Experiment 35
Preparation of Copper(I) Chloride, CuCl

Goal: To prepare copper(I) chloride by reducing copper(II) chloride with sulfite ions in the presence of chloride ions.

Objectives: At the end of this session students should be able to:
Copper can exist as Cu(I), Cu(II) which is most common and Cu(III) which is rare. The relative stabilities of each oxidation state depend on the nature of ligands and anions as well as the nature of the solvent medium. In the experiment below the copper(I) ions once formed, react with chloride ions to form the insoluble copper(I) chloride.

The relevant redox potentials are:
        Cu+ + e- ↔ Cu(s)         E° = 0.52 V
        Cu2+ + e- ↔ Cu+         E° = 0.15 V
        Cu2+ + 2e- ↔ Cu(s)        E° = 0.34 V
 Hence Cu(s) + Cu2+ ↔ 2Cu+         E° = -0.37 V

WARNING! Glacial acetic acid is a dangerous liquid and must be used in a fume cupboard.

Prepare three solutions:
Add slowly, with constant stirring, the sodium sulfite solution to the copper(II) chloride solution. Dilute the suspension of copper(I) chloride so formed with about half the sulfurous acid solution, allow the precipitate to settle, and decant most of the supernatant solution.

Filter the solid by suction on a Buchner funnel or sintered glass crucible (fumehood) and wash it into the funnel by means of the sulfurous acid solution. Take care that the copper(I) chloride is always covered by a layer of solution. Finally, wash the product with 5 cm3 of glacial acetic acid and 10 cm3 of ethanol and dry it in air, or in a vacuum desiccator. Determine and record the yield.

Immediately show the product to a demonstrator for grading. Store in a stoppered bottle and label with your name, formula of compound and yield.

Write a balanced chemical equation for the preparation and calculate the % yield based on the quantity of copper(II) chloride hydrate used.

Note that both CuCl and CuCl2 have crystalline ionic lattice structures. The structure of an ionic salt depends on the relative sizes of the ions that form the solid. The relative sizes of these ions is given by the radius ratio, which is the radius of the positive ion divided by the radius of the negative ion.

If the assumptions of an ionic model hold then the structures of the ionic crystal can be rationalized based on the radius ratio rules shown below:

Radius Ratio CN Holes in Which Cations Pack
0.155 - 0.225 3 triangular holes
0.225 - 0.414 4 tetrahedral holes
0.414 - 0.732 6 octahedral holes
0.732 - 1 8 cubic holes
1 12 close-packed structure

Note: The lower limit in each range is the ideal ratio and is known as the LIMITING RADIUS RATIO. The text-book and this link to the Univ of Sydney might be useful for additional information on the calculation of the radius ratios.

Below 400 °C CuCl has the zinc blende structure (ccp) but above this temperature it converts to the wurtzite structure (hcp). CuCl2 has a CdCl2 type structure (ccp).

Copyright © 2009-2013 by The Department of Chemistry UWI, Jamaica, all rights reserved.

Created and maintained by Prof. Robert J. Lancashire,
The Department of Chemistry, University of the West Indies,
Mona Campus, Kingston 7, Jamaica.
Created Nov 2009. Links checked and/or last modified 30th October 2014.
URL http://wwwchem.uwimona.edu.jm/lab_manuals/c1901exp35.html