NATIONAL QUALIFICATIONS CURRICULUM SUPPORT

Chemistry

A Practical Guide

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[REVISED ADVANCED HIGHER]

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Contents

Introduction 5

Chemical analysis 6
Qualitative and quantitative analysis 6
Volumetric analysis 6
Gravimetric analysis 14
Colorimetric analysis 17

Organic techniques 22
Introduction 22
Preparation 22
Isolation 24
Purification 29
Identification 33
Percentage yield 37

Errors 39
Accuracy and precision 39
Repeatability and reproducibility 41
Quantifying errors 41
Absolute uncertainties and percentage uncertainties 42
Combining uncertainties 43
Some ‘forgotten’ uncertainties 46

Experiments 53
Experiment 1A: Preparation of a standard solution of 0.1 mol l–1 oxalic acid 53
Experiment 1B: Standardisation of approximately 0.1 mol l–1 sodium hydroxide 55
Experiment 1C: Determination of the ethanoic acid content of white vinegar 57
Experiment 2A: Preparation of a standard solution of 0.1 mol l–1 sodium carbonate 59
Experiment 2B: Standardisation of approximately 1 mol l–1 hydrochloric acid 61
Experiment 2C: Determination of the purity of marble by back titration 63
Experiment 3: Determination of nickel in a nickel(II) salt using EDTA 65
Experiment 4A: Gravimetric determination of water in hydrated barium chloride 67
Experiment 4B: Gravimetric determination of nickel using dimethylglyoxime 69
Experiment 5: Preparation of potassium trioxalatoferrate(III) 72
Experiment 6: Determination of vitamin C 75
Experiment 7A: Preparation of aspirin 77
Experiment 7B: Determination of aspirin 80
Experiment 8: Preparation of benzoic acid by hydrolysis of ethyl benzoate 83
Experiment 9: Preparation of ethyl ethanoate 86
Experiment 10: Colorimetric determination of manganese in steel 88
Experiment 11: Preparation of cyclohexene from cyclohexanol 91

Introduction

This material has been written to support students with the practical work of Advanced Higher Chemistry and in particular the Researching Chemistry unit, which includes the investigation.

It is divided into four main sections:

• chemical analysis
• organic techniques
• errors
• experiments

The main aims of the sections on chemical analysis and organic techniques are to introduce a wide variety of techniques, to provide a sound understanding of the underlying chemical principles and to develop laboratory skills.

The section on errors deals with accuracy and precision, repeatability and reproducibility but its main thrust is to show how the overall uncertainty in the final result of an experiment can be quantified in terms of the uncertainties in the individual measurements made in the experiment.

The final section details a number of possible experiments that cover all the skills and techniques required of the Researching Chemistry unit. For each experiment there is a brief introduction, a list of requirements in terms of equipment and chemicals, the hazards and control measures associated with the chemicals used and a detailed experimental procedure.

Chemical analysis

Qualitative and quantitative analysis

There are two types of chemical analysis: qualitative and quantitative. Qualitative analysis is the process of identifying what is in a chemical sample whereas quantitative analysis is the process of measuring how much is in the sample. In this section we are concerned with methods of quantitative analysis.

Volumetric analysis

Volumetric analysis relies on methods involving the accurate measurement of volumes of solutions, although mass measurements may also be required. Essentially, we measure the volume of a standard solution (one of accurately known concentration) needed to react exactly with a known volume of another solution (one of unknown concentration) in a chemical reaction for which the stoichiometric or balanced chemical equation is known. From the data, we are then in a position to calculate the accurate concentration of the second solution.

In practical terms, volumetric analysis is achieved by a titration procedure. In a titration, one of the solutions is added from a burette to a pipetted volume of the other solution in a conical flask. The point at which the reaction between the two is just complete is usually detected by adding a suitable indicator to the solution in the flask. It is customary, although not essential, to have the solution of known concentration in the burette.

There are numerous types of titration but the most common are:

• acid-base titrations, which are based on neutralisation reactions
• redox titrations, which are based on oxidation–reduction reactions
• complexometric titrations, which are based on complex-formation reactions.

The principal requirements of a titration reaction are that it goes to completion and proceeds rapidly.
Standard solutions
As mentioned above, a standard solution is one of accurately known concentration and it can be prepared directly from a solute if that solute is a primary standard. To be suitable as a primary standard, a substance must meet a number of requirements.

• It must have a high purity. This is to ensure that the mass of the sample weighed out is composed entirely of the substance itself and nothing else. Were impurities present, then the true mass of the substance present would be less than the measured mass and this would lead to the solution having a concentration less than the calculated value.
• It must be stable in air and in solution. If this were not the case then some of the substance would be used up in reacting with chemicals in the air or with the solvent. As a result, the true concentration of the resulting solution would be less than its calculated value.
• It must be readily soluble in a solvent (normally water) and its solubility should be reasonably high so that solutions of relatively high concentrations can be prepared.
• It should have a reasonably large relative formula mass in order to minimise the uncertainty in the mass of substance weighed out.

As a result of these exacting criteria, there are a limited number of primary standards available. Some examples of acids, bases, oxidising, reducing and complexing agents used as primary standards are outlined in the following table.

|Primary standard |Examples |
|Acid |Hydrated oxalic acid, (COOH)2.2H2O |
| |potassium hydrogenphthalate: |
| | |
| | |
|Base |Anhydrous sodium carbonate, Na2CO3 |
|Oxidising agent |Potassium dichromate, K2Cr2O7; potassium iodate, KIO3 |
|Reducing agent |Sodium oxalate, (COONa)2 |
|Complexing agent |Hydrated disodium salt of EDTA: |
| |[pic] |

Chemicals are supplied in various grades of purity but for analytical work AnalaR grade primary standards must be used. AnalaR grade guarantees high purity.

You will notice that sodium hydroxide, although commonly used in quantitative analysis, is not included in the table as a primary standard. This is because it absorbs moisture from the air and dissolves in it to form a very concentrated solution. Furthermore, both solid sodium hydroxide and a solution of it react with carbon dioxide from the air. Consequently, it is unstable in air and so does not meet the exacting requirements of a primary standard.

The procedure involved in preparing a standard solution directly from a primary standard is detailed below.

You must first calculate the mass of the primary standard required given the volume and concentration of solution you desire. The sample of the primary standard must be dried in order to remove any traces of water that may have been adsorbed from the atmosphere. This is particularly important when using older samples of the substance. The water impurity can be removed by placing some of the substance in a crystallising basin and storing it in a desiccator for several hours.

A desiccator is a closed vessel that contains a desiccant (a drying agent) in its base. Desiccants include phosphorus pentoxide, anhydrous calcium chloride and concentrated sulfuric acid, but the one that is most commonly used is self-indicating silica gel: it is blue when dry and turns pink when it absorbs moisture. An airtight seal is maintained in the desiccator by lightly greasing the ground-glass surfaces on the lid and base.

Alternatively, primary standards can be dried by heating, although this runs the risk of them decomposing if too high a temperature is used.

Once the primary standard is dry, the next step in the procedure is to weigh out accurately the approximate mass of substance you need to make the desired solution. The words ‘accurately’ and ‘approximate’ may sound ambiguous but what it means is that while the mass of the sample of primary standard has to be known accurately, it doesn’t need to be exactly that calculated – just close to it.

It is good practice to use a weighing bottle when weighing out samples of primary standards. There are various types and the one illustrated is a cylindrical glass container fitted with a ground-glass stopper.

The weighing technique described below is known as ‘weighing by difference’.

A clean dry weighing bottle is first weighed empty and then, using a spatula, a sample of the primary standard of mass close to the calculated value is added to it. The accurate mass of the weighing bottle and its contents is then measured and recorded. The next step is to transfer the sample of the primary standard from the weighing bottle to a clean glass beaker containing some deionised water. Gentle tapping on the base of the weighing bottle will ensure that the bulk of the sample is transferred but it is unimportant if traces of the sample remain. Finally, the weighing bottle and any residual material are accurately weighed and the mass recorded. The accurate mass of the primary standard transferred is the difference between the two recorded masses.

Throughout the weighing process it is important that the stopper be removed from the weighing bottle only when necessary. This reduces the time the sample is exposed to the atmosphere and so minimises the chances of it re-adsorbing moisture.

A balance reading to 0.01 g should be adequate in weighing out samples of primary standards but if greater accuracy is required then a balance reading to three decimal places should be used.

With the sample of the primary standard successfully transferred to the beaker of deionised water, the mixture can be stirred to aid dissolving. A glass rod should be used for this purpose and not a spatula since the latter may react with the solution and so contaminate it. On removing the stirring rod, make sure that any solution on its surface is washed back into the beaker. A wash bottle can be used to achieve this.

Once the primary standard has dissolved, the resulting solution is carefully poured into an appropriately sized standard (volumetric) flask via a filter funnel placed in the neck of the flask. Both the flask and the funnel must be clean but neither need be dry just so long as they are wet with deionised water. Using a wash bottle, the interior surface of the beaker should be washed with deionised water and the washings transferred to the flask. The washing process should be repeated at least two more times to ensure that all the primary standard has been transferred to the flask. Deionised water is then added directly to the flask until the level of the solution is within about 1 cm of the graduation mark. With the funnel removed, deionised water is carefully added from a dropper until the bottom of the meniscus is level with the graduation mark. During this last operation, a white tile or a piece of white paper should be held behind the neck of the flask so that the meniscus can be seen more clearly. The graduation mark must be at eye level in order to avoid error due to parallax.

The standard flask should then be stoppered and inverted several times to ensure the solution is thoroughly mixed and is of uniform concentration. The solution of the primary standard should finally be transferred to a clean, dry reagent bottle. If the reagent bottle happens to be wet with deionised water, then it must first be rinsed with a little of the standard solution before the bulk of the solution is transferred to it. Were it not rinsed, then the solution would be diluted by the water, making its true concentration slightly less than its calculated value.

Titrations
Once a standard solution has been prepared, it can be used to determine the accurate concentration of another solution. This is achieved by titration – a procedure whereby one of the solutions is slowly added from a burette to a pipetted volume of the other solution contained in a conical flask. The point at which reaction between the two is just complete is usually detected by adding a suitable indicator to the solution in the flask. It is customary, although not essential, to have the standard solution in the burette and the solution of unknown concentration, often referred to as the analyte, in the conical flask. The practical aspects of a titration are detailed below.

A clean burette has first to be rinsed with a small portion of the standard solution. This involves tilting the burette almost to a horizontal position and rotating it to make sure the standard solution ‘wets’ the entire inner surface. The burette tip is rinsed by draining the solution through it. It is good practice to repeat the rinsing procedure at least one more time – this ensures that all impurities adhering to the inner surface are washed away. The burette is then filled with the standard solution up to the region of the zero mark and the tip is filled by opening the tap for a second or two.

The next task is to transfer a fixed volume of the solution of unknown concentration, ie the analyte, to a clean conical flask. A pipette is used and like the burette it too has to be rinsed. This is done by drawing a small volume of the analyte solution into the pipette and wetting its inner surface by tilting and rotating it. The ‘rinse’ solution is allowed to drain through the tip and discarded. After repeating the rinsing procedure, the pipette is filled with the analyte solution to a point above the graduation mark. With the pipette held vertically and with the graduation mark at eye level, the solution is allowed to slowly drain from the pipette until the bottom of the meniscus coincides with the graduation mark. Holding a white tile or a piece of white paper behind the stem of the pipette defines the meniscus more clearly. With the pipette tip placed well within the conical flask, the analyte solution is run into the flask. When free flow ceases, the tip should be touched against the inner wall of the flask to allow the remaining solution to drain. A few drops of the appropriate indicator are then added to the analyte solution in the flask.

Incidentally, if the conical flask had been wet with deionised water before adding the analyte solution to it, then no problem results – although the solution would be diluted, the number of moles of analyte would be unchanged and this is the critical factor.

Before reading the burette, its vertical alignment should be checked both from the front and the side. With a white tile behind the burette and with the eye level with the top of the standard solution, the burette is read from the bottom of the meniscus and the reading recorded. If the solution is dark and coloured, the bottom of the meniscus may not be clearly visible, in which case the reading is taken from the top of the meniscus. In reading a burette, it is important that the filter funnel used to fill it has been removed. If it were left in place, some drops of solution could drain from it during the titration, leading to a false titre volume.

The conical flask containing the analyte solution and indicator is placed underneath the burette, making sure that the tip of the burette is well within the neck of the flask. It is also good practice to have a white tile underneath the flask so that the colour change at the end-point can be seen more clearly.

The first titration is usually a rough one and its purpose is to see what the colour change is and to provide an approximate titre volume. In this rough titration, portions of the standard solution, about 1 cm3 at a time, are run from the burette into the conical flask. During and after the addition of each portion, the contents of the flask should be swirled – this helps the mixing process and gives the reactants time to react. These 1 cm3 additions are continued until the end-point is reached. The final burette reading can then be recorded. If the end-point proves difficult to assess, it is worthwhile keeping this rough titrated mixture to aid the detection of end-points in subsequent titrations.

A second but more accurate titration is then performed. A portion of the analyte solution is pipetted into a clean conical flask along with a few drops of indicator. The burette is refilled with the standard solution and the initial reading is recorded. Suppose the rough titre volume had been 20 cm3 then in the second titration it would be safe to add about 18.5 cm3 of the standard solution without any danger of over-shooting the end-point. However, care must be taken to ensure that the rate of delivery is not too fast otherwise the burette may not drain cleanly. This would leave drops of solution adhering to the walls of the burette, which in turn would lead to an inaccurate titre volume.

The titration is completed by adding the standard solution very slowly, drop by drop, while vigorously swirling the contents of the flask. The end-point of the titration is finally reached when the indicator just changes colour. The final burette reading should then be recorded. During the titration, should any of the standard solution splash onto the walls of the conical flask then wash it into the mixture with deionised water from a wash bottle. If near the end-point, you find a drop of the standard solution hanging from the tip of the burette, remove it by touching the tip to the wall of the flask and washing it into the solution.

The titrations are then repeated until concordant results, ie two consecutive titre volumes that are within 0.1 cm3 of each other, are obtained.
To carry out a titration quickly and efficiently, the recommended method of adding the solution from the burette to that in the conical flask is illustrated below.

The burette tip is manipulated with the left hand and this leaves the right hand free to swirl the contents of the conical flask as the burette solution is added. This technique is likely to feel awkward and clumsy at first but with practice it will become second nature to you.

Ideally what we try to obtain in a titration is the equivalence or stoichiometric point. This occurs when the quantity of reagent added from the burette is the exact amount necessary for stoichiometric reaction with the amount of reagent present in the conical flask. In practice, what we actually measure in a titration is the end-point and not the equivalence point and there is a subtle difference between the two. To illustrate the difference, let’s consider a permanganate/oxalate titration for which the stoichiometric equation for the titration reaction is:

5C2O42− + 2MnO4− + 16H+ → 10CO2 + 2Mn2+ + 8H2O

Up to and including the equivalence point all the permanganate ions added from the burette are consumed by the oxalate ions in the conical flask and the flask solution remains colourless. It is the first trace of a permanent pink colour that marks the end-point of the titration and for this colour to be exhibited extra permanganate ions, beyond those needed to react with the oxalate ions, are required. This means the end-point overshoots the equivalence point very slightly and hence the end-point of a titration can never coincide with the equivalence point.

As mentioned earlier, the three main titration types are:

• acid-base titrations in which the titration reaction is simply a neutralisation in which protons are transferred from the acid to the base
• redox titrations in which an oxidising agent is titrated against a reducing agent or vice versa. In such redox reactions, electrons are transferred from the oxidising agent to the reducing agent
• complexometric titrations, which are based on complex formation, ie a reaction between metal ions and ligands in which the ligand molecules or ions use their lone pairs of electrons to bind with metal ions. The most common ligand or complexing agent used in complexometric titrations is ethylenediaminetetraacetic acid – commonly abbreviated to EDTA. In alkaline conditions, EDTA has the following structure:

The EDTA ion is a hexadentate ligand and forms 1:1 complexes with metal ions. For example, nickel(II) ions react with EDTA ions to form a complex with the following octahedral structure:

[pic]

Most titrations are direct, ie one reagent is added directly to the other until the end-point is reached. In some situations, however, a direct titration may not be possible, in which case we have to resort to a technique known as a back titration. This involves adding a known but excess amount of one standard reagent to a known mass of the substance being determined (the analyte). After reaction between the two is complete, the excess amount of the standard reagent is determined by titration against a second standard reagent. Back titrations are used when:

• no suitable indicator is available for a direct titration
• the end-point of the back titration is clearer than that of the direct titration
• the reaction between the standard reagent and analyte is slow
• the analyte is insoluble.

Let’s consider an example. Suppose we wished to determine the percentage calcium carbonate in a sample of marble. Back titration has to be used here since marble is insoluble in water. In practice, a sample of the marble of accurately known mass is treated with a definite amount of hydrochloric acid, ie the volume and concentration of the acid are accurately known. An excess of acid is used and the amount remaining after neutralising the calcium carbonate is determined by titrating it against a standard solution of sodium hydroxide. The difference between the initial and excess amounts of hydrochloric acid tells us how much acid reacted with the marble, and with a knowledge of the stoichiometry of the calcium carbonate/hydrochloric acid reaction, the percentage calcium carbonate in the marble sample can be calculated.

Indicators
Indicators are compounds that allow us to detect the end-points of titrations. Typically they undergo an abrupt colour change when the titration is just complete. In general, an indicator reacts in a similar manner to the substance being titrated and so indicator choice will depend on the titration type: acid–base, redox or complexometric.

An acid–base indicator is normally a weak organic acid that will dissociate in aqueous solution, establishing the following equilibrium:

HIn(aq) + H2O(l) H3O+(aq) + In−(aq)

It is able to act as an indicator because it has one colour in its acid form (HIn) and a different colour in its conjugate base form (In−).
If we examine the following table in which the properties of a selection of some common indicators are presented, we see that an acid–base indicator changes colour over a range of about 2 pH units and not at a specific pH.

|Indicator |HIn colour |pH range of colour change |In− colour |
|Bromophenol blue |Yellow |3.0–4.6 |Blue |
|Methyl red |Red |4.2–6.3 |Yellow |
|Bromothymol blue |Yellow |6.0–7.6 |Blue |
|Phenol red |Yellow |6.8–8.4 |Red |
|Phenolphthalein |Colourless |8.3–10.0 |Pink |

Choosing an indicator for a titration depends on the type of acid–base reaction taking place. There are four different types and these are outlined in the following table together with the pH values at their equivalence points.

|Acid–base reaction type |pH at equivalence point |
|Strong acid/strong base |7 |
|Weak acid/strong base |>7 |
|Strong acid/weak base |