absorbance of permanganate solution = 0.25

Total volume of permanganate solution = 100 cm3

Use the graph and the results to calculate the percentage, by mass, of manganese in

the sample of stainless steel.

2. Margaret analysed a sample of 100 cm3 of contaminated water for its copper content.

Suitable chemicals were added to the water to produce a blue coloured copper

compound, and the intensity of the colour was measured using a colorimeter. A series

of 5 solutions with known copper concentrations were prepared in a similar way and the

colorimeter reading produced by each was recorded as shown in the table below.

a. The solutions appeared blue because they were absorbing part of the visible region of

the electromagnetic spectrum. Which colours of the spectrum were being absorbed?

b. Suggest the most appropriate coloured filter for use in the colorimeter and explain

your choice.

Concentration of copper/

milligrams per litre (mg l-1)

Absorbance

0 0.00

50 0.12

100 0.27

150 0.40

200 0.53

c. Draw a calibration graph of the colorimeter results.

d. The colorimeter reading for the contaminated water sample was 0.35. Calculate the

number of moles of copper in the water sample.

3. The diagram shows white light passing through a cyan coloured solution.

a. The steps shown below outline how the concentration of the cyan solution could be

determined by colorimetry. Put the steps in the correct order.

b. Explain why the absorbance graph shown below is NOT that of the cyan coloured

solution.

Practical organic chemistry is primarily concerned with synthesising (making) organic

compounds and the purpose of a ‘synthesis’ is to prepare a pure sample of a specified

compound. Essentially, there are five steps involved:

preparation – the appropriate reaction is carried out and a crude sample of the

desired product is prepared

isolation – the crude sample of the product is separated from the reaction mixture

purification – the crude product is purified

identification – the identity of the pure compound is confirmed

calculation of the percentage yield.

The next two experiments will explore these steps in detail. You will need to refer to the

reference booklet regularly to help with the questions in this section of the booklet.

Preparation of aspirin

Introduction

Aspirin (acetyl salicylic acid) is an analgesic (pain-killing), anti-inflammatory and antipyretic

(fever-reducing) drug. It is an ester and can be prepared by the condensation reaction

between 2-hydroxybenzoic acid (salicylic acid) and ethanoic anhydride:

In this experiment you will prepare aspirin then separate it from the reaction mixture using

vacuum filtration. The aspirin will be purified by recrystallisation. The purity of the

prepared aspirin will be checked by measuring its melting point and by thin-layer

chromatography.

Requirements

50 cm3 conical flask, 2-hydroxybenzoic acid, 100 cm3 conical flasks, 85% phosphoric acid,

measuring cylinders (10 cm3 and 50 cm3), ethanoic anhydride,250 cm3 glass beakers,

ethanol, thermometers, anti-bumping granules, dropper, deionised water, glass stirring rod,

ice, balance (accurate to 0.01 g), hot plate, iodine, Buchner funnel and flask,

dichloromethane, water pump, ethyl ethanoate, filter papers, clock glass, oven,

capillary tubes, melting point apparatus, chromatography chamber,TLC plate, test-tubes,

UV lamp, pure aspirin.

Hazcon


2-Hydroxybenzoic acid is harmful by ingestion, causing nausea, vomiting etc. It is also a

severe skin and eye irritant. Wear gloves. Ethanoic anhydride is corrosive. The liquid

irritates and burns the eyes and skin severely while the vapour irritates the respiratory

system and may cause bronchial and lung injury. It is also flammable. Wear gloves and

handle in a fume cupboard. 85% phosphoric acid is corrosive: it burns and irritates the skin

and eyes. It is a systemic irritant if inhaled and if swallowed causes serious internal injury.

Wear gloves. Aspirin irritates the eyes and skin. Ethanol is volatile, highly flammable,

irritating to the eyes and intoxicating if inhaled or ingested. Dichloromethane irritates the

eyes and skin and is at its most harmful if inhaled. Wear gloves. Ethyl ethanoate is irritating

to the eyes, volatile and can irritate the respiratory system. It is highly flammable. Wear

gloves.

Procedure – aspirin preparation and separation.

1. Weigh a 50 cm3 conical flask and to it add about 5 g of 2-hydroxybenzoic acid.

Reweigh the flask and its contents.

2. In a fume cupboard, add 10 cm3 of ethanoic anhydride from a measuring

cylinder to the 2-hydroxybenzoic acid. During the addition, swirl the contents

of the flask to ensure thorough mixing.

3. Add five drops of 85% phosphoric acid to the mixture, again with swirling.

4. Place the flask on a hot plate (in the fume cupboard) and heat the mixture to

about 85°C. Keep it at this temperature for about 10 minutes and constantly

stir the mixture.

5. Cool the mixture in an ice/water bath and then pour it into approximately

150 cm3 of cold water contained in a 250 cm3 beaker.

6. Filter off the precipitate at the water pump and wash it thoroughly with

several portions of cold water.

7. Transfer the crude product to about 15 cm3 of ethanol in a 100 cm3 conical

flask. Add a couple of anti-bumping granules and heat the mixture gently on a

hot plate until it dissolves.

8. Pour this solution into a 100 cm3 conical flask containing about 40 cm3 of water.

If an oil forms, reheat the mixture on a hot plate to dissolve it. If the oil

persists, add a few drops of ethanol and reheat the mixture.

9. Set aside the mixture and allow it to cool to room temperature.

10. Filter off the crystals of aspirin at the water pump and wash them with a small

volume of cold water. Allow air to be drawn through the crystals for a few

minutes in order to partially dry them.

11. Weigh a clock glass and transfer the crystals to it. Dry the crystals in an oven

at about 100°C and then reweigh the clock glass and crystals.

Procedure – identification and purity

1. Determine the melting point of the aspirin product.

2. Take a TLC plate and using a pencil lightly draw a line across the plate about 1

cm from the bottom. Mark two well-spaced points on the line.

3. Place small amounts (about a third of a spatulaful) of your aspirin product and a

pure sample of aspirin in two separate test-tubes.

4. Add about 1 cm3 of solvent (a 50:50 mixture of ethanol and dichloromethane) to

each of the test-tubes to dissolve the aspirin samples.

5. Use capillary tubes to spot each of the two samples onto the TLC plate. Allow

to dry and repeat two or three more times.

6. After the spots have dried, place the TLC plate into the chromatography

chamber, making sure that the pencil line is above the level of the solvent

(ethyl ethanoate). Close the chamber and wait until the solvent front has risen

to within a few millimetres of the top of the plate.

7. Remove the plate from the chamber, immediately marking the position of the

solvent front, and allow it to dry.

8. Place the TLC plate in a beaker containing a few iodine crystals and cover the

beaker with a clock glass. Once any brownish spots appear, remove the plate and

lightly mark the observed spots with a pencil. Alternatively, observe the dried

TLC plate under UV light and lightly mark with a pencil any spots observed.

9. Calculate the Rf values of the spots. This will give you some indication of the

purity of the aspirin you have prepared.

Mass of conical flask = ___________________________________

Mass of conical flask and 2-hydroxybenzoic acid = ______________

Mass of 2-hydroxybenzoic acid ___________________________

Mass of clock glass = ____________________________________

Mass of clock glass and aspirin = ____________________________

Mass of aspirin _______________________________________

a. Calculate the number of moles of 2-hydroxybenzoic acid used in the

experiment.

b. Use the value calculated in (a) to find the theoretical mass of aspirin which

should be produced.

c. Calculate the % yield of aspirin and give reasons why it is not 100%

Theoretical melting point of aspirin = ______________________

Melting point of prepared aspirin = ________________________

Draw a diagram of the developed TLC plate and calculate the Rf values of any spots. Your

diagram should show how the Rf values were determined

a. Using the results of the TLC and melting point comment on the purity of your

product.

Preparation of cyclohexene from cyclohexanol

Introduction

Cyclohexene can be prepared by dehydrating cyclohexanol using concentrated phosporic

acid. The product can be separated from the reaction mixture by distillation, and after

purification it can be weighed and the percentage yield determined.

Procedure

See animation of this experiment

1. Calculate the percentage assuming 20.00 g of cyclohexanol produced 4.10 g of

cyclohexene.

2. List the chemicals that are present in the separating funnel

and explain why a separating funnel is being used.

3. What test could be used to show the product is unsaturated.

A list of “learning outcomes” for the topic is shown below. When the topic is

complete you should review each learning outcome.

You are expected to consult your teacher if you have any

issues with the learning outcomes.

State that a standard solution is a solution of known concentration

Know how to weigh chemicals by the weighing by difference method.

State that a solution may be standardised by using a primary standard.

State the properties required of a primary standard.

Be able to name some primary standards.

State that sodium hydroxide is not a primary standard and state reasons why.

State the procedure required to prepare a standard solution.

Be able to carry out simple uncertainty calculations and be aware of the

uncertainties in laboratory equipment.

State that a control experiment allows the validity of a technique to be

evaluated and to be able to suggest chemicals which can be used as controls.

State that volumetric analysis involves measuring volumes and the main

technique is titration.

Sate what a back titration is and state the reasons it may be necessary to

carry out a back titration.

State that gravimetric analysis involves weighing reactants and products.

Be able to explain what is meant by heating to constant mass.

State that colorimetry can be used to analyse the concentration of coloured

solutions.

State that when carrying out a colorimetric analysis the filter which provides

maximum absorption for the test substance is always chosen.

Need Help

Understand

Revise

Be able to carry out stoichiometric calculations based on both volumetric and

gravimetric analysis.

Be able to carry out percentage yield calculations

State that recrystallisation is a technique used to purify a solid organic chemical

based on its solubility.

State that the melting point of a chemical is one way to identify the chemical

and it can be used to evaluate the purity of a chemical.

State that mixed melting point involves taking the pure form of the product

and mixing it with the prepared compound. If the melting point of the mixture

is the same as the pure product then the prepared compound is identical to the

pure compound.

State that thin – layer chromatography (TLC) can be used to identify

chemical and assess their purity.

State that the technique of solvent extraction (using a separating funnel) is

based on the relative solubility of a compound in two different immiscible

liquids, usually water and an organic solvent. It is a separation technique.

State that refluxing is a technique used to apply heat energy to a chemical

reaction mixture over an extended period of time which avoids the loss of a

substance through evaporation.

Vacuum filtration using a Buchner or sintered glass funnel. This method is

carried out under reduced pressure and provide a faster means of separating

the precipitate from the filtrate.

State that distillation is the process of heating a liquid until it boils, capturing

and cooling the resultant hot vapours, and collecting the condensed vapours.

In the modern organic chemistry laboratory, distillation is a powerful tool, both

for the identification and the purification of organic compounds.

I have discussed the work of this section with my teacher!

Date. __________________________________

Pupil Signature. __________________________

Teacher Signature. _______________________

http://en.wikipedia.org/wiki/Miscibility

http://en.wikipedia.org/wiki/Liquid

http://en.wikipedia.org/wiki/Water_(molecule)

http://en.wikipedia.org/wiki/Solvent

good use of the “researching chemistry reference … unit covers the key areas of gravimetric...

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