The concentration of copper ions



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This Project Page first appeared in the September 1999 issue of Chemistry Review, Volume 9, Number 1, Pages 8-9. Chemistry Review is published four times during the academic year by Philip Allan Updates and is a journal for post-16 students. It contains a variety of interesting and colourful articles aimed at 16-19 year-olds taking mainly AS and A2 courses in chemistry.

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The concentration of copper ions

When we carry out chemical analyses we usually take it for granted that the results we obtain are correct. But are we justified in making this assumption? One idea for an interesting project would be to compare a set of methods designed to measure the same thing - for example, the concentration of copper(II) ions in a solution.

If you set out to find the concentration of copper(II) ions in a solution, there are many techniques you can use. Quite a few of them involve the chemistry you meet during an A-level course. But do all the methods give you the same result? If all the techniques are in agreement at a concentration of 0.1mol dm-3 copper(II) ions, what happens if you use a concentration of 0.01 or even 0.001mol dm-3? Do they all work well if the solution is made up starting with a brass screw? A comparison of the effectiveness and relative accuracy of different methods of finding the copper(II) ion concentration will provide you with a challenging and wide-ranging investigation.

You will need to start by making up a large quantity of the copper(II) ion solutions which you are going to analyse by different methods. You could start your investigation with a redox titration. Copper(II) ions react with excess iodide ions to form a precipitate of copper(I} iodide and molecular iodine. You can titrate the iodine against a standard solution of sodium thiosulfate, using starch indicator.

2Cu2+(aq) + 4I-(aq) → 2Cul(s) + I2(aq)

I2(aq) + 2S2O32-(aq) → 2I-(aq) + S4O62-(aq)

The copper(II) ions can also be titrated against a solution of EDTA in a complexiometric titration. The indicator to use is called Fast Sulphon F.

An acid-base titration would not at first sight seem to be very helpful. But it can be. If you add sufficient cation exchange resin in its hydrogen form, all the copper(II) ions should stick on the resin and displace twice as many hydrogen ions.

2Resin-H(s) + Cu2+(aq) → Resin-Cu(s) + 2H+(aq)

Filter off the resin and titrate the solution with a standard sodium hydroxide solution.

A different approach is to use a gravimetric or weighing method. Zinc powder will displace copper metal from a solution of copper(II) ions.

Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

You will need to remove excess zinc with sulfuric acid before drying and weighing your copper. This reaction between zinc metal and copper(II) ions is exothermic. Is it possible to measure the heat evolved and to link this with the concentration of copper(II) ions in the solution?

The blue colour of your solution is caused by the Cu2+(aq) ion. The more concentrated the solution the more intense the colour is. You can make up known concentrations of copper(ll) ion solution and use a colorimeter to produce a copper(ll) ion concentration/absorbance calibration curve. If you now put the solution you are trying to analyse in the colorimeter, you will be able to read off its concentration from the curve. Adding excess concentrated ammonia solution to a solution of copper(ll) ions produces a deep blue solution containing the [Cu(NH3)4(H20)2+ complex ion. You could devise a colorimetric method based on this solution. You might also like to try solutions in which the Cu2+ ion is complexed with other ligands such as Cl- or EDTA.

A completely different approach to the task is based on electrochemical cells. You can make a half-cell by dipping a strip of copper into a solution of copper(ll) ions. Its electrode potential depends upon the concentration of the copper(ll) ions. If you connect this to another half-cell, such as a zinc/ zinc ion half-cell, then the cell potential you can measure with a high resistance voltmeter also depends on the concentration of copper(ll) ions. You can investigate the link between copper(ll) ion concentration and cell potential by using different copper ion solutions, and you can represent your results in the form of a copper(ll) ion/cell potential calibration curve. Alternatively, you can look up and use the Nernst Equation which relates electrochemical cell potential to a number of variables, including ion concentration.

As you can see, there is quite a lot of chemistry involved in finding the concentration of copper(II) ions in a solution.

|[pic] |

|The depth of colour of a copper sulfate solution is caused by the concentration of the |

|Cu2+(aq) ion. The more concentrated the solution, the more intense the colour. |

Derek Denby

Derek Denby is Director of the Science Centre of Excellence at John Leggott College, Scunthorpe

The original article was written by Derek Denby. We are grateful to Derek for allowing us to reproduce it here.

This page is free for your personal use, but the copyright remains with Philip Allan Updates. Please do not copy it or disseminate it in any way.

Chemistry Review is indebted to Don Ainley, who has helped to prepare this article for the Web.

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John Olive

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