Iodometric Determination of Cu in Brass

by Ulrich de la Camp and Oliver Seely [1] (Copied with kind permission and with no liability accepted for the current content.)

Discussion

In acid solution practically all oxidising agents will oxidise iodide ions to iodine quantitatively. The iodine formed in the reaction can then be titrated by means of a standard sodium thiosulphate solution. This type of indirect titration is given the general name of iodometry.

Iodometric methods of analysis have a wide applicability for the following reasons:

1. Potassium iodide, KI, is readily available in high purity.

2. A good indicator, starch, is available to signal the equivalence point in the reaction between iodine and thiosulphate. Starch turns blue-black in the presence of iodine. Therefore, when the blue-black colour disappears, the iodine has been completely reduced to the iodide ion.

3. Iodometric reactions are rapid and quantitative.

4. A precise and stable reducing agent, sodium thiosulphate (Na₂S₂O₃), is available to react with the iodine.

The amount of iodine liberated in the reaction between iodide ion and an oxidising agent is a measure of the quantity of oxidising agent originally present in the solution. The amount of standard sodium thiosulphate solution required to titrate the liberated iodine is then equivalent to the amount of oxidising agent. Iodometric methods can be used for the quantitative determination of strong oxidising agents such as potassium dichromate (VI), manganate (VII) ('permanganate'), hydrogen peroxide, copper (II) ('cupric') ion and oxygen.

As has been mentioned above, the endpoint in a titration of iodine with thiosulphate is signaled by the colour change of the starch indicator. When starch is heated in water, various decomposition products are formed, among which is beta-amylose which forms a deep blue-black complex with iodine. The sensitivity of the indicator is increased by the presence of iodide ion in solution. However, if the starch indicator solution is added in the presence of a high concentration of iodine, the disappearance of the blue-black colour is very gradual. For use in indirect methods, the indicator is therefore added at a point when virtually all of the iodine has been reduced to iodide ion, causing the disappearance of the colour to be more rapid and sudden. The starch indicator solution must be freshly prepared since it will decompose and its sensitivity is decreased. However, a properly prepared solution will keep for a period of a few weeks. A preservative such as a small amount of mercuric ions may be added to inhibit the decomposition.

Solutions of sodium thiosulphate are made up to an approximate concentration by dissolving the sodium salt in water that has previously been boiled. Boiling the water is necessary to destroy micro-organisms which metabolize the thiosulphate ion. A small amount of Na₂CO₂ is added to the solution in order to bring the pH to about 9. The solution is standardised by taking a known amount of oxidising agent, treating it with excess iodide ion and then titrating the liberated iodine with the solution to be standardised. Oxidising agents such as potassium dichromate (VI), bromate (V), iodate (V) or copper (II) can be employed for this procedure. You will be using potassium iodate (V), KIO₃, as your primary standard. The reaction between IO₃- and I- is given as

6 H⁺ + IO₃^- + 5 I^- ---> 3 I₂ + 3 H₂O

Reactions Involved in Iodometric Processes.

Iodometric methods depend on the following equilibrium:

I₂ + I^- <===> I₃^-

Since the solubility of I₂ in water is quite low, the formation of the tri-iodide ion, II₃^-, allows us to obtain useful concentrations of I₂ in aqueous solutions. The equilibrium constant for this reaction is approximately 700. For this reason iodometric methods are carried out in the presence of excess iodide ion.

The reaction between iodine and the thiosulphate ion is:

I₂ + 2 S₂O₃^2- <===> 2 I^- + S₄O₆^2-

This reaction proceeds quantitatively in neutral or slightly acidic solutions. In strongly alkaline or acidic solutions the oxidation of the thiosulphate does not proceed by a single reaction. In the former, the thiosulphate ion is oxidised to sulphate as well as to the tetrathionate. In the latter, the thiosulphuric acid formed undergoes an internal oxidation-reduction reaction to sulphurous acid and sulphur. Both of these reactions lead to errors since the stoichiometry of the reactions differs from that shown above for the thiosulphate as a reducing agent. The control of pH is clearly important. In many cases the liberated iodine is titrated in the mildly acidic solution employed for the reaction of a strong oxidising agent and iodide ion. In these cases the titration of the liberated iodine must be completed quickly in order to eliminate undue exposure to the atmosphere since an acid medium constitutes an optimum condition for atmospheric oxidation of the excess iodide ion.

The basic reaction in the determination of copper using the iodometric method is represented by the equation:

2 Cu²⁺ + 4 I^- <===> 2 CuI(s) + I₂

This is a rapid, quantitative reaction in slightly acidic solutions, if there is a large excess of iodide ion present and if the copper is in the form of a simple ion rather than a complex one. The iodine that is liberated can be titrated in the usual manner with standard thiosulphate solution. The reaction involving copper (II) ion and iodide takes place quantitatively since the copper (I) ('cuprous') ion formed as result of the reduction is removed from the solution as a precipitate of copper (I) iodide.

Iron interferes since iron(III) ions will oxidise iodide. Since the iron will be found in the +3 oxidation state as a result of the dissolution of the brass sample, a means of preventing this interference is necessary. This can be accomplished by converting the iron(III) to a soluble iron(III) phosphate complex using phosphoric acid. At a pH of 3.0-4.0 the iron phosphate complex is not reduced by iodide ion. If arsenic and antimony are present, they will provide no interference at this pH if they are in their higher oxidation states. Brass formulations also may contain up to 39% Zn, 2.5% Sn and 8.5% Pb. When dissolved in concentrated nitric acid, the zinc and the lead become Pb²⁺ and Zn²⁺ . These do not interfere with the analysis of copper because they are not reduced to the Pb⁺ and Zn⁺ states by the action of iodide ion under the conditions of this experiment. The tin is oxidised to Sn⁴⁺ by the concentrated nitric acid and after dilution and adjustment of pH this form becomes SnO₂ which is insoluble and may be observed as an inert white precipitate at the bottom of your flask. Under these conditions the tin does not interfere with the analysis.

Sources of Error.

The following are the most important sources of error in the iodometric method:

1. Loss of iodine by evaporation from the solution. This can be minimised by having a large excess of iodide in order to keep the iodine tied up as tri-iodide ion. It should also be apparent that the titrations involving iodine must be made in cold solutions in order to minimise loss through evaporation.

2. Atmospheric oxidation of iodide ion in acidic solution. In acid solution, prompt titration of the liberated iodine is necessary in order to prevent oxidation.

3. Starch solutions that are no longer fresh or improperly prepared. The indicator will then not behave properly at the endpoint and a quantitative determination is not possible.

Experimental

Click here to see images of what one can expect in this experiment.

Preparation of a 0.10 mol dm^-3 Standard Na₂S₂O₃ Solution

With a graduated cylinder measure out 1 dm³ (1 litre) of distilled water. Place it in your 1 dm³ beaker and boil the water for at least 5 minutes. Weigh out 25 g of Na₂S₂O₃ · 5 H₂O and 0.1 g of Na₂CO₃. Dissolve the thiosulphate in the hot water and then cool this solution with the aid of an ice bath to room temperature. Then add the carbonate and stir until it is completely dissolved. Transfer the solution to your plastic 1 dm³ bottle. When not in use store this bottle in the darkness of your equipment cabinet as the decomposition of thiosulphate is catalysed by light.

Blank Determination.

Potassium iodide may contain appreciable amounts of iodate ion which in acid solution will react with iodide and yield iodine. The liberated iodine would react with thiosulphate and thereby cause the apparent molarity of the thiosulphate to be too low. The following procedure allows for the determination of a blank correction which will properly correct for any iodate that might be present. Prepare a solution of exactly 2.00 g of KI dissolved in 50 cm³ of distilled water and then acidify the solution with 5 cm³ of 3 mol dm⁻³ sulphuric acid and then immediately add 5 cm³ of starch indicator. If a blue-black colour appears right after mixing, use the thiosulphate solution in the burette to determine the volume of solution required to cause the colour to disappear. This volume must be subtracted from the standardisation and analyses volumes. If the potassium iodide is completely iodate-free no colour will of course develop and no blank correction is necessary.

Standardisation of the Na₂S₂O₃ Solution

Dry approximately 2 g of KIO₃ at a temperature of 110 °C for one hour. Weigh to a precision of ± 0.0001g three samples of the iodate having masses near 0.12 g directly into three 250 cm³ conical flasks. Dissolve the iodate in 75 cm³ of distilled water. Cover the flasks with parafilm and store them. Rinse and fill your burette with the solution. If a blank correction is required add exactly 2.00 g of KI to each of the potassium iodate solutions. If no blank determination is required, the exact amount of KI is not crucial but should be close to 2 g. Then add 10 cm³ of 1 mol dm⁻³ HCl to one of the solutions. It will turn a dark-brown colour. Immediately titrate it with the thiosulphate solution. When the colour of the solution becomes very pale yellow add 5 cm³ of starch indicator. Continue the titration until the blue colour of the starch complex just disappears. Follow the same procedure with each of the other two solutions, first adding the HCl then titrating. Correct your titration data for burette error and if necessary apply the blank correction. Calculate the molarity of the Na₂S₂O₃ solution. Results should agree to within 0.2% of the average. If you do not achieve that kind of precision, titrate additional samples.

Dissolution of the Brass Sample.

The following procedures in this section make use of the hot plates in the fume hoods. The solutions of dissolved brass generally have a low volume and high acid and salt concentrations. "Bumping" or little explosions of steam in the superheated liquid can occur. You do not want your hand to be close to the mouth of the flask should the solution suddenly "bump" because drops of acid (not to mention part of your sample) will fly out of the flask and possibly onto your hand. For that reason you must use your tongs to place the flasks on the hot plate and to remove them. Do not use strips of paper towel or the rubber Hot Hands because your real hand will end up being too close to the mouth of the flask.

The brass sample which you will receive does not have to be dried before use. Accurately weigh out three brass samples, of about 0.3 g each, directly into separate 250 cm³ conical flasks. In the fume hood add 5 cm³ of 6 mol dm⁻³ HNO₃. Warm the solution on a hot plate in the fume hood until dissolution is complete. Add 10 cm³ of concentrated (not 3 mol dm⁻³) H₂SO₄ and continue heating until white SO₃ fumes appear. It is important that you do not mistake ordinary water vapour for SO₃ fumes. It is also important at this point that the flask not be removed from the hood. SO₃ fumes are dangerous and ought not to be inhaled. Only when the slightly denser white fumes of SO₃ are observed can you be sure that all HNO₃ has been removed. NO₃^- will oxidise I^- and hence will seriously interfere with the procedure. Cool the flask in air for one or two minutes and then in an ice bath, then carefully add 20 cm³ of distilled H₂O. Boil for one or two minutes then again cool in an ice bath. Continue to keep the flask in the ice bath and using your medicine dropper add concentrated NH₃(aq) dropwise, and with adequate mixing, until the light-blue colour of the solution is completely changed to the dark-blue colour of the copper tetraammine complex. As many as 400 drops (20 cm³) may be required. The solution must be kept cool in an ice bath since the reaction between the concentrated H₂SO₄ and concentrated NH₃ is highly exothermic. Now add 3 mol dm⁻³ H₂SO₄ dropwise until the dark-blue colour just disappears. You do not have to produce a complete disappearance of the dark blue colour but you need to approach that point. The subsequent addition of phosphoric acid will lower the pH appropriately to around 3.5. If you add too much 3 mol dm⁻³ H₂SO₄ the pH may turn out to be sufficiently low to cause unwanted side reactions to occur when you reduce the Cu²⁺ with iodide. If you are uncertain about the disappearance of the dark blue colour you may put 50 cm³ of 0.06 mol dm⁻³ Cu²⁺ in a clean 250 cm³ flask and add 12 mol dm⁻³ ammonia dropwise until you have that unmistakable dark blue colour. Then add 3 mol dm⁻³ H₂SO₄ dropwise until the blue colour almost disappears. Then add 2 cm³ concentrated phosphoric acid and you ought to see the dark colour completely disappear. You may copy that procedure to achieve an appropriate pH of around 3.5 for subsequent steps in the analysis. Now, back to your real sample: Once you are confident that you haven't added too much 3 mol dm⁻³ H₂SO₄, but that you have caused the dark colour of the copper tetraammine complex almost to disappear, add 2.0 cm³ of concentrated phosphoric acid, H₃PO₄, to each sample. Verify to yourself that they exhibit the light copper colour rather than the dark colour and cover the flasks with parafilm and set them aside until you are ready to proceed with the titration.

Titration of the Dissolved Brass Sample.

If you have let the dissolved samples stand overnight, be sure to warm the sample on a hot plate (this can be done at your desk) in order to dissolve all larger crystals of copper sulphate that might have formed. Be sure to cool the samples to room temperature, or below, with the aid of an ice bath. The solutions will still contain a fine, white precipitate at this point; however, this will not interfere with the rest of the procedure. From this point on work with only one sample at a time. Add 4.0 g of KI to one of your samples and titrate immediately with the standard thiosulphate solution. The sample contains white CuI precipitate and the color of I₃^- must be observed against that precipitate. The slurry will at first appear brown or dark yellow-brown. Continue adding thiosulfate until the slurry is a light mustard colour. At this point add 5 cm³ of starch indicator and titrate until the mixture in the flask takes on a milky pink or lavender hue. Now add 2 g of KSCN and mix well; the solution will darken somewhat. After the addition of thiocyanate, continue to add more thiosulphate dropwise. You should observe a sudden change to a white or cream colour. That is the endpoint of the titration. After you have titrated all three samples calculate the percentage of Cu in each of the brass samples, the average percentage and the average deviation.

The description above applies for brass samples with low concentrations of zinc (<10%). Some of you may have brass samples with higher concentrations of zinc. Such samples will become quite dark after the addition of KI and will lighten only slightly as thiosulphate is added. The "mustard colour" will be darker than samples having low percentages of copper. When the starch is added the sample will become dark blue-black again and as you approach the end point with the thiosulphate the slurry will turn a violet colour rather than milky pink or lavender hue. With the addition of the KSCN the solution will darken somewhat as in the case of the other samples, but the final end point will be a bit darker than the white or cream colour described above. If you think that you have a sample with high zinc content, observe your progress carefully and take notes which will allow you to achieve repeatability.

Explanation: The reduction of Cu²⁺ to Cu⁺ occurs as the result of the oxidation of I^- to I₂. The I₂ combines with iodide ion to produce the dark brown triodide ion, I₃^-. The excess iodide ion also causes the reduced copper to precipitate as white copper (I) iodide, CuI. I₂ and I₃^- in solution tend to adsorb on the surface of CuI thus becoming unavailable for rapid reduction by the thiosulphate. As a result, iodometric titrations involving reduced copper tend to yield lower results unless the adsorbed I₂ can be liberated by adding thiocyanate ion, SCN^-, which competes with the adsorbed iodine molecules on the surface of solid particles of CuI. After the addition of thiocyanate, continue to add more thiosulphate dropwise. You should observe a sudden change to a white or cream colour. That is the endpoint of the titration. After you have titrated all three samples calculate the percentage of Cu in each of the brass samples, the average percentage and the average deviation.

Report

Your report must include the following information in the two sections below.

1. Unknown number

2. The three weights of KIO₃ used for the standardisation of thiosulphate

3. Volume in cm³ of thiosulphate for each of your standardisation titrations

4. Average molarity of the thiosulphate solution

5. Weight of brass used for each sample

6. Volume of thiosulphate solution used for each sample

7. Cu percentage for each sample

8. The average percentage of Cu in your brass sample

9. The average deviation from the mean of the percent Cu for the three samples

10. Pages in your lab notebook containing the pertinent data.

Questions on Cu in Brass Analysis

1. Why is it necessary to boil the water used to prepare the thiosulphate solution?

2. Why is Na₂CO₃ added to the thiosulphate solution?

3. Why is the thiosulphate solution stored in the dark?

4. Why is HCl added to the IO₃^- mixture and why must the solution be titrated immediately?

5. Why is the solution containing the dissolved brass sample heated to expel SO₃ fumes?

6. Why is H₃PO₄ added to the brass sample?

7. What is the purpose of the KSCN that is added just before the endpoint in the titration?

8. Why is the solution containing the dissolved brass made basic with concentrated NH₃ and then again acidified with H₂SO₄?

9. What is the formula of the tetrammine copper(II) complex?

10. Why do Zn²⁺ and Pb²⁺ not interfere in this procedure?

11. What sort of complications would arise if the iodine/thiosulphate titration were carried out in a highly acidic solution?

12. If the solution were highly basic, how would the iodine/thiosulphate reaction be influenced? 13. Why is the starch indicator not added at the beginning of the titration?