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University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 1: MELTING POINT DETERMINATION Introduction The determination of melting points in capillary tubes is a very simple, effective and rapid way to confirm the identity and purity of organic solid substances. The process of ‘melting’ is usually characterized by:

(a) Pronounced softening or shrinking. (b) Most of the substance suddenly running down and forming a meniscus. (c) All the substance becomes liquid.

Although in some cases (a)-(c) occur within a range of 1oC, melting usually occurs (especially impure samples) over a temperature range of c. 5-10oC. In such cases, this range should be recorded (rather than estimating an 'average' reading). Note that even in such case we still refer to the ‘range’ as the melting ‘point’. Extensive databases now exist listing melting points of pure substances including:

o The data collection of the National Institute of Standards and Technology (NIST) which is available online on http://webbook.nist.gov/;

o The Organic Compounds Database maintained by Colby College available online on http://www.colby.edu/chemistry/cmp/cmp.html;

o The Sigma-Aldrich catalogue available online on http://www.sigma-aldrich.com/ o ChemFinder from Cambridge Software, available online from

http://chemfinder.cambridgesoft.com/ o The CRC Chemistry and Physics databook (available in the library) o The Materials Safety Data Sheets (MSDS, may be accessed from manufacturers

websites, see website for details.). Published melting points are usually within a 0 - 360oC temperature range and refer to values taken in capillary tubes sealed at one end*. It is interesting to note that melting point values for the same material quoted from different data sources may be slightly different. One should also bear in mind that the melting point ranges listed in the company catalogues (e.g. on the Sigma-Aldrich website) are the melting points of the compounds as they are sold, i.e. if the compounds are sold slightly impure, then the company catalogue melting point will deviate from those listed in the CRC data book or on the MSDS which always refer to the melting point of the pure compound.

* Your laboratory demonstartor should show you how to prepare these capillary tubes sealed from one end from a normal capillary tube.

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The recommended procedure for melting point determination is to first place a small amount of the sample on a clean dry watch-glass or glass slide and then powder it by means of a spatula or the rounded end of a glass rod. You should then fill four capillary tubes each to a depth of 2-4 millimetres with the powdered sample. (This is best performed by tapping the open end of the capillary tube into the powdered sample and then tapping the sealed end on the bench to move the powder down the capillary tube. This filling/tapping may be repeated if necessary until you obtain a loose but continuous column of c. 3mm in depth at the sealed end of the capillary tube.) One of these tubes is then placed in the melting point apparatus and the temperature is raised at a relatively fast rate (about 10 °C/minute) until the compound first begins to melt. This temperature is probably an over-estimation of the true melting point and should only be treated as a rough estimate. The melting point apparatus should then be allowed to cool to about 20 °C below that estimated temperature, and the process is repeated with another sample, this time using a much slower heating rate (less than 1oC per 30 seconds for the last 3 or 4 degrees). This reading should provide you with the first accurate melting point result. At least one other accurate reading should then taken using the third sample. If these two accurate readings are inconsistent, then a third, and possibly more determinations have to be performed until the results become consistent. The melting point should then be taken as the average range. For example, if the three accurate readings were made, and these are a1-b1, a2-b2 and a3-b3 then the average range is from (a1+a2+a3)/3 to (b1+b2+b3)/3 . NOTE: You should always discard melting point tubes once they have been used. If the melting point obtained is doubtful repeat using a fresh tube. Procedure: NOTE: Before you commence the practical work, please ‘read’ carefully the Safety Data for A –E and M1 provided in the table below. You have been provided with six samples A – E and M1, the molecular formula and melting points for some of which are given below. Sample Molecular

Formula Melting pt. (oC) *

Risks / Safety data†

A C7H7NO2 54 Toxic, R23/24/25, R36/37/38, S45, S26, S22, S36/37/39

B C7H6O2 Harmful, R22, R41, R37/38, R42/43, S26, S36, S22

C C7H6O3 159 Toxic, R61, R22, R41, R37/38, S45, S26, S36/37/39, S22

D C7H5NO4 239 Harmful, R22, R41, S26, S39

E C9H8O2 131-134 Irritant, R36/37/38, S26, S36

M1 not available similar to one / a combination of A-E (*) Melting points quoted directly from the information supplied by manufacturer of the respective

chemicals on the labels of the original containers. † Please refer to the RISK and SAFETY terms and symbols on the course website, http://staff.um.edu.mt/jgri1/teaching/che1302

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You should first determine the meting point of substances B and M1 individually using the procedure outlined above. Enter these results in your laboratory notebook. (After the practical, you should the search the online Organic Compounds Database maintained by Colby College to try and identify substances A-E. Report the name and structural formula of these compounds in your lab reports. You may also wish to consult other databases to check whether there are any differences in the reported values, etc.) You should find that substance M1 has a melting point that is very similar to one of the other five samples. This similarity in melting points may indicate that substance M1 is of the same chemical composition as one of the samples A-E (henceforth referred to M2). You can test this hypothesis by (1) Preparing a ‘new’ sample (=M3) by grinding mixture of equal amounts of M1 and M2, (2) Filling three separate tubes with M1, M2 and M3 respectively, and (3) placing them together in the melting point apparatus with the mixture in the middle.

If M1 and M2 are identical, then M1, M2 and M3 should all melt simultaneously. However, if this is not the case, M3 will melt over a range of temperature lower than M1 and M2. In such case, you should determine the melting point of M3 more accurately using the procedure outlined above and enter the value in your laboratory notebook. Once you have convinced yourself that your results are consistent you should fill in the R-CHE1700-01 form attached (University Copy) and hand it to the laboratory demonstrator. The demonstrator should also check your laboratory notebook and fill in the Chemistry Practical Front Matter sheet. Note that this procedure should be repeated for every practical.

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University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 2: RECRYSTALLISATION AS A TECHNIQUE FOR PURIFICATION Introduction The purification of organic compounds is one of the commonest procedures encountered in experimental organic chemistry. This is because it is very rare that a product from a chemical reaction can be extracted in a pure form. The common practice to purify a product which is crystalline in its pure form is through a process of one / a series of crystallisation / recrystallisation process/es, In its simplest form, the recommended procedure for recrystallisation is as follows*:

1. An inert solvent is chosen such that the substance to be crystallised is only slightly soluble in cold solvent but very soluble in hot solvent. (This is discussed at a later stage.)

2. The impure substance to be re-crystallized is dissolved in the minimum amount of hot solvent (at or near its boiling point). Note: (i) Please refer to page 3 for a discussion on how to choose a solvent. (ii) The only case when it is justified (and recommended) to use more solvent than the minimum amount is when there is a very steep solubility curve, i.e. when slight cooling causes considerable crystallisation.

3. The hot solution is filtered rapidly (to remove any insoluble particles e.g. dust) using a stemless funnel and a fluted filter paper (see Fig. 1) into a conical flask. Note that both the funnel and the flask should be warmed just a little before filtration by passing them over a flame so as to prevent premature crystallisation. (If this occurs, the crystals have to be re-dissolved.)

4. The conical flask containing the hot filtered solution is then corked lightly and then left undisturbed to cool down very slowly hence allowing the dissolved substance to crystallise out. Note that slow cooling favours formation of large crystals, as opposed to quenching which will result in the formation of a powder. If an ‘oil’ emerges rather than crystals, one should induce crystallisation by scratching the inside of the conical flask with a glass rod. If this is unsuccessful, one may (a) inoculate (seed) the solution with some of the solid materials, or, (b) cool the solution to lower temperature (e.g. by immersing in an acetone/dry ice bath.

* This procedure has to be slighly modified if the amount of sample to be purified is less than 1g and expensive to obtain (e.g. the final product of a multi-step reaction). In such cases, the filtartion in step (2) should be avoided unless absolutely necessary whilst the filtration in step (5) is carried out using a Willstatter ‘filtration nail’. For further details on such techniques, or on other modifications in special cases, please refer to specialist textbooks, such as Vogel’s Textbook of Practical Organic Chemistry, 5th Edn. (Longman, 1988).

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5. Once the hot solution has cooled down to room temperature, the conical flask with the solution is immersed in an ice bath to maximise the extraction of the substance from the supernatant solution.

6. The crystals are then separated from the supernatant solution by filtration. This filtration is usually carried out under suction (see fig. 2a) using through a Buchner funnel fitted with a filter paper wetted with the solvent (for large quantities of crystals) or a Hirsch funnel (for medium quantities see Fig. 2b).

7. The crystals are then washed from any remaining impure supernatant solution using minimum amount of ice-cold pure solvent (taking care that the crystals do not dissolve). To do this one must (i) discontinue the suction, (ii) pour small amounts of the chilled solvent, (iii) stir cautiously the solvent/crystals to ensure that the solvent washes all the crystals, and (iv) re-apply suction to ‘dry’ the crystals. Note that this ‘drying’ may be aided by pressing the crystals on the sieve using a clean spatula/flattened glass rod, and by leaving the ‘suction process’ running for an extra few minutes.

8. The solid crystals are then transferred to a watch glass and dried to constant weight in a desiccator under vacuum. Other drying methods include the use of suitably heated ovens or even by placing on a water bath if the substance melts over 100oC (although this technique may result in contamination of the crystals by steam).

9. The DRY crystals are tested for purity (e.g. by a melting point verification, or by thin layer chromatography). If the substance is still found to be impure, then the whole process is repeated from step 2 until purity can be ascertained or successive recrystallisations result in very little change in the purity (indicated for example by consistency in the melting point values).

Fig. 1: To prepare a fluted filter paper you must first fold a normal filter paper in half and again in quarters, and then open it up as shown in (a). The edge 2,1 is then folded on to 2,4 and edge 2,3 on to 2,4, hence producing new folds at 2,5 and 2,6 when the paper is opened. The folding is continued, 2,1 to 2,6 and 2,3 to 2,5, thus producing folds at 2,7 and 2,8 respectively (b); further 2,3 to 2,6 giving 2,9, and 2,1 to 2,5 giving 2,10 (c). The final operation consists in making a fold in each of the eight segments - between 2,3 and 2,9, between 2,9 and 2,6, etc. - in a direction opposite to the first series of folds, i.e. the folds are made outwards rather than inwards as first. The result is a fan arrangement (d), and upon opening, the finished fluted filter paper. (e). (Taken from Vogel’s Textbook of Practical Organic Chemistry, 5th edn.).

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Fig. 2: (a) The set-up to be used for filtration under suction using a Buchner funnel, (b) An illustration of a Hirsh funnel which should be used for crystal quantities in the range of 1-2g. Note that the smaller versions of Hirsch funnels will accommodate filter papers 3-4mm in diameter. Note also that apart from the porcelain versions of the Buchner and Hirsch funnels where the diameter of the sieve pores are in the millimetre region, one may also find sintered glass funnels available in a number of porosities (coarse, medium and fine). The success behind purification by recrystallisation relies on the fact that the impurities are present in much lower quantities that the substance to be purified, and not necessarily that the impurities are more soluble than the substance to be purified. To illustrate this we present here idealised calculations for two particular (simplified) examples of recrystallisation where we want to extract A from two 10g mixtures of A/B containing (w/w):

(1) 95% A and 5% B (the impurity), see Table 1, and, (2) 90% A and 10% B (the impurity), see Table 2.

For each of these two examples we shall consider two scenarios: (i) When A is more soluble in cold solvent than the impurity B, and, (ii) When the impurity B is more soluble in cold solvent than A.

In the discussion, we shall assume that both A and B are freely soluble in hot solvent and that 100mL of solvent is used in each recrystallisation. The calculations presented in Tables 1 and 2 illustrate two important ideas:

1. In all cases, it is possible to obtain A in a pure form, even when irrespective on whether A is more soluble than B in cold solvent, or not.

2. The small change of a 5% increase in the percentage of the impurity B will result in a substantial increase in (a) effort required to obtain 100% pure A in terms of the number of recrystallisations required, and, (b) loss of percentage recovery of A (and we are assuming that there are no losses due to experimental errors/limitations!).

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Scenario:

(i) (ii)

Substance: A B A B Solubility in cold solvent (g / 100mL) 0.6 0.3 0.3 0.6 1st Recrystallisation: Initial percentage composition (%): Initial amount present (g) Amount dissolved in solution at room temperature (g) Amount recovered (g) Final percentage composition: (%)

95 9.5 0.6 8.9 97.2

5 0.5 0.3 0.2 2.8

95 9.5 0.3 9.2 100

5 0.5 0.5 (all) 0 0

2ns Recrystallisation: Initial percentage composition (%): Initial amount present (g) Amount dissolved in solution at room temperature (g) Amount recovered (g) Final percentage composition (%):

97.2 8.9 0.6 8.3 100

2.8 0.2 0.2 (all) 0 0

not required

not required

Total % of A recovered: 87.4% 96.8% Table 1: Example 1 - Initial composition is 95%A, 5%B Scenario:

(i) (ii)

Substance: A B A B Solubility in cold solvent (g / 100ml) 0.6 0.3 0.3 0.6 1st Recrystallisation: Initial percentage composition (%): Initial amount present (g) Amount dissolved in solution at room temperature (g) Amount recovered (g) Final percentage composition: (%)

90 9.0 0.6 8.4 92.3

10 1.0 0.3 0.7 7.7

90 9.0 0.3 8.7 95.6

10 1.0 0.6 0.4 4.4

2ns Recrystallisation: Initial percentage composition (%): Initial amount present (g) Amount dissolved in solution at room temperature (g) Amount recovered (g) Final percentage composition (%):

92.3 8.4 0.6 7.8 95.1

7.7 0.7 0.3 0.4 4.9

95.6 8.9 0.6 8.3 100

4.4 0.4 0.4 (all) 0 0

3rd Recrystallisation: Initial percentage composition (%): Initial amount present (g) Amount dissolved in solution at room temperature (g) Amount recovered (g) Final percentage composition (%):

95.1 7.8 0.6 7.2 98.4

4.9 0.4 0.3 0.1 1.6

not required

not required

4th Recrystallisation: Initial percentage composition (%): Initial amount present (g) Amount dissolved in solution at room temperature (g) Amount recovered (g) Final percentage composition (%):

98.4 7.2 0.6 6.6 100

1.6 0.1 0.1 (all) 0 0

not required

not required

Total % of A recovered: 73.3% 92.2% Table 2: Example 2 - Initial composition is 90%A, 10%B

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Choice of Solvent The most desirable characteristics of a solvent for recrystallisation are as follows:

1. The solvent must be inert, i.e. it must not react chemically with the substance to be purified;

2. The substance to be purified must be highly soluble in the solvent at elevated temperatures and ‘insoluble’ at room temperature;

3. The solvent should dissolve the impurities readily or to only a very small extent (so that the impurity may be removed in the first filtration using the fluted filter paper);

4. The solvent should allow the formation of well-formed crystals of the purified compound;

5. The solvent must have a relatively low boiling point so that it may be easily removed from the crystals of the purified compound;

6. Ideally, the solvent should not be toxic, flammable and expensive. Some common solvents used in recrystallisation are listed in Table 3 below (broadly in the order of decreasing polarity). Solvent b.p. (°C) Comments: Water (distilled) 100 To be used whenever suitable Methanol* 64.5 Flammable, toxic Ethanol 78 Flammable Industrial spirit 77-82 Flammable Rectified spirit 78 Flammable Acetone 56 Flammable Ethyl acetate 78 Flammable Acetic acid (glacial) 118 Not very flammable, pungent vapours Dichloromethane (methylene chloride)* 41 Non-flammable; toxic Chloroform* 61 Non-flammable; vapour toxic Diethyl ether 35 Flammable, avoid whenever possible Benzene*+ 80 Flammable, vapour highly toxic Toluene*+ 110 Flammable, vapour toxic Carbon tetrachloride* 77 Non-flammable, vapour toxic Light petroleum (petroleum ether, see note) 40-60** Flammable Cyclohexane 81 Flammable Table 3: Common solvents for recrystallisation (Taken from from Vogel’s Textbook of Practical Organic Chemistry, 5th edn.). Note: (*) To be used in a fume hood, (+) Benzene is now known to be extremely toxic. Use toluene instead whenever possible. (**) See note on petroleum ether on page 6. It should be noted that:

• The use of ether (diethyl ether) as a solvent for recrystallisation should be avoided wherever possible, partly owing to its great flammability (and its potential to form explosive peroxides with prolonged exposure to air, an important factor to consider when handling ‘aged’ ether) and partly owing to its tendency to creep up

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walls of the containing vessel, thus depositing solid matter by complete evaporation instead of preferential crystallisation.

• The vapours of methanol, dichloromethane, benzene, toluene and carbon tetrachloride are toxic and therefore recrystallisations involving their use must he conducted in an efficient fume cupboard: excessive inhalation of any vapour should be avoided.

• Toluene is much less toxic than benzene and should be used in place of benzene whenever possible.

• Note on petroleum ether: Petroleum ether (or light ether) is a mixture of alkanes and is available in different compositions labelled according to the boiling points. Available fractions include b.p. 40-60, 60-80, 80-100 and 100-200 oC. When the boiling point exceeds 120oC the fraction is usually called 'ligroin'. Pure pentane (b.p. 36oC), hexane (b.p. 69oC), and heptane (b.p. 98oC), are also frequently used as recrystallisation solvents, although these are much more expensive than the mixtures.

There are no hard and fast rules on the choice of solvents. In general, it should be noted that a substance is likely to be most soluble in a solvent to which it is most closely related in chemical and physical characteristics. In particular, a polar substance is more soluble in polar solvents and less soluble in non-polar solvents (“like dissolves like”). It is also common practice to use a mixture of two or more solvents when one cannot find a single solvent with the right characteristics. In practice the choice of a solvent for recrystallisation must be determined experimentally if no information is already available. The recommended procedure to do this is as follows:

1. About 0.lg of the powdered substance is placed in a small test tube 2. The cold solvent is added a drop at a time with continuous shaking of the test tube.

If the sample dissolves easily in 1 mL of cold solvent the solvent is unsuitable. 3. If after about 1 mL of the solvent has been added, the substance is still

undissolved, then the mixture is heated to boiling to check the solubility in hot solvent. If the sample dissolves easily upon gentle warming, the solvent is un-suitable.

4. If all the solid does not dissolve, more solvent is added in 0.5mL portions, and again heated to boiling after each addition. If 3 mL of solvent is added and the substance does not dissolve on heating, the substance is regarded as sparingly soluble in that solvent, and hence unsuitable.

5. If the compound dissolves in the hot solvent, the tube is cooled to determine whether crystallisation occurs. (Note: If crystallisation does not take place rapidly, you should check if this is due to the absence of suitable nuclei sites by scratching the tube below the surface of the solution with a glass rod. The fine scratches on the walls and the minute fragments of glass produced may serve as excellent nuclei for crystal growth.) If crystals do not separate (even after scratching for several minutes and cooling in an ice-salt mixture), the solvent is considered as unsuitable.

6. If crystals separate, the amount of these should be noted. 7. This process may be repeated with other solvents (using a fresh test tube for each

experiment) until the best solvent is found. For each experiment, the approximate proportions of the solute and solvent giving the most satisfactory results should be

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recorded. Note that if no solvent is found to be satisfactory, one should attempt using mixed solvents (see note below).

Note on the use of mixed solvents:

If the substance is found to be far too soluble in one solvent and much too insoluble in another solvent to allow recrystallisation, mixed solvents or solvent pairs may be used. The two solvents must, of course, be completely miscible. (Miscible solvents must not have very different polarities.) The procedure for use in such cases is as follows:

o The compound is dissolved in the solvent in which it is very soluble.

o The hot solvent, in which the substance is only sparingly soluble, is added cautiously until a slight turbidity is produced.

o The turbidity is then just cleared by the addition of a small quantity of the first solvent and the mixture is allowed to cool to room temperature. This should result in crystal formation.

Pairs of liquids which may be used include: alcohols and water; alcohols and toluene; toluene and light petroleum; acetone and light petroleum; diethyl ether and pentane; glacial acetic acid and water; dimethylformamide with either water or toluene.

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Procedure In this practical you shall first check the suitability of the following eight solvents

o water o hexane o acetone

o ethanol o dichloromethane o ethyl acetate

o petroleum ether 60-80oC o toluene

for recrystallisation of: (A) Benzoic acid (B) Acetanilide, and (C) Naphthalene

and then purify benzoic acid (m.pt. 122oC) from a mixture of benzoic acid and 5% naphthalene (w/w). In particular:

(a) Before the lab session, you should look up: I. The safety information for all the chemicals to be used from the MSDS

datasheets and fill in the special assessment forms where appropriate.

II. The boiling points of the solvents and the meting points of A-C. (b) Test the suitability of the different solvents for recrystallisation of A, B and C by

following a modified (because of time constraints) version of the procedure on page 6, i.e.:

i. Place about 0.2g of the powdered substance in a small test tube; ii. Add 1mL of cold solvent & check for solubility;

iii. Bring to the boil using the water baths provided† and check for solubility in hot solvent;

iv. Allow the mixtures to cool (slowly) and check for the formation of crystals.

(c) Having identified the best solvent to use for the purification of benzoic acid though recrystallisation from a mixture of benzoic acid and 5% naphthalene (w/w), you should place about 1g (weighed) into a flat-bottomed flask and carry out one recrystallisation procedure as outlined on pages 1-2. The only modification (because of time) is that you should dry your crystals (step 9) first between two filter papers, and then in a flameless oven set at a temperature in between the boiling point of the solvent and the melting point of the benzoic acid. Having dried the solvent you should:

i. Measure the percentage extraction of benzoic acid. ii. Check the purity of your crystals through a melting point

determination. (Recall that pure benzoic acid melts at 122oC.)

† Because all (apart from water and dichloromethane) of the solvents are flammable, you should not heat the solvents on a naked flame. Note that all the organic solvents you shall be using apart from toluene (b.pt. 110oC) have boiling points less than 100oC.

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University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 3: INVESTIGATING THE BROMINATION OF ACETANILIDE – An Example of an Electrophilic Aromatic Substitution Introduction Bromine reacts with the aromatic amide acetanilide (N-phenylethanamide, 1) in a substitution mode in which one or more hydrogen atoms bonded to the benzene ring are replaced by bromine atoms.

NH

CH3

O

(1)

The reaction can be described as follows:

3 6 5 2 3 6 5-CH CONHC H Br CH CONHC H Br HBrn nn n+ → + Although at a first glance it may appear that a large number of products can result from this reaction, in practice only one major product forms when this reaction is carried out at room temperature*. For example, meta monosubstituted products can be excluded since the–NHCOCH3 group is ortho / para directing. Table 1 lists the melting points of the possible products. It may be observed that these different potential products have very different melting points, and hence, the identity of the reaction product should be determined without difficulty even if the crude compound is used in the melting point determination†.

2-Bromoacetanilide 99 4-Bromoacetanilide 165.4 2,4-Dibromoacetanilide 145.4 2,6-Dibromoacetanilide 208-209 2,4,6-Tribromoacetanilide 232

Table 1: The Melting Points (oC) of possible bromination products (from Beilstein).

* The Risks/Safety phrases for this product (taken from the Fisher Scientific (Canada) website, http://www.fishersci.ca/ ) are listed in Appendix 1. † The presence of impurities, (e.g. any unreacted acetanilide, or other minor reaction products) will result in a lowering of the melting point. Thus, unless properly purified, the product of bromination will most probably not melt at the exact temperature listed in the table.

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Experimental Procedure

(a) Before your practical, you should look up the safety information for acetanilide, glacial acetic acid, bromine liquid (used to prepare bromine solution) and ethanol from the MSDS datasheets and fill in the special assessment forms where appropriate (see http://staff.um.edu.mt/jgri1/safety/download ).

(b) Carry out the reaction using the following procedure: i. Place 0.5g of acetanilide (accurately weighed) and 5mL of glacial acetic acid

in a boiling tube which is secured using a clamp and placed in a beaker containing water at room temperature.

ii. Stir carefully but effectively the acetanilide/glacial acetic acid mixture until all the solid acetanilide dissolves.

iii. Add with stirring, 6mL bromine solution and continue stirring for another 20 minutes.

iv. At the end of this period, quench the reaction by adding 30mL of water and 6mL of saturated sodium hydrogen sulfite solution, again keeping the mixture well stirred. Continue stirring until the red colour (indicating an excess of the bromine) disappears.

v. Cool the tube in an ice-water bath for 15 minutes. vi. Collect the product by vacuum filtration, wash with cold distilled water and

air-dry. Determine the weight of crude product, and hence determine the percentage yield.

(c) Product characterization: Determine the melting point of the product and hence identify it using the data provided in Table 1.

(d) If you have time, recrystallise your product from ethanol and repeat the melting point determination.

(e) In your discussion, use the information given in Appendix 2, and other information you may regard are important to explain the outcome of this experiment.

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Appendix 1: The Risks/Safety information for the main product of this reaction (taken from the Fisher Scientific (Canada) website, http://www.fishersci.ca/ European Labeling in Accordance with EC Directives Hazard Symbols: Not available. Risk Phrases: Safety Phrases: S 24/25 Avoid contact with skin and eyes. S 28A After contact with skin, wash immediately with plenty of water. S 37 Wear suitable gloves. S 45 In case of accident or if you feel unwell, seek medical advice immediately (show the label where possible). Appendix 2: Orientation of Nitration in Substituted Benzenes (at 25oC)‡

HNO3 / H2SO4

X X

NO2

Meta-directing deactivators

X % Ortho % Meta % Para

( )+

3 3N CH− 2 89 11

2NO− 7 91 2

CN− 22 77 2

COOH− 17 81 2

2 3COOCH CH− 28 66 6

3COCH− 26 72 2

CHO− 19 72 9

‡ Taken from Jojn McMurry, ‘Organic Chemistry’, 3rd ed., Brooks/Cole Publsihing Company, USA (1992), pg. 579.

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Ortho/para-directing deactivators

X % Ortho % Meta % Para

F− 13 1 86

Cl− 735 1 64

Br− 43 1 56

I− 45 1 54

Ortho/para-directing activators

X % Ortho % Meta % Para

3CH− 63 3 34

OH− 50 0 50

3NHCOCH− 19 2 79

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University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 4: SEPARATION BY EXTRACTION Introduction Separation by extraction is a technique which allows the separation of compounds based on their difference in solubility in two immiscible solvents. If a solute X is allowed to distribute itself between two immiscible solvents, A and B, then the following equilibrium will be established: ( ) ( )X A X B!!"#!! The equilibrium constant for this process is known as the distribution coefficient or the partition coefficient, KD, or BDA and is given by:

( ) ( ) [ ][ ]

BB A

A

Xthe concentration of X in phase B X Xthe concentration of X in phase A XDK D= = =

Similarly, for solutes X and Y distributed into two immiscible solvents, A and B, we may define the following equilibria: ( ) ( )X A X B!!"#!!

( ) ( )Y A Y B!!"#!! and for effective separation by extraction of X and Y using solvents A and B, we require that: ( ) ( )X YD DK K>> in which case most of X will find itself in solvent B and most of Y will find itself in solvent A, or, ( ) ( )X YD DK K<< in which case most of X will find itself in solvent A and most of Y will find itself in solvent B. Solvents A and B are usually water (the aqueous phase) and a water-immiscible organic solvent, such as diethyl ether respectively (henceforth referred to as ether), di-isopropyl ether, toluene, dichloromethane and light petroleum. The selection of the organic solvent will depend upon the solubility of the substance to be extracted in that solvent and upon the ease with which the solvent can be separated from the solute. Diethyl ether is very frequently used since:

1. It is very immiscible with water; 2. It has very powerful solvent properties for non-polar solutes; 3. It has a very low boiling point (35 °C) hence it can be evaporated easily

after the extraction

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However, special care must be taken when handling this solvent since it is extremely volatile and flammable*. In particular, all naked flames in the immediate vicinity should be extinguished. Extractions are carried out using a separatory funnel (conical, as in Fig. 1, or pear-shaped) which should have a short stem and fitted with a ground glass interchangeable stopper. The size of the separatory funnel must be about twice the volume to be extracted. It is usually convenient to have this separatory funnel mounted in a ring on a stand with a firm base. It is also important to check that the stopcock of the funnel is not jammed. For best performance, the barrel and plug of the stopcock are dried with a linen cloth and in the case of glass stopcocks lightly treated with a suitable lubricant such as silicon grease or petroleum jelly.

Fig. 1: A conical separating funnel The solution and the extraction solvent (in this case exemplified by ether/water) are introduced into the funnel, and with the stopper firmly held in place the funnel is shaken gently so as to avoid a quick build-up of the internal pressure. Furthermore, it is essential to periodically invert the funnel and to open the stopcock to relieve the excess pressure. This gentle shaking / relieving of pressure should be continued until the atmosphere in-side the funnel is saturated with ether vapour. At this stage, further shaking develops little or no additional pressure. The separating funnel is then shaken vigorously for 2-3 minutes to ensure the maximum possible transfer of the solutes to the appropriate phase and * The fire hazard is reduced also by employing di-isopropyl ether (b.p. 67.5 °C), but this solvent is much more expensive than diethyl ether.

3

returned to the stand to allow the mixture to settle. When two sharply defined layers have formed, the lower aqueous layer is run off and separated as completely as possible. However, care must be taken not to run off drops of the ethereal phase. In fact, it would be appropriate to run down the last few drops of the aqueous phase and the first few drops of the ethereal phase into a small beaker so as to avoid mutual contamination of the two phases. The separated solutes are then removed from solution. This can either be done by evaporation of the solvent (especially for the organic phases where the solvents are usually very volatile), or by precipitation. The solutes are then dried, purified by recrystallisation, and tested for purity using a melting point determination or a chromatographic method.

Procedure For this practical, you have been provided with a mixture containing benzoic acid (50% by mass) and 4-nitrotoluene. You are required to:

!"Isolate these two components through extractions; !"Purify the separated samples through recrystallisations; !"Determine the purity of your separated samples through a melting point

determination. The procedure to be used is as follows:

1. Before your practical, you should fill you should look up the safety information for the chemicals to be used, i.e.: benzoic acid, 4-nitrotoluene, diethyl ether, sodium bicarbonate solution, concentrated hydrochloric acid and ethanol from the MSDS datasheets and fill in the special assessment forms where appropriate. (see http://staff.um.edu.mt/jgri1/safety/download).

2. To separate the benzoic acid and 4-nitrotoluene you should: a. Place 2g benzoic acid / 4-nitrotoluene mixture in a 100mL conical flask,

add 20 mL diethyl ether (no flames!) and shake. Transfer this mixture to a separating funnel and add 20mL sodium bicarbonate solution (5% by mass). This will convert the benzoic acid to the water soluble sodium benzoate.

b. Stopper the separating funnel and shake for a short while; then invert the funnel and open the tap gently to release pressure. Close the tap and continue shaking, frequently releasing the pressure until further shaking develops little or no additional pressure. Then shake vigorously (but carefully) for 2-3 minutes to ensure the maximum possible transfer of the solutes to the appropriate phase.

c. Allow the layers to separate and run off the bottom layer into a 100mL conical flask. Then add 10mL water to the funnel, shake and run the aqueous bottom layer into the same conical flask.

3. To recover benzoic acid and to test its purity: a. Add a few drops of concentrated hydrochloric acid to the aqueous phase

until the pH is less than 2. Allow to solution to stand until crystallisation of the benzoic acid is complete.

4

b. Filter the solution under suction and collect the crude benzoic acid crystals in a Hirsch funnel. Wash the benzoic acid crystals twice with water and allow to drain.

c. Recrystallise from water and dry in an oven set at 105oC. d. Determine the mass, percentage recovery and melting point of the

recovered/purified benzoic acid. 4. To recover 4-nitrotoluene and to test its purity:

a. Evaporate the ether solution to dryness over a steam bath in a fume hood, (no flames!).

b. Collect the crystals so formed and recrystallise from ethanol c. Determine the mass, percentage recovery and melting point of the

recovered/purified 4-nitrotoluene. 5. In your discussion:

a. Explain carefully the chemistry involved in the separation. b. Discuss how would you modify the above protocol to deal with the

separation of a mixture of benzoic acid and 4-nitrophenol.

1

University of Malta Department of Chemistry

CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 5: SEPARATION INLOVING THE FORMATION OF ADERIVATIVE OF ONE OF THE COMPONENTS

Introduction

The separation of a binary mixture of substances with very similar physical properties (e.g. two similar hydrocarbons) cannot be accomplished by simple crystallisation or solvent extraction. In such cases, one needs to:

1. Chemically modify one of the components to render it significantly different from the other constituents;

2. Separate the derivative from the mixture; 3. Transform the derivative back to the original form.

This practical involves the separation of naphthalene, 1, (m.pt. 80oC) and biphenyl, 2,(m.pt. 69oC).

1 2

Although these two chemicals are extremely similar, 1 (a Lewis base) reacts with picric acid to form a picrate, 3, which 2 does not. The separation is made possible since 3 has different solubility properties from 1 and 2.

Procedure:

1. Before your practical, you should fill you should look up the safety information for the chemicals to be used, i.e.: naphthalene, biphenyl, picric acid and concentrated ammonia and fill in the special assessment forms where appropriate.(see http://staff.um.edu.mt/jgri1/safety/download ).

2. To form the naphthalene derivative (naphthalene picrate) and isolate it from biphenyl:

a. Place 1g of the naphthalene (50% w/w) and biphenyl mixture in a boiling tube, add 2mL of methanol and heat on a steam bath until the mixture is completely dissolved.

b. In another boiling tube dissolve 1.3g of picric acid in 10mL methanol by warming on a steam bath until a yellow solution is obtained.

c. Transfer the solution of the mixture slowly to the picric acid solution. Warm for 1 minute on a steam bath and allow to cool. This will result in the formation of solid naphthalene picrate.

d. Collect the crystals formed in a Hirsch funnel, wash with methanol, dry and determine the melting point.

c. After putting aside a few drops of the picric acid solution for step 2e, to the rest of the picric acid solution add slowly the solution of the mixture from step 2a. Warm for 1 minute on a steam bath and allow to cool. This will result in the formation of solid naphthalene picrate. d. Collect the crystals formed in a Hirsch funnel, wash with methanol, dry and determine the melting point.

2

e. Check for the presence of naphthalene in the filtrate by adding a few drops of fresh picric acid solution. If no picrate forms then separation has been complete.

3. To re-extract biphenyl, purify it and test its purity: a. Add 1mL concentrated ammonia (using a measuring cylinder) and 30mL

water to the filtrate. This will result in the precipitation of biphenyl. b. Collect the biphenyl by filtration and recrystallise from hot water. c. Record its melting point.

4. To re-convert the picrate to naphthalene, purify it and test its purity: a. Place the picrate crystals in a 100mL conical flask, and dissolve them in a

little methanol by warming on a steam bath. b. Add 1mL concentrated ammonia and 30mL water. Filter the naphthalene

and recrystallise it from methanol. c. Record its melting point.

1

University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 6: FRACTIONAL DISTILLATION Introduction Distillation is a very useful method of purifying liquids. It is used routinely for purifying solvents and reagents and, with care and the correct apparatus it can be used to separate liquids whose boiling points are only a few degrees apart. The conventional apparatus for use in simple distillation is shown in Fig. 1a. This apparatus is suitable for distilling compounds from involatile residues, or for separating liquids whose boiling points differ by at least 50°C. The separation of liquids whose boiling points are between 5°C and 50°C apart requires the use of a fractionating column (see Fig. 1b). In its most basic form, a fractionating column may be regarded as a long vertical tube with a temperature gradient (hottest at the bottom). In practice, the use of a very long vertical tube is very inconvenient and hence it is normally replaced with a shorter ‘column’ designed to offer the same properties (e.g. the Vigreux column in Fig. 1b).

(a) Taken from M. Casey, J. Leonard, B. Lygo and G. Procter,

Advanced Practical Organic Chemistry, 1990, Blackie & Professional (UK) (b) Taken from A.I. Vogel, Vogel’s Textbook of Practical Organic

Chemistry, 5th Edn., 1988, Longman (UK) Fig. 1: (a) The setup for a simple distillation, and (b) The setup for a fractional distillation using a Vigreux column.

2

Fig. 2: The temperature-composition diagram, corresponding to an ideal mixture A-B where A is more volatile than component B. In fractional distillation, successive boilings and condensations of a liquid originally of composition a1 lead to a condensate that is pure A. Let us illustrate the process of distillation with reference to an A-B ideal* liquid mixture where A is more volatile than B (i.e. A has a lower boiling point than B, see Fig. 2). We shall assume that the initial liquid mixture has a composition a1 and boils† at a temperature of T2. At T2, an equilibrium between the liquid and the vapour phases is established, but whilst the liquid phase still has a composition of a1 (=a2), the vapour phase has composition a2` which is richer in the more volatile component A. In a simple distillation, the vapour of composition a2` is withdrawn and condensed. Thus, in this example, the first drop of the condensate would have given a liquid of composition a3 (the same as a2

`, i.e. a mixture which is richer in the more volatile component, A, than the original liquid mixture of composition a1). In fractional distillation, the vapour of composition a2` is not collected, but it is allowed to cool slowly as it ascends the fractionating column, until eventually it returns to the condensed phase of composition a3 (the same composition as the vapour a2`). However, as this condensate descends the column, it exchanges heat with the ascending hot vapour and the vapour is enriched with the more volatile component at the expense of the liquid, in an attempt to reach equilibrium within the liquid-vapour system. These boiling-condensation cycles are repeated successively until ‘almost pure’ A is obtained at the top of the column. * In reality, some substances do not form ideal solutions, and non-ideality will lead to the formation of azeotropic mixtures, which results in discrepancies from the theory described here. For a discussion on distillation of non-ideal solutions, please refer to standard physical chemistry textbooks, such as P.W. Atkins, Physical Chemistry, 6th Edn., 1998, OUP (UK), Chapter 8. † At a given pressure, the boiling point is the temperature when the liquid is in equilibrium with the vapour.

3

All this suggests that the conditions necessary for a good separation are:

o Comparatively large amounts of liquid continually returning through the column; o Thorough mixing of liquid and vapour; o A large active surface of contact between liquid and vapour.

This means that cooling in the fractionating column must occur very slowly. Excessive cooling (e.g. because of drafts) should be avoided by, for example, insulating or lagging the outer surface of the column or, if possible, by surrounding it with a vacuum jacket or an electrically heated jacket. This difficulty is particularly important when separating liquids of high boiling points. It is important to ensure that the procedure is carried out safely. In particular:

1. Never heat a closed system. If you are worried about moisture entering your system, you can attach a drying tube to the vents.

2. Take special care when distilling organic solvents to avoid fires and explosions. Therefore:

a. Ensure that the vapour does not come into contact with flames, sources of sparks (electrical motors), or very hot surfaces (hot plates).

b. Do not heat (especially round bottomed flask) using a Bunsen burner – Always use an electrically heated oil/water bath or a heating mantle.

c. Never allow a distillation pot to boil dry. The residues may ignite or explode.

d. Beware of the possibility that ethers and hydrocarbons may be contaminated with peroxides that form on exposure with air.

e. Always check whether any of the liquids you plan to distil are thermally unstable.

3. Always add some boiling chips (or if possible, stir the liquid using a magnetic stirrer), in order to prevent bumping.

Procedure: In this practical, you have shall be separating the two organic liquids, acetone and toluene from a 1:1 mixture of the two, determine their boiling point, and determine which of the two is the carbonyl compound (acetone) through the 2,4-dinitrophenylhydrazine (2,4-DNPH) and the iodoform tests.

1. Before your practical, you should fill you should look up the safety information for the chemicals to be used and fill in the special assessment forms where appropriate. (see http://staff.um.edu.mt/jgri1/safety/download). You should also look up the boiling points of pure acetone and toluene.

2. Separate about 50mL of an acetone/toluene mixture and measure the boiling points of acetone and toluene:

a. Set up the apparatus for fractional distillation as in Fig. 1b using a 100mL round-bottomed flask and the column provided (insulated with cotton wool). Note that (1) the bulb of the thermometer should be just below the

4

level of the side arm, (2) a very small piece of glass wool (not cotton wool) should be used in the fractionating column to support the glass beads.

b. Place about 50mL of the acetone/toluene mixture in the 100mL round-bottomed flask together with a few boiling chips (porous pot etc.).

c. Heat the flask slowly until the reading on the thermometer reaches a steady state (the boiling point of the first fraction) and drops are observed to condense out of the Leibig condenser‡. Record this steady state temperature and collect the distillate in a conical flask. Allow the distillation to proceed until no more liquid gets out of the condenser/measuring cylinder. Record the volume of the first fraction.

d. When all of the first fraction has been distilled, the temperature at the top of the column will be observed to increase and then reach a second steady state (the boiling point of the second fraction) and whist drops of the second fraction will start to condense out of the leibig condenser‡. Record this second steady state temperature (the boiling point of the second fraction) and collect the distillate in a clean conical flask/measuring cylinder§. You should collect as much of the second fraction as practically possible making sure that you so not allow the distillation pot to boil dry. Record the volume of the second distillate.

3. Identify which of the two liquid distillates has a carbonyl group (acetone) using the 2,4-DNPH and iodoform tests:

a. For the 2,4-DNPH test: Add 2-3 drops of the liquid to be tested to 3mL of 2,4-dinitrophenylhydrazine and shake. If no precipitate forms immediately, allow to stand for 5-10 minutes. Record the observations and inferences in your laboratory notebook.

b. For the iodoform test: Dissolve 4 micro drops of the liquid to be tested in water (2mL). Add 10% sodium hydroxide (2mL), and then slowly add 2mL of the iodine solution. If the substance is insoluble in water, dissolve it in dioxane (2mL). Proceed as above and at the end dilute with water (10mL). Record the observations and inferences in your laboratory notebook.

‡ You should collect in a separate container (and then discard it), any liquid collected (i) before the boiling point of the first distillate, and (ii) as the temperature is rising between the boiling point of the first fraction and that of the second fraction. § Since you have a binary A-B mixture, the second main distillate can be obtained without the use of a fractionating column.

1

University of Malta Department of Chemistry

CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 7/8: THE OXIDATION OF CYCLOHEXANOL TO CYCLOHEXANONE & ITS EXTRACTION THROUGH STEAM DISTILLATION.

Introduction

This practical involves: i. The oxidation of cyclohexanol to cyclohexanone using acidified sodium

dichromate; ii. The extraction of the cyclohexanone so formed from the rest of the reaction

mixture though steam distillation (the distillate is a water/cyclohexanone mixture); iii. The extraction of cyclohexanone from the water/cyclohexanone mixture in ether*; iv. Simple distillation to purify the cyclohexanone (the boiling points of ether and

cyclohexanone are very different)*. (* To be performed in the second week)

Steam distillation is a technique where a water immiscible substance A with a high boiling point can be ‘distilled’ with the help of steam at a temperature that is lower than 100oC.

A substance A is immiscible with water (substance B) when the mutual solubilities of A and B are very small. At equilibrium, an A/B system will consist of two liquid phases and a single gaseous phase, i.e.:

(1) A liquid phase containing mainly A which is saturated with B (the low solubility of B in A means that this phase is almost pure A);

(2) A liquid phase containing mainly B which is saturated with A (the low solubility of A in B means that this phase is almost pure B); and

(3) A vapour phase where A and B coexist (A and B are miscible in the gaseous phase).

The total vapour pressure, p, for this A/B system is approximately given by: * *A Bp p p= +

where * *A B and p p are the vapour pressures of pure A and pure B. These vapour pressures

are functions of temperature, T, such that *ip increase as T increases (volatility increase

with an increase in temperature). The system will start boilingi (i.e. distilling) at the point when the total vapour pressure becomes equal to the atmospheric pressure, i.e. p = patm.This occurs at a temperature A-B

boilT that is less than either of AboilT or B

boilT , the boiling points of pure A and B respectivelyii.

i Boiling can be considered as the point when the dissolved substances are purged out of their solution. ii For boiling to progress sucessfully at a single temperature, there must be continuous presence of the saturated solutions ‘A in B’ and ‘B in A’. This ‘continuous presence’ is ensured at all times since these very dilute solutions can be very quickly replenished to their saturatiuon composoition aided by the vigorous agitation induced by boiling.

2

For example, assuming that B Aboil boilT T< , boiling of A/B must occur before the

temperature reaches BboilT (=100oC if B is water) since ( )*

B Bboil

atmp T p= i.e.:

( ) ( ) ( )* * *A B B B A B

boil boil boilatm atmp T p T p T p p+ = + >

There is however one important snag to this technique, namely that the composition of the vapour (and hence of the distillate) is in proportion to the vapour pressures of the components, i.e.:

( )( )

*A AA*

B B B

boil

boil

p Tnn p T

=

This means that substance of a high boiling point with a low volatility will distil in low abundance.

To illustrate this we shall look at a nitrobenzene (=A) and water (=B) mixture. These substances boil (i.e. their respective vapour pressures reaches 760 mmHgatmp = ) at 210-211oC (i.e. A 483 484 KboilT = − ) and 100oC (i.e. B 373 KboilT = ) respectively, but a mixture of the two boils at 99.4oC (i.e. A-B 372.4 KboilT = ), as illustrated in Fig. 1. This results in a reduction of the boiling point of approximately 111K. However, at 99.4oC,

*A 20 mmHgp ≈ and *

B 740 mmHgp ≈ , i.e.:

( )( )

*A AA*

B B B

20 1740 37

boil

boil

p Tnn p T

= ≈ =

i.e. ( )( )

AA

B B

RMM A 1 123 123RMM B 37 18 666

nww n

× ×= ≈ =× ×

i.e. the wt% of nitrobenzene in the distillate will only be:

A

A B

wt% 100% 15%ww w

⎛ ⎞= × ≈⎜ ⎟+⎝ ⎠

Nevertheless, the low wt% of nitrobenzene in the distillate is by far overweighed by the fact that the distillation could be carried out at 99.4oC rather than at the boiling point of nitrobenzene, which is 210-211oC.

Note that the technique of steam distillation is very important in natural product chemistry where it is employed on a industrial scale to extract essential oils from plantsiii.

iii See, for example: ‘A Guide to Aromatherapy’ web-site, http://www.fragrant.demon.co.uk/, for a list of oils used in aromatherapy and their method of production, ‘The English Chamomile Company’ web-site, http://www.chamomile.co.uk/ for a discussion of steam distillation in the extraction of chamomile, or http://www.distillation.co.uk/, a web-site for a company which specialises in the production of distillation equipment for use in the production of essential oils.)

3

Fig. 1: The vapour pressure vs. temperature curve for a nitrobenzene-water system. Note that whilst pure nitrobenzene boils at 483-484K, the when it is mixed with steam, it will boil at the much lower temperature of 372.4K when the vapour pressure of the nitrobenzene-water system reaches 760 mmHg (atmospheric pressure).

Procedure:

(a) Before your practical, you should look up the risks/safety information for the known chemicals used in this practical.

(b) In the first session of your practical you will be oxidising the cyclohexanol to cyclohexanone and extracting the cyclohexanone by steam distillation.

1) At the beginning of the practical, fill the steam generator (to be used for the steam distillation) with water (3/4 full) and heat it with a Bunsen burner. (Boiling of such a quantity of water requires a plenty amount of time!)

2) To oxidise the cyclohexanol to cyclohexanone:i. Dissolve 15g sodium dichromate dihydrate in 25mL 2M H2SO4 by

heating the mixture in a beaker. Once all the sodium dichromate dihydrate dissolves, cool the mixture to 15oC by immersing the beaker in ice.

ii. Prepare a mixture of 15g cyclohexanol in 10mL 2M H2SO4 (10mL). Cool in ice to below 15oC.

iii. Add the dichromate onto the cyclohexanol and stir. Note that the reaction is exothermic, and you will observe that the temperature starts to rise very rapidly. This increase in temperature is beneficial

4

as it increases the rate of the oxidation reaction. However, you should not allow the temperature to exceed 60oC.

iv. Keep the reaction going until the temperature begins to drop and the solution turns green, indicting that the reaction is ‘complete’. (This should take about 15 minutes from the addition of the start of the reaction). Allow the reaction mixture to stand for another 10 minutes.

3) To extract the cyclohexanone so formed by steam distillation:i. Pour the solution into a 250mL two-necked round-bottomed flask.

Rinse the reaction flask with water (100mL, split between three rinses) and add the washings to the distillation flask.

ii. Set up the steam distillation apparatus according to the instructions given by your demonstrator.

iii. In theory, the steam generated from the steam generator should be enough to heat the reaction mixture. In practice, this would take a long time, and hence you are advised to accelerate the process by heating up the round bottomed flask directly until the reaction mixture starts boiling. The steam distillation should then be allowed to proceed on its own.

iv. When about 50mL of distillate is collected, add another 50mL of distilled water to the round bottomed flask and continue distilling until no more oil passes over with the water. This is an indication that the steam distillation is ‘complete’.

v. Cover the distillate, label it, and store it in your laboratory cupboard for the next session.

(c) In the second session of your practical you will be extracting the cyclohexanone from the cyclohexanone-water distillate, and further purifying it through simple distillation.

1) To extract the cyclohexanone from the distillate:,i. Salt out the distillate by adding sodium chloride (approx. 20mL of

0.2g/mL solution) and extract with ether (25mL). ii. Wash the ether layer with 10% sodium hydroxide (25mL) and then

re-wash with the salt solution (approx. 25mL of 0.2g/mL solution). iii. Dry the ethereal extract over anhydrous sodium sulphate and filter

into a distillation flask. 2) To purify the extracted cyclohexanone:,

i. Set up the simple distillation apparatus according to the instructions given by your demonstrator.

ii. Distil the ether over a steam bath into a Buchner flask having a side-arm fitted with a rubber tubing leading well below the level of the bench. This precaution is essential when ether is being distilled because of the risks of fire and explosion (Air and ether vapour form an explosive mixture when the latter reaches a flame.).

iii. Once all the ether has been distilled off, heat the remaining mixture with a heating mantle to distil the cyclohexanone.

iv. Record the temperature of collection and the yield in grams and in moles. Calculate the percentage yield.

25mL

1

University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 9: DETERMINATION OF THE COMPOSITION OF A ‘BORAX / BORIC ACID' AQUEOUS SOLUTION. Introduction In this practical you shall be measuring quantitatively the concentration of boric acid and borax in a boric acid / borax solution through acid/base titrations. Boric acid, 1, is too weak an acid to be titrated quantitatively in an aqueous medium using a visual indicator. However, it can be titrated with standard alkali (NaOH) in the presence of certain polyols (e.g. propane-1,2,3-triol, 2, commonly known as glycerol), using phenolphthalein as indicator. This is because 1 and two molecules of 2 can form a complex, (borospiranic acid, 3, which may be regarded as a diester of glycerol) in which two rings are attached to the boron atom, in mutually perpendicular planes.

(eqn. 1)

Borospiranic acid is a strong enough acid to give a satisfactory end-point with strong aqueous alkali (e.g. NaOH) using phenolphthalein as indicator (see Appendix 1). Note that the overall stoichiometry of the reaction can be represented as follows (eqn. 2):

( ) ( )glycerine3 4B OH + NaOH NaBa OH →

(eqn. 2) (Question: Given that the borospiranic acid is a monobasic acid, which is the labile proton in the structure? Give a reason for your choice.)

2

The propable 3D structure of the proposed borospiranic acid complex, 3.

The sodium salt of boric acid, borax (4, Na2B4O7.10H2O, also known as sodium tetraborate decahydrate) is a weak base. It can be quantitatively titrated directly with standard acid (e.g. HCl) using methyl red as an indicator (see Appendix 1) as in eqn. 3:

2 4 7 2 3 3 2Na B O .10H O + 2HCl 4H BO + 2NaCl + 5H O →

(4) (1)

(eqn. 3) One should note that these two determinations are independent of each other, and hence, if performed carefully, they can be used to quantitatively measure the amounts of boric acid and borax present in a mixture of the two. Procedure: In this practical you shall be applying the theory described above to measure quantitatively the concentration of boric acid and borax in a aqueous solution containing both boric acid and borax solution. However, you shall first be practicing the titrations on the two components present on their own, and at the same time, use these ‘practice titrations’ to determine quantitatively the mass percentage (i.e. percentage purity) of the boric acid and borax as supplied.

(a) Before your practical, you should look up the risks/safety information for the known chemicals used in this practical.

(b) To determinate of the mass percentage of boric acid: Accurately weigh about 1g of sample and transfer it to a clean conical flask together with 50mL water. Carefully boil the solution for about 1 minute so that the boric acid dissolves completely and any carbon dioxide in the distilled water is completely exsolved. Be careful not to loose any solute during boiling. Cool, add about 15mL glycerol, and titrate with (about) 1M NaOH using phenolphthalein as indicator. (The actual molar concentration will be provided during the practical.) Check whether the titre values that you obtain conform to

3

the amount of boric acid used. Repeat this procedure two more times. Use the results of the titration to calculate the mass percentage (i.e. purity by weight) of the boric acid as supplied.

(c) To determine the mass percentage of borax: Dissolve about 1.2g of borax (accurately weighed) in about 10mL water, and heat to help dissolution. Cool and titrate with (about) 0.25M HCl using methyl red as indicator. Check whether the titre values that you obtain conform to the amount of borax used. Repeat this procedure two more times. Use the results of the titration to calculate the mass percentage (i.e. purity by weight) of the borax as supplied.

(d) To determine the composition of an aqueous solution labeled M containing borax and boric acid: Pipette 25mL of M into a conical flask and titrate with (about) 0.25M HCl using methyl red solution as indicator. Record the volume of acid required. Reserve the titration liquid and repeat this portion of the assay to obtain concordant results. This should give you a measure of the amount of borax present. To determine the amount of boric acid, boil the titration liquid to expel carbon dioxide. Cool the solution, add glycerol (15mL), and titrate with (about) 1M NaOH using phenolphthalein as indicator, ensuring that you obtain concordant results. Use these measurements to calculate the concentration of borax and boric acid in the solution mixture M.

Note: You are expected to provide a quantitative estimate of the errors involved in these determinations. A discussion on ‘Errors in Quantitative Analysis’ is provided on http://staff.um.edu.mt/jgri1/teaching/ch130/errors.pdf .

4

Appendix 1: pH Transition ranges and colour of some common indicators*

* Taken from G.D. Christain, Analytical Chemistry, 4th Ed., John Wiley & Sons, USA, 1996.

1

University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 10: PERMANGANATE METHODS Introduction Redox (reduction/oxidation) titrations are extremely important in analytical chemistry. In these reactions, one reactant (the oxidising agent or oxidant) will have a stronger affinity for electrons than a second species (the reducing agent or reductant) and will tend to extract electrons from it. It is convenient to analyse a redox reaction in terms of two half-reactions, one which describes the oxidation, and one which describes the reduction. For example, the equation for the oxidation of iron(II) ions by permanganate ions in an acidic solution (eqn. 1) may be viewed in terms of the two half-reactions: (i) the reduction of 4MnO − to 2Mn + (eqn. 2), and (ii) the oxidation of Fe2+ to Fe3+ (eqn. 3). These two half reactions show very clearly that the permanganate ion is the electron acceptor (i.e the oxidising agent) whilst the iron(II) ion is the electron donor (i.e. the reducing agent): The full reaction: 2+ - + 3+ 2+

4 25Fe + MnO + 8H 5Fe + Mn + 4H O→ (eqn. 1) The corresponding half reactions:

4- + 2+2MnO + 8H +5e Mn +4H O− → (eqn. 2)

2 35Fe 5Fe 5e+ + −→ + (eqn. 3) The ability (potential) of a substance to act as an electron acceptor (i.e. as oxidising agent) can be measured in terms of the magnitudes of the standard reduction potentials, E0, (sometimes referred to as standard electrode potentials) of the half reactions. The higher the standard reduction potential of a species, the better is its ability to act as an oxidising agent (i.e. be reduced). Standard reduction potentials are always quoted relative to the half reaction where two hydrogen ions accept two electrons to form a hydrogen molecule which is arbitrarily given a standard reduction potential of 0V (see Appendix 1). Tables of standard potentials can be found in reference books or physical /analytical chemistry textbooks. One should note that the standard reduction potential measures the relative intensity of the driving force for the reduction half reaction. This means that, for example, the half reactions 3 2Fe e Fe+ − ++ → and 3 25Fe 5e 5Fe+ − ++ → both have the same standard reduction potential, E0 = 0.771V. Thus referring to Appendix 1 and by re-writing eqns. 2 & 3 as reductions with their associated standard reduction potentials, we may observe that:

- + 2+4 2MnO + 8H +5e Mn +4H O− → E0 = 1.51V (eqn. 4)

3 25Fe 5e 5Fe+ − ++ → E0 = 0.771V (eqn. 5)

2

that is, the standard reduction potential of the permanganate ion is higher (more positive) than that of the iron(II) ions, and hence, the forward of eqn. 4 and the reverse of reaction eqn. 5 will take place, as indicted in eqns. 1-3. Note that for a redox reaction to go to completion there must be a difference in the standard reduction potentials of the two half reactions of at least 0.1V. This is an important requirement if a redox reaction is to be used in quantitative analysis. This practical involves a number of titrations against potassium permanganate, KMnO4, namely:

i. Oxalic acid against KMnO4 to standardize the KMnO4 solution; ii. Iron(II)sulphate against KMnO4 to determine the concentration of Fe++;

iii. Hydrogen peroxide against KMnO4 to determine the volume strength of the peroxide solution.

The standard potentials for the half reactions involved in these process (taken from Appendix 1) are as follows:

- + 2+4 2

2-2 2 4

3 2

2 2 2

E ,MnO + 8H +5e Mn +4H O(1) 1.51

(2) 2CO 2e C O -0.440(3) 0.771Fe e Fe(4) 0.682O 2H 2e H O

Vθ

−

−

+ − +

+ −

→

+ →

+ →+ + →

Since the standard reduction potential of the half reactions (2)-(4) are lower than that of (1) (i.e. the permanganate has the greatest tendency to be reduced), the reverse of reactions (2)-(4) occur in these titrations. This means that in all of these reactions, the potassium permanganate is acting as the oxidising agent. Note that in all of these reactions, there is no need of an indicator since a 0.02M solution of potassium permanganate is deep purple whilst the product of its reduction, Mn2+, is nearly colourless. Thus, before the endpoint, the purple colour of the permanganate ion is removed practically as soon as it is added (see below) whilst as soon as the titration is complete, a fraction of a drop of excess permanganate will impart a definite pink colour to the solution. The mechanism of the reactions between oxalate and permanganate ions is extremely complicated. The first few drops of permanganate react slowly with the oxalic acid but after a small amount of manganous salt has been formed, the reaction occurs almost instantaneously in hot solution. It appears that the Mn2+ ions catalyze the reaction between permanganate and oxalic acid. This can be verified by adding a small amount of manganese(II) salt at the very start of the titration. The permanganate is then reduced rapidly from the start. Although the titration of oxalic acid (obtained by dissolving sodium oxalate (Na2C2O4), previously dried at 100-110oC) against potassium permanganate is a standard technique to standardise permanganate solutions, special care must be taken since hot solutions of oxalic acid decompose slowly:

2 2 4 2 2H C O CO + CO + H O→ This decomposition reaction is catalysed by manganous salt and hence the titration should not be carried out too slowly as this could lead to an error (a too low titre value).

3

However, if the titration is too rapid, or there is insufficient stirring, the added permanganate will tend to spend a longer time than necessary in the presence of hot acid. This will result in the auto-decomposition of the still unreacted permanganate with the evolution of oxygen, and this will be reflected in a too high titre value. (Note that if the titration is carried out according to the standard procedure, these two side reactions will only occur to a negligibly small extent.) Hydrogen peroxide solution is an aqueous solution of hydrogen peroxide containing not less than 5% and not more than 7% w/v of H2O2, corresponding to about 20 times its volume of available oxygen. In the titration of hydrogen peroxide with permanganate, we observe a similar phenomenon as with oxalic acid, that is, the first few drops of permanganate are decolourised slowly but again the reaction is catalysed by adding a trace of a manganous salt at the very beginning of the titration. Procedure:

(a) Before your practical, you should look up the risks/safety information for the known chemicals used in this practical.

(b) Preparation of approximately 0.02M KMnO4: Dissolve 0.8g KMnO4 in distilled water and make up to 250mL. Transfer to a beaker and heat to about 90oC for 15 minutes. Use after cooling to room temperature.

(c) Standardisation of potassium permanganate: i. Prepare an approximately 0.04M sodium oxalate solution by dissolving

about 0.54g of sodium oxalate (Na2C2O4, previously dried at 100-110oC) in distilled water and making up to 100.0mL (using a 100mL volumetric flask). Calculate the accurate concentration of the sodium oxalate solution.

ii. Pipette 25mL of the 0.04M sodium oxalate solution to a 250mL conical flask and add 50mL 2M sulphuric acid.

iii. Heat to about 80oC and titrate slowly with permanganate whilst shaking the flask well. If the solution remains pink after adding the first drop or two, wait until it clears before continuing the titration.

iv. Continue the titration slowly until a faint permanent pink tinge is obtained. The temperature should not be less than 60oC at the end-point. If it falls below this during the titration, the solution must be reheated.

v. Repeat this process twice more, or until you obtain concordant results to obtain a mean value for the molar concentration of the permanganate. Calculate the molar concentration of the potassium permanganate.

(d) Determination of the percentage of iron(II) in iron(II)sulphate crystals using standardised potassium permanganate solution:

i. Weigh out accurately about 3.5g of iron(II)sulphate crystals. Dissolve, and make up to 100.0mL using a 100mL volumetric flask using 2M sulphuric acid.

ii. Pipette 25.0mL of iron(II)sulphate solution into a 250mL conical flask, add 25mL 2M H2SO4 and titrate with KMnO4 solution to the first permanent pink end-point. Repeat the titration two or three times.

iii. Calculate the percentage of ferrous iron in the crystals.

4

(e) Determination of the concentration and volume strength of a peroxide

solution using standardised potassium permanganate solution i. Dilute 10.0mL of the hydrogen peroxide solution (measured using 10mL

pipette or a burette) to 250mL with water in a volumetric flask. ii. Pipette 25.0mL of the diluted solution to a conical flask, add 5mL of

sulphuric acid (50% w/v) and titrate with the potassium permanganate solution to the first permanent pink end-point. Repeat the titration two or three times until you obtain concordant results.

iii. Calculate the molar concentration and the ‘volume strength’ of the original peroxide solution.

5

Appendix 1: Standard potentials at 298K in electrochemical order*

Others†:

2-2 2 42CO 2e C O−+ → E0 = -0.440 V

* Taken from P.W. Atkins, Physical Chemistry, 6th Ed., OUP, UK, 1998. † Taken from G.D. Christain, Analytical Chemistry, 4th Ed., J. Wiley & Sons, Inc., USA, 1986.

1

University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 11: DETERMINATION OF WATER HARDNESS Introduction This practical involves:

(i) The determination of the total, permanent and temporary hardness of two different samples of tap water through complexometric titrations;

(ii) The determination of the total calcium and magnesium content of two different tap water samples through complexometric titrations;

(iii) The assessment whether the two water samples are significantly different in terms of their hardness.

(a) Water hardness

Water hardness is a measure of the mineral levels (usually calcium and magnesium salts) dissolved in the water. Hard water requires more soap and synthetic detergentsi and contributes to scaling in boilers and other industrial and domestic equipment. Beacuse the concentrations of calcium and magnesium ions in natural water are usually much higher than the concentrations of the other ions, it is common practice to measure water hardness by measuring the concentrations of Ca++ and Mg++. Hardness can either be temporary or permanent, this being determined by the anions associated with the Ca2+/Mg2+ cations. If the anion is a carbonate or bicarbonate, the hardness is temporary and can be removed by boiling the water and allowing it to cool. Sulphates and chlorides result in permanent hardness that cannot be removed by boiling. This type of hardness will not contribute to scaling but will cause waste of soap and detergent. (Temporary hardness that will contribute to scaling and soap precipitation). Hardness is measured as the equivalent to the amount of calcium carbonate (CaCO3) in water. It is measured in mg L-1 (milligrams per litre) or as ppm (parts-per-million) CaCO3 (even though all the hardness may not be due to CaCO3). Various international agencies have produced guidelines for water hardness classification. For example, the United States Environmental Protection Agency (EPA) classification for total hardness is as follows: Equivalent CaCO3 content (ppm or mg L-1) Description 0-17.1 Soft 17.2-59.9 Slightly hard 60-128.2 Moderately hard 128.3-179.6 Hard

i Historically, water hardness was understood to be a measure of the capacity of the water for precipitating soap. Soap precipitaion occurs in the presence of polyvalent metals such as calcuium, magnesium, aluminum, iron, manganese, strontium and zinc, and by hydrogen ions.

2

179.7+ Very Hard (b) Complexometric titrations: The quantitative analysis of water hardness Cations (acting as Lewis acids) can form complexes with substances that have a pair of unshared electrons (known as ligands, i.e. molecules with the atoms such as N, O and S that have a lone pair and can act as Lewis bases) provided that the Lewis base can satisfy the coordination number of the Lewis acid cations. The number of ligand molecules that can attach to a single cation will depend on (i) the coordination number of the metal ion, and (ii) the number of complexing groups on the ligand molecules. An organic molecule that has two or more groups capable of complexing with a metal ion is known as a chelating agent and the resulting complex is known as a chelate. The formation of metal complexes can form the basis of quantitative analysis through titrations (complexometric titrations) for a quantitative determination of the amount of metal present in a sample. Furthermore, these methods and are particularly useful when the concentrations of these metal ions are present at the millimole levelii. A chelate that is very commonly used in complexometric titrations is ethylenediaminetetraacetic acid, EDTA the structural formula of which is as follows: EDTA is a tetraprotic acid and commonly represented as H4Y where the four hydrogens represent the four ionisable hydrogens on the four carboxyl groups. Since H4Y has a very low solubility in water, the disodium salt Na2H2Y.2H2O is more commonly used. The pH used for the reaction must also be carefully chosen and controlled (see standard analytical chemistry textbooks for details). The determination of the hardness due to Ca2+ and Mg2+ can be carried out at pH 10 (controlled using an ammonium chloride / ammonium hydroxide buffer) using the indicator Eriochrome Black T. Eriochrome Black T is itself a chelating agent having three ionisable protons (hence representable by H3In) which can complex with Mg2+ to from the red complex MgIn− .iii When EDTA (in the from of H2Y2-) is added to a sample containing free Ca2+ and Mg2+ and MgIn− , the EDTA will first complex with all the free Ca2+, then with all the free Mg2+, and finally, once all the Ca2+ and Mg2+ free ions are used up, the strong EDTA ligand displaces indicator from the magnesium. This gives rise to uncomplexed indicator

2HIn − which has a very distinctive blue colour: 2 2 2

2(red) (colourless)(colourless) (blue) MgIn H Y MgY HIn H− − − − ++ → + +

Note that this colour change from red to blue (titration end-point) can be made sharper if the indicator is kept as dilute as practically possible.

ii Note that similar titrations can be carried out for the quantitative determination of anions, e.g. the determination of chlorides through argenometric methods. iii Note that the calcium-indicator complex is less stable than the magnesium-indictor complex, and if no Mg2+ is present in solution, then Eirochrome Black T cannot be used to give a sharp end-point. If no Mg2+ is present, a few drops of Mg-EDTA (Na2MgY(aq)) can be added to obtain a detectible end-point.

3

This titration can be carried out on tap water ‘as supplied’ to determine the total hardness of the water. It may also be carried out on water that has been previously boiled for 20-30 minutes (after it has been cooled, filtered and made to volume, see procedure below) to determine the permanent hardness. (Boiling precipitate the carbonates and bicarbonates as CaCO3 and MgCO3.) Furthermore, it the titration is carried out at higher pH (≥12), then the Mg2+ would not contribute to the titre value as at this high pH, Mg2+ precipitates as Mg(OH)2. The end point in such determinations may be carried out using a suitable indicator, e.g. Patton and Reeder’s indicator or calcon. (c) Statistical testing for differences between means

Chemists are often entrusted with the task of assessing whether the chemical composition (or more simply whether the concentration of some particular component) of two different samples are the same or not. Because of experimental error, such an assessment needs to be done using the appropriate statistical tool. In this practical you are required to establish whether the total hardness of the two water samples as supplies are statistically different or not. Let us assume that on the first sample (sample A), you have carried out NA titrations against EDTA to give titre values

AA,1 A,2 A,, ,..., Nx x x , and that on the second sample (sample B) you have carried out NB titrations against EDTA to give titre values

BB,1 B,2 B,, ,..., Nx x x . You may recall that a good estimate of the true mean titre value at 95% confidence limit for sample A (and similarly for sample B) is given byiv:

A AA A

A

t sxN

µ = ±

where Ax is the arithmetic mean of the titre values AA,1 A,2 A,, ,..., Nx x x given by:

AA,1 A,2 A,A

A

... Nx x xx

N+ + +

=

sA is the standard deviation of the titre values AA,1 A,2 A,, ,..., Nx x x given by:

( )

A 2A, A

1A

A 1

N

ii

x xs

N=

−=

−

∑

and tA is a statistical factor (the t-value) which depends on NA and the confidence level desired. This factor may be obtained from standard statistical tables (see appendix 1) where it is usually quoted in terms of ν and α, where ν is the number of degrees of freedom given by ν = NA – 1, and α relates to the confidence limit (The percentage confidence limit is given by (1 – 2α)100%. Thus for example, a 95% confidence limit is equivalent to an α value of 0.025). Fore example if NA = 3 (i.e. ν = 2), then the t-value at a 95% confidence level (i.e. α = 0.025) is 4.303 (see Appendix 1).

iv These equations are valid for NA < 30.

4

The test to check whether the computed average titre values of two samples are statistically different is based on this equation of µ. A pooled standard deviation, sp, for the NA + NB measurements is calculated from the two standard deviations sA and sB for the two samples. This is given by:

2 2

2A A B B

pA B

N s N ssN N

+=+ −

We can that say at 95% confidence limit, there is no significant difference between the hardness of the two samples if:

A BA B

A Bp

N Nx x tsN N

+− <

where t is obtained from statistical tables at α = 0.025 and ν = NA + NB – 2. For example, if NA = NB = 3, the number of degrees of freedom of the spooled data sample is given by ν = 3+3-2=4, i.e. at 95% confidence limit (α = 0.025), the statistical factor t will have a value of 2.776. Thus, according to the experiments performed, the two samples A and B do not have a significant different ammount of hardeness (at 95% confidence limit) if:

A B3 32.7763.3px x s +− <

i.e. A B 2.267 px x s− < and a different amount of hardness (i.e. a difference in the hardness which is higher than what can be expected from random errors at 95% confidence limit) if: A B 2.267 px x s− > Note that a higher confidence level (i.e. a smaller value of α) for the same value of ν will result in a higher value of t. For example, at 99% confidence level (α = 0.005), the hypothesis that samples A and B do not have a significant different ammount of hardeness (at 99% confidence limit) will hold provided that A B 3.759 px x s− < , i.e. a considerable

relaxation from the requirement that A B 2.267 px x s− < at 95% confidence limit. This is because the acceptable range of values where we can be sure that at 99% confidence level

A Bx x− is solely due to random errors will be higher than if we only want to be sure at a 95% confidence level. An alternative approach for this analysis is to compute a value of t (tcalc) for ν = NA + NB – 2 from:

A B

A B

A B

calc

p

x xt

N NsN N

−=

+

and then look where this value lies in the table for ν = NA + NB – 2. Then if for example, tcalc is found to relate to an α value between 0.025 and 0.01, then we can conclude that the experimental evidence suggests that we can be more than 95% and less than 98% confident that the difference between the means is due to other factors apart from random experimental error.

5

Note: This analysis will only hold if the two samples are analysed using exactly the same procedure (i.e. same method, same apparatus, etc.). Procedure: You have been provided with:

1. Two tap water samples, Sample A and Sample B taken from different locations in Malta

2. 0.01000 M EDTA solution (±0.5%). The solution was prepared from analytic reagent quality of the disodium salt Na2H2Y.2H2O, RMM 372.24. The disodium salt first dried at 80oC. This ensures that the composition agrees with the molecular formula of the dehydrate. A standard solution of 0.01000 M was then prepared using deionised water. The concentration EDTA solution prepared in this way can be assumed to be within ±0.5% of the target concentration.

3. Eriochrome Black T indicator in aqueous solution. This was prepared by dissolving 0.3g of solid Eriochrome Black T in 100mL of deionised water.

4. Calcium hardness indicator tablets for hardness of water determinations, supplied by BDH.

5. A pH 10 buffer solution prepared by adding 142 mL of concentrated ammonia to 17.5g ammonium chloride and diluting to 250 mL with deionised water.

6. 4M NaOH(aq) solution prepared using deionised water. Before your practical, you should look up the risks/safety information for the known chemicals used in this practical. The procedure to be followed is as followsv: (a) To determine the total hardness in water in Sample A and Sample B:

1) Pipette 50.0mL of Sample A in a 250mL conical flask; 2) Add 1mL of pH 10 buffer solution. 3) Add 1-2 drops of the Eriochrome Black T indicator; 4) Titrate against the EDTA. 5) Repeat until you obtain 3 concordant titre values. 6) Repeat for Sample B

(b) To determine the permanent water hardness in Sample A: 1) Measure a volume of 250.0 mL of sample A and transfer this to a 500mL

beaker. 2) Boil for 20-30 minutes. This will remove the temporary out of solution

by precipitating it as MgCO3 or CaCO3. Note that inevitably, some of the water will be lost as steam. However, all the mineral content will remain in solution.

3) Let the boiled water cool to room temperature, making sure that there is no contamination.

4) Filter the remaining solution under gravity using a fluted filter paper such that the filtrate is collected a 250.0mL volumetric flask and make up to 250.0mL using deionised water. (Caution: (i) You should NOT wash the

v Beacuse this practical involved at least (1+3)x4=16 titrations, you may choose to work in pairs. Also note that Part (b) involves boiling for 20-30 mins, and then cooling to room temperature (i.e. a lot of waiting). Thus you are advised to start with Part (b).

6

filter paper - otherwise you will re-dissolve some of residue. (ii) Pay special attention to minimise losses during the transfer.)

5) Proceed as in (a), i.e. pipette 50.0mL of solution so produced in a 250mL conical flask, add 1mL of pH 10 buffer solution and 1-2 drops of the Eriochrome Black T indicator and titrate against the EDTA. Repeat the titrations until you obtain 3 concordant titre values.

(c) To determine the calcium content in Sample A: 1) Pipette 50.0mL of Sample A in a 250mL conical flask; 2) Add 1mL of 4M NaOH solution. This will precipitate any Mg2+ present

as Mg(OH)2 3) Add 1 calcium hardness indicator tablet. Stir the tablet until it completely

dissolves. (This may take a a few minutes.) 4) Titrate against the EDTA. 5) Repeat until you obtain 3 concordant titre values.

(d) Use the data you collected from (a) – (c) to: 1) Determine the total, permanent and temporary hardness of Sample A and

the total hardness of Sample B (as the equivalent to the amount of CaCO3 in water, exprssed as ppm);

2) Determine the hardness (in ppm CaCO3 equivalent) that is due to Ca2+ and Mg2+;

3) Determine whether there is any a significant difference in the ammount of hardeness at 95% confidence limit between Sampes A and B.

Note: For further information please refer to:

o The AAA water testing web-page, http://aaawatertesting.com/ o The US water resources web-page, http://water.usgs.gov/ o Standard analyitical chemistry textbooks

7

Appendix I:

αααα 0.10 0.05 0.025 0.01 0.005 0.001 ν = 1 3.078 6.314 12.706 3.1621 63.657 318.310

2 1.886 2.920 4.303 6.965 9.925 22.327 3 1.638 2.353 3.182 4.541 5.841 10.215 4 1.533 2.132 2.776 3.747 4.604 7.173 5 1.476 2.015 2.571 3.365 4.032 5.893 6 1.440 1.943 2.447 3.143 3.707 5.208 7 1.415 1.895 2.365 2.998 3.499 4.785 8 1.397 1.860 2.306 2.896 3.355 4.501 9 1.383 1.833 2.262 2821 3.250 4.297

10 1.372 1.812 2.228 2764 3.169 4.144 11 1.363 1.796 2.201 2.718 3.106 4.025 12 1.356 1.782 2.179 2.681 3.055 3.930 13 1.350 1.771 2.160 2.650 3.012 3.852 14 1.345 1.761 2.145 2.624 2.977 3.787 15 1.341 1.753 2.131 2.602 2.947 3.733 16 1.337 1 .746 2120 2.583 2.921 3.686 17 1.333 1.740 2.110 2.567 2.898 3.646 18 1.330 1.734 2.101 2.552 2.878 3.610 19 1.328 1.729 2.093 2.539 2.861 3.579 20 1.325 1.725 2.086 2.528 2.845 3.552 21 1.323 1.721 2.080 2.518 2.831 3.527 22 1.321 1.717 2.074 2.508 2.819 3.505 23 1.319 1.714 2.069 2.500 2.807 3.485 24 1.318 1.711 2.064 2.492 2.797 3.467 25 1.316 1.708 2.060 2.485 2.787 3.450 26 1.315 1.706 2.056 2479 2.779 3.435 27 1.314 1.703 2.052 2.473 2.771 3.421 28 1.313 1.701 2.048 2.467 2.763 3.408 23 1.311 1.699 2.045 2.462 2.756 3.396 30 1.310 1.697 2.042 2.457 2.750 3.385 40 1.303 1.684 2.021 2423 2.704 3.307 60 1.296 1.671 2.000 2.390 2.660 3.232

120 1.289 1.658 1.980 2.358 2.617 3.160 ∞ 1.282 1.645 1.960 2.326 2.576 3.090

Percentage points of the Student’s t-distribution (Taken from The University of Malta booklet of Mathematical Formulae, Malta University Press, 1987).

1

University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 12: THE DETERMINATION OF THE CALCIUM CONTENT IN LIMESTONE Introduction For this practical you shall be determining the calcium content in limestone. In particular you will be:

i. Dissolving the limestone sample in acid; ii. Transforming the Ca2+ to calcium oxalate and removing it from

solution; iii. Quantifying the amount of calcium oxalate so formed though the

classical oxalate/permanganate titration. (Note: A potassium permanganate solution (~0.05M) will be provided but this needs to be standardised against sodium oxalate.)

The quality of the final results is highly dependent on the success of the preparation of the titratable calcium oxalate. You are thus advised to take special care so as to avoid losses (or contamination) of the sample. Procedure:

(a) Before your practical, you should look up the risks/safety information for the known chemicals used in this practical.

(b) Standardisation of the approximately 0.05M potassium permanganate provided:

i. Prepare an approximately 0.10M sodium oxalate solution by dissolving about 3.375g of sodium oxalate (Na2C2O4, previously dried at 100-110oC) in distilled water and making up to 250.0mL (using a 250mL volumetric flask). Calculate the accurate concentration of the sodium oxalate solution.

ii. Pipette 25mL of the 0.10M sodium oxalate solution to a 250mL conical flask and add 50mL 2M sulphuric acid.

iii. Heat to about 80oC and titrate slowly with permanganate whilst shaking the flask well. If the solution remains pink after adding the first drop or two, wait until it clears before continuing the titration.

iv. Continue the titration slowly until a faint permanent pink tinge is obtained. The temperature should not be less than 60oC at the end-point. If it falls below this during the titration, the solution must be reheated.

2

v. Repeat this process twice more, or until you obtain three concordant results. Calculate the molar concentration of the potassium permanganate.

(c) Converting the calcium carbonate in limestone to calcium oxalate: i. Weigh accurately 0.3-0.4g of the finely-ground limestone sample in a

250mL beaker, add 5mL water and 10mL of concentrated hydrochloric acid and cover the beaker with a watch-glass (Care! Effervescence may be very copious. You advised to carry out this reaction in a fumehood). Warm the solution to complete the decomposition of the calcium carbonate to yield Ca2+ in aqueous solution.

ii. Dilute the mixture to 50mL, heat almost to boiling, and add 100mL of hot 5% ammonium oxalate solution. Add a few drops of methyl orange indicator and precipitate calcium oxalate (with the beaker kept warm) by dropwise addition of ammonia solution (1 volume concentrated ammonia diluted by 1 volume of water) until the pH of the medium is about 4, corresponding to a yellowish pink colour (imparted to the solution by methyl orange indicator). Allow the solution to stand for about 30 minutes.

(d) Quantifying the amount of calcium present through titration of calcium oxalate against potassium permanganate:

i. Filter the calcium oxalate suspension through a glass sinter (No. 4) under suction and wash the oxalate precipitate with eight to ten 5mL portions of ice-cold water.

ii. Remove the glass sinter and transfer it with its contents to a 600mL beaker and add 200mL of hot 4M H2SO4. Leave the sinter in the beaker. (Be extra careful not to loose any calcium oxalate whilst transferring it!).

iii. Heat the solution to ~85oC with gentle stirring and titrate against the standardised (~0.05M) KMnO4. Use a stirring rod to mix the solutions. A faint pink colour that persists for at least 30 seconds is taken as the end-point. (Note: The temperature of the solution should be kept above 60oC during the entire titration and it may take a minute or two for the first portion of titrant to react.)

(e) Repeat the determination (i.e. parts c & d) at least once, or until you obtain two concordant titre values.

(f) Use the results from the titration to calculate the percentage of calcium in limestone.

1

University of Malta Department of Chemistry

CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 13: ARGENTIMETRIC METHODS

Introduction

In an earlier practical, you have carried out complexometric titrations to determine the concentration of Ca2+ and Mg2+ in a tap water sample, hence determining its hardness. In this practical you will be performing another type of complexometric titration, this time with the scope of quantifying the concentration of specific anions through their complexation with silver cations (argentimetric titrations). Argentimetric titrations are precipitation titrations as the one of the products of the titration reaction (the silver salt) is a precipitate.

You will be determining the concentration of thiocyanate ions [CNS-] in solution by a technique known as Volhard's method. In this method a silver nitrate solution (standardized, see below) is titrated against the solution containing the thiocyanate ions using iron(III) ammonium sulphate (or iron(III) nitrate) as an indicator. The addition of thiocyanate solution produces first a precipitate of silver thiocyanate (AgSCN) and when all the SCN- is used up, a slight excess of the silver will react with the Fe3+ to from the complex [FeSCN]2+ which causes a ‘reddish brown coloration’.

Before you may carry out this titration, you will need to standardize the silver nitrate supplied (approximately 0.02M) by titration against sodium chloride (a primary standard if prepared from dried reagent grade NaCl). This standardisation may be carried out using:

1. The Mohr procedure, i.e. in the presence of a small quantity of potassium chromate solution to act as an indicator;

2. Adsorption indicators (e.g. fluorescein, dichloroflurescein). Both of these standardisation reactions need to be carried out in neutral solution and are discussed in detail in Vogel’s Textbook of Quantitative Chemical Analysis, 6th Edition.

Procedure:

(a) Before your practical, you should look up the risks/safety information for the known chemicals used in this practical.

(b) Standardisation of the approximately 0.02M silver nitrate solution provided (Method 1 - Mohr method):

i. Weigh out accurately about 0.3g of oven-dried sodium chloride. Transfer quantitatively to a 250mL volumetric flask and make up to the mark with distilled water.

2

ii. Pipette 25mL of the solution into a conical flask. Add 1mL of 5% potassium chromate solution to act as indicator and titrate with silver nitrate whilst swirling the liquid. The end-point is indicated by the appearance of the first reddish-brown colour in the solution.

iii. Repeat the titration until concordant results are obtained. iv. Determine the indicator blank as follows: Add 1mL of the indicator

solution to 50mL of distilled water. Titrate with a solution of silver nitrate obtained by diluting 5.0mLof the 0.02M silver nitrate to 50mL ina volumetric flask. The end-point is reached when the colour of the blank is the same as that of the solution that has just been titrated. Repeat the blank titration until concordant results are obtained. Correct the mean blank titre to correspond to the value it would have been if the original solution of the silver nitrate had been employed. Subtract this adjusted value from the mean volume of silver nitrate used in the titration.

v. Calculate the molar concentration of silver nitrate. (c) Standardisation of the approximately 0.02M silver nitrate solution

provided (method 2): i. Repeat the above procedure (part b, i-iii) but use fluorescein (10 drops)

as the indicator. (Note: Fluorescein is an adsorption indicator). The end-point is indicated when the precipitated silver chloride suddenly assumes a pink colour. Repeat the titration until concordant results are obtained. An indicator blank need not be determined with this indicator (though you may wish to verify this for yourselves).

ii. Calculate the molar concentration of silver nitrate as obtained by this titration.

iii. Compare the two titre values and comment on the advantages of the two indicators.

(d) Volhard titration:i. Pipette 25mL of the standardised silver nitrate solution into a conical

flask and add 5mL of 6M nitric acid and the ferric indicator solution. (This indicator is a saturated aqueous solution of iron(III) ammonium sulphate to which 2M nitric acid (10 drops) has been added.)

ii. Titrate this mixture with the potassium thiocyanate solution provided. Near the end-point, the silver thiocyanate becomes flocculent and settles out. The end-point is indicated by the appearance of a permanent faint brown colour. (Question: What causes this faint brown colour?) Repeat the titration until concordant results are obtained.

iii. Calculate the molar concentration of the thiocyanate solution.

10mL 100mL

1

University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 14: AN INVESTIGATION OF THE REACTION BETWEEN COPPER(II) SULFATE AND SODIUM HYDROXIDE Introduction The result of a reaction between NaOH(aq) and CuSO4(aq) is a gelatinous blue precipitate of the binary salt 2 4Cu(OH) CuSOa b which forms between Cu(OH)2 and CuSO4 i.e.: ( ) 4 2 4 2 4CuSO 2 NaOH Cu(OH) CuSO Na SOa b a a b a+ + → +

Chem. Eqn. 1

In this practical you are required to determine the ratio b:a in the formula of the binary salt 2 4Cu(OH) CuSOa b by determining experimentally the amount of NaOH that is required to turn a given quantity of copper (II) sulphate (n1 moles = the number of moles CuSO4, in 25.0 mL of 0.1 M CuSO4(aq)) to the binary salt. Thus, if we know that exactly n2 moles of NaOH are required to turn n1 moles of CuSO4 to the binary salt we know that:

1

22a b n

a n+ =

Eqn. 1 i.e.:

1

2

1 2b nna

+ =

Eqn. 2 or:

1

2

2 1b nna

= −

Eqn. 3 The question thus simplifies to determining the ratio n1 / n2. In this experiment, known volumes of 0.1M NaOH (5.0 mL of 0.1 M, 10.0 mL of 0.1 M NaOH, 15.0 mL of NaOH, etc.) are added to 25.0 mL of 0.1 M CuSO4 solution. The added NaOH reacts with CuSO4 to form the blue gelatinous precipitate. However, if not enough NaOH is added to convert all the CuSO4 to the binary salt, some some of the Cu2+ ions (the excess) will remain in solution. These excess Cu2+ ions can be quantified by:

(1) Adding an excess of I- so that two Cu2+ ions liberates one I2 molecule, 2

22Cu 2I 2Cu I+ − ++ → + Chem. Eqn. 2 (2) Titrating the released iodine against thiosulphate:

2

2- 2-2 2 3 4 6I 2S O 2I S O−+ → + Chem. Eqn. 3

Through this procedure, we can determine the amount of NaOH that is required so that the titre value in the iodine vs thiosulfate titration becomes 0, indicating that there was no Cu2+ in excess, i.e. that just enough NaOH (=n2) was added to turn all the CuSO4 (=n1) to the binary salt. If we let x be the volume of NaOH added and y be the equivalent titre value (in mL) for the iodine against thiosulfate titration, then, given the stoichiometries of the reaction, we may assume that x and y are linearly related, i.e.: y = m x + c

Eqn. 4

Given various values of the pair ( ),i ix y , we may obtain an estimate of the gradient m and the y-intercept c graphically or by regression. Whilst the graphical analysis is probably the best way to visualise the quality of the fit of the ( ),i ix y data to the y = m x + c expression, it is only through regression analyses that we can properly determine m m± ∆ and c c± ∆ given n data points ( )1 1,x y , ( )2 2,x y , …, ( ),n nx y . Linear regression is a statistical method which gives the gradient m and the intercept c of the best straight line through a set of n points having coordinates (x, y). This method is most appropriate when the errors in the x-coordinate of the points are negligible or zero as is often the case in practical work. The expressions for m m± ∆ and c c± ∆ are as follows:

( )22

i i i i

i i

n x y x ym

n x x

−=

−∑ ∑ ∑∑ ∑

( )

2

22

i i i i i

i i

x y x x yc

n x x

−=

−∑ ∑ ∑ ∑

∑ ∑

( )( )2

2

2, 2

12

i in

i

y mx cm t

n x xα −

− −∆ =

− −∑

∑

2ix

c mn

∆ = ∆ ∑

Eqn. 5-8 where

2 , 2ntα − is the t-value at ( )1 2 100%α− confidence limit (i.e. 0.025α = for a 95%

confidence limit) at n-2 degrees of freedom where n is the number of data points. Another quantity which is very useful in these calculations in the value of the value of r, the ‘sample correlation coefficient’ or the ‘Pearson product moment correlation coefficient’, or sometimes simply ‘the r-squared value’ given by:

( ) ( )2 22 2

i i ii

i i i i

n x y yxr

n x x n y y

−=

− −

∑ ∑∑∑ ∑ ∑ ∑

Eqn. 9 This value (r) ranges can be interpreted as a ‘measure’ of the correlation between the x and y data - the greater this correlation, the closer r is to 1. Although equations 5-9 seem very complicated, you should note that most scientific calculation are programmed to include packages with can help in linear regression analysis. Regression analysis can also be performed using ‘spreadsheet’ packages such as MS Excel. This is discussed in a separate document available on the CHE1700 website.

3

Returning back to our discussion, assuming that we have now evaluated m m± ∆ and c c± ∆ , we may determine the volume in mL of 0.1M NaOH that is required to turn all the 25.0 mL of 0.1M copper sulphate solution by determining the x-intercept, i.e. the value of x when y = 0, i.e.:

Vol. of NaOH requried c cm c

± ∆= −± ∆

Eqn. 10 This volume can then be used to determine n2, the number of moles of NaOH that are required to convent n1 moles of CuSO4 to the binary salt (Recall that n1 is the number of moles CuSO4, in 25.0 mL of 0.1 M CuSO4(aq).). The values of n1 and n2 may then be substituted into eqn. 3 to obtain the ratio b:a. Procedure:

(a) Before your practical, you should look up the risks/safety information for the known chemicals used in this practical.

(b) Experimental section1: i. Pipette eight 25mL aliquots of the 0.1M copper sulphate into eight

separate beakers. ii. To the first beaker add (using a pipette) 5mL of the 0.1M sodium

hydroxide, to the second beaker 10mL, to the third 15mL, continuing in this way up to 40mL of sodium hydroxide. Thoroughly mix the contents in each beaker.

iii. Remove the precipitate in the first beaker by filtration through a fluted filter paper and collect the filtrate in a conical flask. Wash the precipitate with cold water (20mL) and collect the washings in the flask. This procedure should be repeated for the contents of each beaker.

iv. Estimate the quantity of unused copper(II) ions in each flask by the following procedure: To each flask add potassium iodide (3g) and titrate the liberated iodine with standard 0.1M sodium thiosulphate using starch solution (2mL) as the indicator.

(c) Calculations i. Plot (using MS Excel or a similar package) the titre value (i.e. the

volume thiosulphate required to reduce all the liberated iodine) against the volume of NaOH added. Fit the best straight line and note the values of the gradient, intercept, and r.

ii. Perform a full regression calculation to obtain values of m m± ∆ , c c± ∆ and r. Check that the values for m, c and r are as obtained from (i).

iii. Calculate 1 1n n± ∆ from the concentration of the copper sulphate solution and the volume of copper sulphate solution used.

iv. Use the value of m m± ∆ and c c± ∆ to obtain an estimate of 2 2n n± ∆ . 1 You are advised to fisrt try the procuedure on just two 25mL aliquits rather than eight. To the first add 5mL of NaOH and to the second add 40mL of NaOH. These trial runs will give you an idea of the range of titre values to expect.

4

v. Use the values of 1 1n n± ∆ and 2 2n n± ∆ to calculate a ab b

± ∆ . Hence

determine the likely formula for the binary salt.

1

University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 15: THE PRARTITION COEFFICIENT OF IODINE BETWEEN iso-OCTANE AND WATER Introduction If a solute X is allowed to distribute itself between two immiscible solvents, A and B (usually water (the aqueous phase) and a water-immiscible organic solvent) then the following equilibrium will be established: ( ) ( )X A X B!!"#!! The equilibrium constant for this process is known as the distribution coefficient or the partition coefficient, KD, or BDA and is given by:

( ) ( ) [ ][ ]

BB A

A

Xthe concentration of X in phase B X Xthe concentration of X in phase A XDK D= = =

In this practical you will be determining the partition coefficient of iodine with respect to the iso-octane (solvent B) and water (solvent A) through a graphical technique. Procedure:

(a) Before your practical, you should look up the risks/safety information for the known chemicals used in this practical (i.e.

(b) Experimental section: In this practical you shall be working in pairs to obtain a total of six data points. In particular, referring to the procedure below, you should let:

V1 (mL) V2 (mL) V3 (mL) Student 1 10.0 20.0 30.0 Student 2 15.0 25.0 35.0

i. Pipette a volume V1 of iodine solution in iso-octane (approximately 0.3%

w/v) into your separatory funnel, add 100mL of distilled water and acidify with 1mL of 2M sulphuric acid. Shake vigorously for 10 minutes, and allow the two phases to separate. (Swirl if necessary to detach droplets held up on the inside.)

ii. Drain the lower aqueous layer into a clean, dry beaker and titrate a 50mL aliquot against the 0.008M sodium thiosulphate solution provided. Add the starch indicator only as the end point is approached.

iii. Pipette 5mL of the organic layer into a conical flask. Add 25mL of distilled water and about 2g of solid potassium iodide; (Why is this

2

important?). Swirl the flask, and titrate with the 0.008M thiosulphate solution, as above.

iv. Repeat parts (i-iii) using volumes V2 and V3 of iodine solution in iso-octane.

(c) Calculations i. Calculate and tabulate the concentration of iodine in the aqueous and

organic layers ([I2]aq and [I2]org) for V = 10.0mL, 15.0mL, 20.0mL, …, 35.0mL.

ii. Plot [I2]org against [I2]aq. Fit the best straight line (passing through the origin, i.e. with c = 0) and note the values of the gradient and r.

iii. Perform a full regression calculation to obtain values of m m± ∆ and r and hence estimate a value of KD (at a 95% confidence level).

1

University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 16: AN INVESTIGATION OF THE REACTION BETWEEN IRON(II) IONS AND NITRATE(V) IONS IN ACID SOLUTION Introduction In this practical you shall be investigating the reaction between iron(II) and nitrate(V) ions. FeII (used in excess) will first be allowed to react with a known amount of nitrate(V) anions and the amount of excess iron(II) is then estimated through a redox titration against standard potassium dichromate. The end-point of this reaction is determined by the use of the redox indicator sodium diphenylamine sulphonate. Procedure: For this practical you are going to be supplied with1:

o Solid iron(II) ammonium sulphate, (NH4)2Fe(SO4)2.6H2O o Standard potassium dichromate solution (0.017M) o Standard potassium nitrate solution (2.0gL-1) o Sodium carbonate solution (10%) o Solid sodium bicarbonate o A redox indicator: aqueous sodium diphenylamine sulphonate (0.3%) o Concentrated sulphuric acid o Concentrated phosphoric acid

(a) Before your practical, you should look up the risks/safety information for the known chemicals used in this practical.

(b) Preparation of iron(II) ammonium sulphate standard solution: Prepare a standard solution of iron(II) ammonium sulphate of approximate concentration 0.1M by dissolving an accurately weighed amount of the salt in 250mL of solution.

(c) Investigation2: i. Pipette 25mL of the iron(II) solution into a conical flask, add 5mL

concentrated sulphuric acid and rapidly add about 1g sodium bicarbonate so that the carbon dioxide displaces the air from the flask.

ii. Add an additional 20mL of concentrated sulphuric acid and pipette 25mL of the potassium nitrate solution into the solution. Swirl the mixture and boil for 5 minutes.

1 The actual concentrations of the standard solutions will be supplied by your lab officers. 2 This procedure (apart from the titrations) should be carried out in a fume-hood.

2

iii. Place a dropping funnel containing 50mL of 10% sodium carbonate solution pointing into the conical flask. As the boiled solution cools, add small aliquots of the sodium carbonate solution so that the carbon dioxide released prevents any air entering the flask.

iv. When the solution has cooled, add 5mL of concentrated phosphoric acid, 8 drops of the redox indicator, and titrate the excess iron(II) with the standard dichromate until an intense purple coloration is obtained at the end-point. (Note: The titration should not be delayed after all the carbonate has been added.)

v. Repeat parts (i-iv) at least twice or until you obtain concordant results. (d) Calculations and questions: Calculate the mole ratio of iron(II) ions to nitrate

i. Calculate the mole ratio of iron(II) ions to nitrate ions for this reaction and construct an equation for the reaction of iron(II) ions with nitrate ions in acid solution.

ii. Comment of the role of the air expelling agents in this analysis, and on the necessity that the titration is performed soon after the addition of the bicarbonate.

1

University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 17: PREPARATION OF 1-BROMOBUTANE Introduction In this practical you shall be converting butan-1-ol (b.pt. 117.7oC) to 1-bromobutane (p.bt.: 100-104oC) by heating with NaBr under reflux in the presence of concentrated sulphuric acid:

2 4

NaBr3 2 2 2 3 2 2 2H SOCH CH CH CH OH CH CH CH CH Br→

This is followed by various purification steps in an attempt to isolate pure 1-bromobutane. Procedure:

(a) Before your practical, you should look up the risks/safety information for the known chemicals used in this practical.

(b) Preparation of bromobutane (p.bt.: 100-104oC): Place sodium bromide (27g), water (30mL), and n-butanol (20mL, 16.2g) in a 250mL round-bottomed flask. Cool the mixture in an ice-bath and slowly add concentrated H2SO4 (23mL, 42.3g) with swirling and cooling. Heat under reflux for 30 minutes.

(c) Extraction and purification of bromobutane from the reaction mixture: i. Distil the reaction mixture until no more water-insoluble droplets come

over, by which time the temperature should have reached 115oC. (Why does the boiling point steadily increase?)

ii. Pour the distillate into a separating funnel and shake with 20mL water. Drain the lower layer of butyl bromide into a clean flask.

iii. Wash and dry the separating funnel and return the butyl bromide to it. iv. Cool concentrated sulphuric acid (20mL) in an ice-bath, add to the

separating funnel, shake well and allow 5 minutes for settling of the layers.

v. Separate the layers, allow 5 minutes for further drainage, and separate again. (Note: The densities of 1-bromobutane and concentrated sulphuric acid are respectively 1.27 and 1.84g/mL.)

vi. Wash the butyl bromide with 10% sodium hydroxide (20mL), separate and be careful to save the proper layer.

vii. Dry the cloudy butyl bromide by adding calcium chloride (5g) and warm the mixture gently on a steam-bath with swirling until the liquid clears.

2

viii. Decant the dried liquid through a funnel fitted with a small loose plug of cotton into a small flask, add boiling chips, distil, and collect the liquid in the range of 99-104oC.

ix. Calculate the yield and express it as a percentage of the theoretical yield. x. Place the sample in a suitable container and label, giving the name and

formula of the product and your name. (d) Questions:

i. How could you distinguish between a sample of your product and that of its isomer 2-bromo-2-methylpropane by simple test-tube reaction? Do this test. A sample of 2-bromo-2-methylpropane is provided.

ii. Write down the mechanism for the reaction of butan-1-ol with NaBr in the presence of H2SO4 to form 1-bromobutane.

iii. What organic impurities are likely to be present in the crude product? Why are these impurities likely to be present? How does each of these impurities react with sulphuric acid when the bromobutane is shaken with this reagent?

1

University of Malta Department of Chemistry CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 18: PREPARATION OF BUTYL ACETATE Introduction

The interaction between a carboxylic acid and an alcohol is a reversible process and proceeds sluggishly towards an equilibrium mixture. The equilibrium is usually reached after refluxing the mixture for several days. However, one may:

1. Increase the reaction rate by adding concentrated sulphuric acid (about 3% of the weight of the alcohol);

2. Increase the yield of ester by using an excess of either the acid or the alcohol. This method of direct esterification works well for primary and secondary alcohols but is unsatisfactory for tertiary alcohols. In particular, it can be used to prepare butyl acetate from ethanoic acid and butan-1-ol:

( ) ( )2 4conc. H SO3 3 2 2 3 2 2 3reflux, 1 hr2 2

CH COOH + CH CH CH OH CH COOCH CH CH → Procedure:

(a) Before your practical, you should look up the risks/safety information for the known chemicals used in this practical.

(b) Preparation of bultyl acetate: Mix together 18.5g (23mL, 0.25mol) of butan-1-ol and 30g (30mL, 0.5mol) of glacial acetic acid in a 250mL round-bottomed flask. Add cautiously 1mL of concentrated sulphuric acid (use a small measuring cylinder). Attach a reflux condenser and reflux the mixture for 1 hour.

(c) Extraction and purification of butyl acetate from the reaction mixture: i. Pour the reaction mixture into about 50mL water contained in a

separating funnel. Remove the upper layer of crude ester and wash it again with about 50mL of water, followed by about 15mL of saturated sodium hydrogencarbonate solution and 25mL of water.

ii. Dry the crude ester with 6g anhydrous sodium sulphate, filter through a small funnel containing a fluted filter-paper.

iii. Distil the dried product (on a wire gauze) and collect the pure butyl acetate at 124-126oC.

iv. Record the yield and express it as a percentage of the theoretical yield. v. Place the sample in a suitable container and label, giving the name and

formula of the product and your name.

2

(d) Questions: i. Give the mechanism of the reaction involved in this preparation.

ii. Why is this method of direct esterification not suitable for esters derived from tertiary alcohols? How are esters derived from tertiary alcohols usually prepared?

iii. A recent procedure for the preparation of methyl esters involves refluxing the carboxylic acid with methanol and 2,2-dimethoxy propane in the presence of toluene-p-sulphonic acid. Explain the chemistry of this method.

1

University of Malta Department of Chemistry CHE1700: Chemistry Practical I - B.Sc.(Hons.) I Year

EXPERIMENT 19: ISOLATION OF LYCOPENE FROM TOMATO PASTE Introduction Lycopene, the red pigment of the tomato, is a coloured liphatic hydrocarbon and belongs to a group of compounds known as C40 carotenoids. Lycopene with its antioxidant properties has generated widespread interest as a possible deterrent to heart disease and some forms of cancers (e.g. prostate cancer). Fresh tomato juice contains about 96% water and 1kg of fruit yields only 0.02g of lycopene. A more convenient source of lycopene is commercial tomato paste which is 26% solid and 0.15g lycopene can be isolated per kg of paste.

CH3 CH3 CH3

CH3CH3 CH3 CH3 CH3

CH3CH3

Lycopene (C40H56)

In this practical you shall be extracting lycopene through chromatography from commercial locally produced tomato paste. The isolation procedure affords amounts of lycopene that are unweighable except on a microbalance, but more than adequate for an illustration of the technique of chromatography. (The expected yield in the present experiment is 0.00075g.) The term chromatography generally refers to the separation of different components of a sample mixture between two phases: a stationary phase (usually a solid or a liquid adsorbed on a solid) and a mobile phase (liquid or gas). Components of the sample mixture are carried through the stationary phase by the flow of the mobile phase and the separation occurs due to differences in the migration rates among the different sample components. In column chromatography, the stationary phase is held in a tube and the mobile phase is forced through the tube under pressure of by gravity. CAUTION: Lycopene is very sensitive to photochemical air oxidation. Protect solutions

and solid from undue exposure to light.

2

Procedure:

(a) Before your practical, you should look up the risks/safety information for the known chemicals used in this practical.

(b) Solvent extraction of pigments from tomato paste: i. Weigh 5g of tomato paste in a 100mL round-bottomed flask. Add 10mL

ethanol and heat under reflux for 5 minutes (using a heating mantle). ii. Filter the hot mixture on a small Hirsch funnel. Let the flask drain

thoroughly. Squeeze liquid out of the solid residue with a spatula. iii. Return the solid residue to the flask, add 10mL dichloromethane, and

reflux for 3-4 minutes. Filter the yellow extract and combine it with the ethanol filtrate.

iv. Repeat the extraction with 2 or 3 further 10mL portions of dichloromethane.

v. Pour the combined extracts into a separating funnel, add water and about 15mL sodium chloride solution (to aid layer separation) and shake.

vi. Run the coloured lower layer through a cone of sodium sulphate into a dry flask and evaporate to dryness IN A FUMEHOOD (avoid overheating). This cone should be prepared using a minimal amount of sodium sulphate (to avoid loss of product) supported by a little glass wool in a funnel.

(c) Preparation of a 12cm column of alumina using light petroleum (40-60oC) as solvent:

i. CAUTION: You have been provided with a glass chromatography column. Take special care to ensure that:

? The chromatography column and other glassware used in the preparation of the column should not be wetted with water. If the column/glassware need cleaning, they should be cleaned with the solvent (in this case pet. ether 40-60oC).

? The stopcock should not be tightened too much as this could result in jamming and/or breakages.

? The properties (activity) of the alumina are dependent on its water content and different forms or alumina with different levels water content can be used in column chromatography. The water content (and hence the activity) is described by the Brackmann Scale, where ‘Alumina Brackmann Grade I’ is the driest and most active form and ‘Alumina Brackmann Grade V’ is the wettest and least active form (least polar). For this practical, you should make use of Neutral Alumina Brackmann Grade IV which can be prepared from the Neutral Alumina Brackmann Grade I by adding 10% water and mixing well.

ii. Mix the stationary phase (activated alumina, an adsorbent) and solvent (pet. ether 40-60°C) in a beaker (stir if necessary) and pour the resulting slurry into the coulumn. The adsorbent settles fairly rapidly under gravity and the process may be assisted by gently tapping the column (avoid air bubbles).

iii. Once all the stationary phase is introduced, run out excess solvent from the column but make sure that the upper surface of the adsorbent remains

3

covered with liquid at all times. Otherwise the column may shrink and lead to oxidation of the adsorbate.

(d) Extraction of lycopene through column chromatography*: i. Dissolve the crude product obtained from (b-vi) in 2-3mL of toluene and

transfer the solution onto the column using a Pasteur pipette. ii. Elute the column with light petroleum (b.pt. 40-60oC) (using a separating

funnel). The rate of solvent flow should be 4-10cm/min for a 2cm (diameter) column or 20-40cm/min for a 5cm column.

iii. You will note that bands will start forming along the column. You must discard the initial colourless eluate and collect the main orange fraction (a further narrower red band is due to ß-carotene, another constituent of tomato paste).

iv. Evaporate the solution at room temperature to obtain pure lycopene. Examination of the material may reveal crystallinity.

(e) Examination of the extracted lycopene: Treat the product with a drop of conc. H2SO4 to give a characteristic colour.

(*) Note: It is sometimes recommended to place a loose plug of cotton wool at the top of the column to protect the surface of the adsorbent from disturbance when the mobile phase is introduced.

ALL WASTE PET. ETHER SHOULD BE DECANTED INTO THE NON-HALOGENATED WASTE BOTTLE.

1

University of Malta Department of Chemistry

CHE1700 - Chemistry Practical I - B.Sc. (Hons.) I Year

EXPERIMENT 20: THIN LAYER CHROMATOGRAPHY OF 2,4 DINITROPHENYLHYDRAZONES

Introduction

In this practical you shall be identiying aldehydes and ketones present in a mixture by converting them to 2,4-dinitrophenylhydrazones and then identifying these though thin-Layer Chromatography (TLC).

TLC is a is a chromatographic technique that is useful for separating organic compounds and because of the simplicity and rapidity, it is often used to monitor reactions, check the purity of products or confirm their identity.

In TLC, the solid stationary phase is coated onto glass or plastic plates (the chromaplate). The most commonly used stationary phases are silica gel, alumina, kieselguhr and cellu-lose powder but could also include polyamides, modified celluloses with ion exchange properties or various forms of organic gel having molecular sieving properties (e.g. Sephadex, Bio-Gel P). Very often, a fluorescent compound (e.g. zinc sulphide) is also incorporated with the stationary phase in order to facilitate the detection of the resolved components of the mixture which is then achieved by viewing the plates under ultraviolet light. In this practical you shall be with 20 x 20cm thin-layer plates which have been coated with a 500 micron layer of silica gel. These plates are very fragile and should be handled with caution, the layer itself must not be touched.

The mobile phase is an organic solvent or a mixture of organic solvents and is placed inside a development jar (~ 1cm). Small 20 x 5cm plates are conveniently developed in a cylindrical glass jar as in Fig. 1a whilst the larger plates require a rectangular glass tank as in Fig. 1b. The atmosphere inside of the jar must be saturated with the solvent before the chromaplate is inserted. This can be done by wetting the sides of the jar with the solvent, lining the jar with a filter paper saturated in the developing solvent (leaving a gap for viewing the chromaplate) and closing the jar for about 10 minutes to allow the atmosphere inside the jar to become saturated with solvent vapour.

2

Fig. 1: The cylindrical and rectangular glass jars / tanks used for developing TLC plates.

The samples (liquids or solids dissolved in a volatile solvent) are deposited as spots on the stationary phase. The spots should be placed around 2cm from the bottom of the plate, and spots of different samples should be separated from each other by about 1.5-2.0cm. It is very important that all spots are placed on a line (the start line) that is parallel to the base of chromaplate. To ensure this, one may (i) draw a line (using a pencil, without damaging the stationary phase) parallel to the base and 2cm from it, and (ii) mark on this line equidistant points where the samples should be placed. It is also useful to draw a finish line, usually about 12cm from and parallel to the start line.

The samples have to be individually applied to the marked points on the adsorbent layer by means of a sample applicator. This is prepared by drawing out a melting point capillary tube in a Bunsen flame and snapping the drawn-out portion in two. The applicator is charged by dipping the capillary end into the solution and after withdrawing, touching the end on a piece of filter paper until the volume is reduced to about 0.5µL. The solution may then be transferred to the plate by touching the tip of the capillary on to the adsorbent layer, taking care not to disturb the surface unduly. The size of the spot should be about 3mm in diameter.

Once the plate is prepares, and the developing jar/tank is saturated with the solvent vapour, one may insert the plate into the developing jar (with the spots towards the bottom of the jar) tilted as shown in Fig. 1 so that the uncoated face is uppermost. If need be, you may carefully pour down the side of the jar more developing solvent so that the bottom of the adsorbent layer is well immersed. It is important to ensure that the solvent level should does not reach as far as the spots. The jar should then be recapped and allowed to stand so that the solvent ascends (by capillary action) to the finishing line.

It is very important that a solvent with the right properties is used as the mobile phase. If the chromatographic behaviour of the substances under investigation is unknown, one should first carry out trail runs to identify teh best solvent system. This can be done by setting up a series of develping jars containing solvent systems of increasing polarity (e.g hexane, toluene, carbon tetrachloride, dichloromethane, diethyl ether, ethyl acetate, acetone, methanol). Identically loaded micro-plates are developed separately using the chosen solvents, dried and examined. Solvents which cause all the components to remain

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near to the spot origin or to move near to the solvent front are clearly unsatisfactory. If it is seen that no single solvent gives a satisfactory chromatogram, with well-spaced compact spots, it is necessary to examine the effect of using mixtures of solvents to provide systems having a range of intermediate polarity. For example, mixtures of toluene and methanol, or hexane and ethyl acetate, arc often suitable when the pure solvents are unsatisfactory.

The time required to complete this development varies greatly with the composition of the solvent and the nature of the adsorbent. If the system is inconveniently slow-running the development process may be terminated before the solvent reaches the finishing line, provided that the position of the solvent front is marked on the adsorbent layer immediately the plate is removed from the development tank. After removal, the plate is dried and the suitably depending upon the volatility and toxicity of the solvent system; for example, dry the plate in the fume cupboard (if necessary) with a warm-air blower or dry in a temperature-controlled oven, etc. The plate may then be examined to determine the positions of components, possibly with the help of a UV lamp, or through other means (e.g. by spraying with an appropriate reagent).

The different components in the mixture move up the plate at different rates due to differences in their partioning behaviour between the mobile liquid phase and the stationary phase. Provided that the experimental conditions are reproducible the movement of any substance relative to the solvent front in a given chromatographic system is constant and characteristic of the substance. The constant is the Rf value and is defined as:

distance moved by substance distance moved by solvent front

=fR

True reproducibility in Rf values is however rarely achieved in practice due to minor changes in a number of variables such as:

(a) the particle size of different batches of adsorbent; (b) the solvent composition and the degree of saturation of the tank atmosphere with

solvent vapour; (c) prior activation and storage conditions of the plates; (d) the thickness of adsorbent layer, etc.

This means that Rf values are not used as a criterion for identification of substances. Instead, reference compounds are usually run on the same plate as the unknown to substantiate the identification.

Procedure:

(a) Before your practical, you should look up the risks/safety information for the known chemicals used in this practical.

(b) Preparation of the 2,4-dinitrophenylhydrazones:i. Preparation of the DNP reagent: Prepare the DNP reagent by taking 4g

of the DNP solid provided, suspending this in 80mL methanol and adding dropwise 10mL concentrated sulphuric acid (CAUTION). This

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should give you a clear warm solution. If undissolved solid remains, filter off.

ii. Preparation of one of the standard DNP derivatives: You shall be provided with one of the following aldehydes/ketones: cyclohexanone, benzaldehyde, salicaldehyde, acetone, butanal which should be converted to its respective DNP derivative.

a. Take 2 mL of this aldehyde/ketone and dissolve it in 2 mL methanol.

b. To this add 40mL of warm DNP reagent. c. Keep the mixture warm for 10min on a water bath, cool in ice-

water, and collect the yellow or orange precipitate indicating the presence of a carbonyl function. (If a precipitate does not form immediately, a little water is added to the mixture and is set aside for 5 minutes.)

d. When the precipitation is complete, the solid is removed by filtering through a Hirsh funnel under suction.

e. If you have time recrystallise the solid derivative from analcohol-water mixture (1:1) and dry in an oven.

f. Record the m.pt. of the DNP derivative you prepared and compare it to the literature value.

g. Dissolve some of the standard 2,4-dinitrophenylhydrazone you prepared in into a small amount of ethanol, label it, and make it available to the rest of the class for use as one of the standard DNP derivatives (see Table 1).

iii. Preparation of the DNP derivatives of the aldehydes/ketones mixture: You shall be provided a mixture of aldehydes and ketones which should also be converted to their respective DNP derivatives.

a. Add 25mL of warm DNP reagent to a mixture of 1mL alcoholicaldehyde-ketone mixture provided and 1mL warm methanol.

b. Keep the mixture warm for 10min on a water bath, cool in ice-water, and collect the mixed crystalline DNP derivatives on a Hirsch funnel. Suck the crystals dry.

(c) Preparation of the developing tank:i. Cut a piece of filter paper to fit inside the tank and use it to cover one of

the larger vertical surfaces. ii. Place in the tank 50mL of a 2:1 mixture of 80-120oC Pet. Ether : Ethyl

Acetate. Wet the inside of the tank and the filter paper with the solvent (using a pipette).

iii. Cover the tank and leave it to stand for 10 minutes so that the air in the tank becomes saturated with the solvent vapour.

(d) Preparation of the plates: TLC plates are very fragile and should be handled with caution, the layer itself must not be touched.

i. Take a 20x20 cm plates and mark (score very gently with a sharp pencil and ruler) a line (the starting line) parallel to one edge and 2cm from it. (NOTE: These plates are very fragile and should be handled with caution, the layer itself must not be touched.)

ii. On this line, mark the position of eight starting points, equidistant from each other.

iii. Score gently a finish line 12cm from and parallel to the starting line.

Take 1mL of the aldehyde-ketone mixture provided, dissolve itin 1mL warm methanol, and add 25mL of warm DNP reagent.

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iv. Dissolve a small portion of the 2,4-dinitrophenylhydrazones you prepared in Part b into a small amount of ethanol. Use this solution, as well as the original DNP reagent and the five standard DNP derivatives prepared by yourselves, for spotting the plate as detailed in Table 1 below1.

(e) Development of the plates: i. Very carefully, place the plate in the devolving jar as illustrated in Fig. 1

above.ii. Run the plate until the solvent reaches the finish line. This should take

about 30 minutes.(f) Examination of the developed TLC plates:

Record in diagram form all the spots visible in daylight and also when inspected under a UV lamp. Account for as many of the other spots as you can. (Note: The stationary phase has an inorganic phosphor incorporated in it and some UV absorbers will be visible as purple spots. The DNP's will be seen as yellow spots and the components of your mixture should be identified by reference to the Rf values of the known derivatives.)

Origin No. No. of Spots Nature of Solution Applied 1 2 Mixed DNP derivatives 2 1 DNP reagent solution 3 1 Cyclohexanone DNP 4 1 Benzaldehyde DNP 5 1 Salicaldehyde DNP 6 1 Acetone DNP 7 1 Butanal DNP 8 1 Mixed DNP derivatives

Table 1

1 You should practise spotting of plates on the small plates provided. Your demonstrator will show you how to do it. When you are confident about the technique, apply the eight spots on the proper chromatography plate.