A large number of inorganic and organic species may be determined spectrophotometrically. The species measured must be capable of absorbing uv or visible radiation efficiently.  If a compound is to be analysed by this method and does not meet this criterion, then it must undergo some form of reaction so that its modified form will do so.  This type of colour-producing reaction is commonly referred to as a chromogenic reaction.

The solution presented for testing should ideally have the following properties:

1.     the absorption should be sufficiently intense as to allow the determination of small amounts of the analyte species,

2.    the absorbing species should form rapidly,

3.    the absorbing species should be stable,

4.    the absorption should be unaffected by small changes in solution temperature and pH,

5.    the reaction forming the absorbing species should be specific and stoichiometric (or reproducibly non-stoichiometric).

In this experiment we will determine iron using a chromogenic reaction with the reagent 1,10-phenanthroline.  This forms a coloured complex with iron.

The reaction is as follows:

Fe2+ + 3(1,10-phen)  ?  [Fe(1,10-phen)3]2+

The complex produced absorbs light strongly and in a different wavelength region to the free species. A chromogenic reaction is one example of a pre-treatment procedure that might be used in a chemical analysis.

Into a 50 cm3 volumetric flask, add hydroxylamine hydrochloride solution (5 cm3, 10% w/v), acetate buffer (10 cm3, 0.10 mol dm?3) and 1,10-phenanthroline solution (5 cm3, 0.25% w/v). Pipette into the flask 20 cm3 of unknown iron solution. Dilute the contents of the flask to the line with deionised water (The bottom of the meniscus should be level with the line, use a dropping pipette to add the last ml or so to prevent overshooting). Mark this as solution A.  Repeat the above procedure in a fresh 50 cm3 volumetric flask and mark the resulting flask, solution B. These are the sample solutions.

Measure the absorbance at 510 nm of the standards provided and of sample solutions using the blank as a reference.  Load the solutions into the cell using a dropping pipette.  Handle the cell by the opaque sides only.  Ensure the clear sides are clean, dry and unmarked and that there are no air bubbles in the cell.

Plot a graph of absorbance against concentration for the standards and use it to determine the concentration of iron in the sample solution and hence in the unknown solution.

Additional Calculation

Suppose the unknown solution (total volume 100cm3) was made up by dissolving 3.7756 g of a solid in 100 cm3 of deionised water.  Calculate the concentration of iron in the solid based on the results obtained above.




Spectrometer used:

Concentration of
iron in solution
/mol dm?3    Absorbance of solution at 510 nm









Absorbance of sample solution 1                = 0.831

Absorbance of sample solution 2                = 0.815

Concentration of sample solution 1                 = 5.211

Concentration of sample solution 2                 =5.110

Average concentration                     =5.1605

Concentration of iron in unknown solution            =



High performance Liquid Chromatography (HPLC) is undoubtedly the most widely used form of chromatography for separating mixtures of compounds and materials into their separate components.  As with all chromatographic methods, two mutually immiscible phases are brought into contact in HPLC; one phase is stationary and the other is mobile.  The stationary phase usually consists of very small (5-10 µm in diameter) porous silica gel beads which are used dry or coated with a liquid – in the latter case the beads and coating are still called the stationary phase.  The mobile phase is usually a mixture of solvents and it is into this phase that the sample mixture is injected.  Clearly the mobile phase is a liquid and the stationary phase is a solid – hence they are immiscible.

The HPLC technique separates components of mixtures by the same principle which applies to all forms of chromatography, i.e. as the sample is carried through the stationary phase by the mobile phase, each component of the mixture interacts with the stationary phase to a different extent and, as a result, is retarded in its progress through the stationary phase to a different extent.  The stationary phase is packed into a column through which the mobile phase and sample mixture pass, and so the least retarded component emerges (or elutes) from the column first, and the most strongly retained component elutes last.  In this way, the components of the mixture can be separated, eluting from the column one after the other.

The choices of stationary and mobile phases are critical to the success of the separation, and considerable effort is needed to optimise these (and other) factors for each particular analysis.  HPLC has an advantage over some other forms of chromatography [e.g. Gas Chromatography (GC)] in that the mobile phase (i.e. the solvent mixture) can be chosen to get the best separation, rather than just the stationary phase.  This greater flexibility of conditions results in HPLC being applicable to a much wider range of problems (than, for example, GC).  HPLC is able to achieve separations for a very wide range of materials, including macromolecules, natural products, pharmaceuticals, polymers, metabolites and the impure results of syntheses!

A successful HPLC analysis not only separates the components, but also allows them to be identified and quantified (i.e. their relative amounts cam be obtained).  Both identification and quantification are achieved by using standard or calibration solutions.

The columns used in HPLC are very tightly packed with the solid support (the silica gel).  If the mobile phase were left to pass through the column by gravity, separations would take weeks!  In order to make the mobile phase pass through the column in a reasonable time (minutes), the solvent is forced through under high pressure.  In comparison to a gravity-fed chromatography column, the HPLC column is indeed High Performance!

As they elute from the column, the components of the mixture are detected by their absorbance in the UV region.  A source provides UV radiation (at a selected wavelength) which is passed through the mobile phase containing the particular component which has just emerged from the column.  The change in intensity of the radiation as a result of it being absorbed by the component is detected, and this gives a measure of how much of the component is present.  Other types of detector are available (e.g. thermal conductivity detector), but UV detector is fitted to the HPLC instruments used in this Experiment.

The HPLC technique is crucial to the research and development work of many industries, but of the pharmaceutical industry in particular – HPLC instruments  are probably the most common pieces of sophisticated analytical equipment in such laboratories.  This Experiments demonstrates the basic technique by challenging you to separate and quantify an unknown mixture of two pharmaceutical products, Paracetamol and Aspirin.


For an explanation of HPLC you could read either

Fundamentals of Analytical Chemistry (7th Edition) Skoog, D.A., West, D.M. and Holler, F.J., Chapter 30.
Exploring Chemical Analysis (4th Edition) Harris, D.C., Chapter 22


In order to get the most (and the best!) out of this experiment, you must come into the laboratory with an understanding both of the essential parts of an HPLC instrument, and the principle by which separations are achieved on the instrument.  Specifically, you should have a knowledge and understanding of the following:

*    the stationary phase and the mobile phase;
*    how components of a mixture are separated on the column;
*    the essential components of an HPLC instrument (eluent, solvent reservoirs, high pressure pump, sample injection loop, column, detector).


You will be provided with two calibration solutions, each of 25 cm3 in volume, and each containing accurately known amounts of the components noted below.  The concentration of each component should be specified on the information sheet provided during the experiment, but one calibration solution should have component concentrations at lower levels and the other calibration solution should have component concentrations at higher level.

An unknown sample will also be provided, the concentration of aspirin and paracetamol will be within the range covered by the two standards.

Injecting a sample

Note:    The pressure in the HPLC system should be checked regularly throughout the experiment.  This is displayed on the HPLC instrument panel under the heading Bar.  If the reading displayed becomes eratic or drops greatly (?5 bars is acceptable), you should ask for assistance immediately.

Set the sample injection loop to LOAD using the control knob.  Choose one of the calibration solutions and rinse the syringe twice with this solution.  Using the syringe, inject 25 µl of the chosen solution into the loop (i.e. the hole in the centre of the control knob); immediately repeat this operation.  Switch the control knob to INJECT and immediately press the [INJECT] button on the data capture unit appropriate to the channel which you are using.  The chromatogram will automatically be displayed on the monitor as it runs, and when complete, the chromatogram will be printed out as a record of the result.

Collection of data for calibration and unknown solutions.

Inject a sample of each calibration solution twice, and on the unknown sample solution twice, hence generating 6 chromatograms in total.  Please wait for FULL elution of each injection before making the next one.

Note:    If for any reason an injection has gone wrong or needs to be aborted, ask for staff assistance immediately.


Using the areas of appropriate peaks computed by the Data Station, calculate the correction factor for paracetamol (Fpara) for each of the calibration injections (only!) using the equation:

Fpara = (Cis . Apara)/ (Cpara . Ais)    where:    Cpara = concentration of paracetamol
Cis = concentration of the internal standard
Ais = area under internal standard peak
Apara = area under paracetamol peak

Use the analogous equation to calculate Fasp for aspirin.

Calculate an average value for Fpara and for Fasp from the calibration injections, and hence calculate the concentrations of paracetamol and aspirin in the unknown solution.

Additional Question

Caffeine is used here as an Internal Standard.  Can you explain why an Internal Standard is used and what criteria might be used to choose one.




Concentration of caffeine in Test Solution            =            g dm-3

Concentration of Aspirin in Low Calibration Solution    =

Concentration of Aspirin in High Calibration Solution    =

Concentration of Paracetamol in Low Calibration Solution    =

Concentration of Paracetamol in High Calibration Solution    =

Concentration of Caffeine in Low Calibration Solution    =

Concentration of Caffeine in High Calibration Solution     =

Area of Aspirin Peak

(Aasp)    Area of Paracetamol Peak
(Apara)    Area of Caffeine Peak
Low Calibration Solution
First Injection
Low Calibration Solution
Second Injection
High Calibration Solution
First Injection
High Calibration Solution
Second Injection
Test Solution
First Injection

Test Solution
Second Injection

Values for Low Calibration:

First Injection:        Fpara =                 Fasp =

Second Injection:    Fpara =                 Fasp =

Average values    Fpara =                 Fasp =

Values for High Calibration:

First Injection:        Fpara =                 Fasp =

Second Injection:    Fpara =                 Fasp =

Average values    Fpara =                 Fasp =

Combined average values:

Fpara =                 Fasp =

Concentration of Aspirin in Test Solution                 =

Concentration of Paracetamol in Test Solution            =

Operational Parameters

Type of column used:

Composition of eluent:

Flow Rate:




This is an experiment which covers two important areas. The glass electrode is an example of a membrane based ion sensitive electrode (ISE). The use of a reference electrode with a “selective” membrane, in this case a special type of glass, produces a system which has an electrical potential which responds to the concentration of a particular species, in this case the hydrogen ion. The practical also illustrates the application of the glass electrode to the measurement of pH, a quantity which is routinely monitored for chemical and biological processes.

References: The appropriate manual for the pH meter that you use – this will be in 10/2. You should read the relevant sections before using the meter.
Skoog, West, and Holler  (7th Edn.), pp. 392-400.


Make up solutions (100 cm3) of NBS standard tetroxalate, tartrate, phthalate, phosphate, borax and calcium hydroxide buffers as follows.

Tetroxalate    Dissolve potassium tetroxalate (1.270 g) in deionised water and dilute to 100 cm3.

Tartrate    This is provided as a ready made solution.

Phthalate    This is provided as a ready made solution.

Phosphate    Dissolve potassium dihydrogen phosphate (0.340 g) and disodium hydrogen phosphate (0.355 g) in carbon dioxide-free water and dilute to 100 cm3.

Borax        Dissolve sodium tetraborate (0.381 g) in carbon dioxide-free water and dilute to 100 cm3.

Choose one of the pH meters and calibrate it with two standards in accordance with the instructions in the manual.  Measure the pH, emf, and temperature of each solution. In your results sheet you will have expected pH values for each buffer at 20?C.

Using Excel , plot graphs of emf against expected pH value and ?pH against expected pH value.
(?pH = observed pH – expected pH.)

Measure the slope of the first graph. The slope should be 59.1 mV for each pH unit.

To draw the second graph, join adjacent points with straight lines.

A saturated calcium hydroxide solution is provided ready made. Measure the pH of this solution. Search on the internet for a reference value – how closely does yours agree?

It should be remembered at all times that pH electrodes are delicate objects. Treat them with care. Do not bump them against the sides or bottom of beakers. When using the electrode, adopt the following procedure.

1.    Always keep the electrode in a beaker of deionised water unless taking a measurement.

2.    When you wish to make a measurement, remove the electrode from the water and pat dry gently with a piece of tissue or a filter paper. Place the electrode into the test solution. The electrode is now ready to give a reading.

3.    After you have taken the reading, remove the electrode and rinse it with deionised water. Pat dry as before and either return it to the deionised water or take another reading.





tetroxalate    Na2HPO4    KH2PO4    borax
boat plus sample /g    2.1534
0.9486    0.9071    0.9523
boat empty /g    0.8809
0.5930    0.5670    0.5712
sample /g    1.2725
0.3556    0.3401    0.3811

pH meter used:

pH    OBS.
/?C    EXP.
AT 20?C
1.68    274    24.1    1.68    0
3.20    190    19.4    3.56    -0.36
4.00    146    20.0    4.00    0
6.88    -13    20.8    6.88    0
9.28    -144    22.4    9.23    0.05

Measurement of pH of Saturated calcium hydroxide solution = 11.42



Titrimetry is a simple and flexible method of absolute analysis. Titrimetric methods have been devised for most simple inorganic species – anions, cations and molecules. As with gravimetry the method relies on accurate measurement of a physical property, in this case volume, and a reliable knowledge of the stoichiometry of any reaction taking place.

In this experiment we will be looking at the concentration of an anion. Here the analyte ion, chloride, reacts one-to-one with a cation, silver, to give a precipitate of silver chloride. Such titrations are often called precipitation titrations. It is possible to use a variety of methods for following the titration such as turbidimetry or conductimetry

Reference: Skoog, West, and Holler, Chapter 6 (particularly pp 100-117)


Obtain a sample of unknown from the stores. Record the approximate concentration given on the sample container on your result sheet. Weigh out accurately ca. 1.46 g of the unknown and dissolve in a sufficient quantity of deionised water in a 250 cm3 volumetric flask. Dilute to volume.

Pipette 25 cm3 of this solution into a 250 cm3 conical flask and add 1 cm3 of indicator (potassium chromate/dichromate) solution. Titrate with standard silver nitrate solution. With each addition of silver nitrate a red colour appears in the solution. As this begins to persist for longer after shaking, the end-point is approaching. The end-point is marked by the appearance of a faint but permanent red-brown colour.

Repeat the titration three times. Your values should agree to within 0.1 cm3. Repeat the titration until you have three values which meet this criterion.

Calculate the concentration of chloride for each of these three titrations. Compute an average value for the concentration and from this calculate the concentration of chloride (in wt%) in the unknown sample.


Your answer for the concentration will contain an error due to the indicator blank. What is this and how might one attempt to nullify its effect? The indicator used for this titration is chromate – how does it work? Find one other indicator that can be used for the titration of chloride with silver nitrate and explain how it works.


NAME:                            DATE:


Weight of bottle plus sample        =

Weight of empty bottle        =

Weight of sample taken        =



/mol dm-3

Mean concentration =

Concentration of chloride in sample =
(Show your calculation for this below)

find the cost of your paper

This question has been answered.

Get Answer