The university further disclaims all responsibility for any loss, injury, claim, liability, or damage of any kind resulting from, arising out or or any way related to (a) any errors in or omissions from this web site and the content, including but not limited to technical inaccuracies and typographical errors, or (b) your use of this web site and the information contained in this web site.the university shall not be liable for any loss, injury, claim, liability, or damage of any kind resulting from your use of the web site. The university expressly disclaims all warranties, including the warranties of merchantability, fitness for a particular purpose and non-infringement. This web site is provided on an "as is" basis. "Do not do demos unless you are an experienced chemist!" Please read the following disclaimer carefullyīy continuing to view the descriptions of the demonstrations you have agreed to the following disclaimer. Is the reaction as written in the forward direction endothermic or exothermic? Predict what effect adding chloride ions will have on the equilibrium. Predict what effect removing chloride ions will have on the equilibrium. Chemical Demonstrations: A Handbook for Teachers of Chemistry Wisconsin 1983 Vol. Ealy Chemical Demonstrations: A Sourcebook for Teachers. These different energies result in different colors absorbed by the cobalt.ġ. Teh different arrangment of ligands cause the electrons in the cobalt system to absorb different energies. The the chloride ions from a tetrahedral arrangement around the Co(II) ion. The wat er molecules form an octahedral arrangement around the Co(II) ion. When the water molecules bonded to the Cobalt ions are replaced, t he negative chloride form an IMF to the positive cobalt ions, and the cobalt appears blue.
When most of the negative species around the cobalt ion are wate r molecules, the ion absorbs light so that it appears pink. Positive ions attract negative particles such as chloride ions (C l -) and the oxygen end of water molecules (H 2O) sine the oxygen in in water has a partial negative charge.
What causes the pink and blue colors? Cobalt ions in salts are positive ions with a 2+ charge. The addition of heat favors the production of the blue complex, whereas removing heat favors the production of the pink complex to restore the lost energy. When reactants predominate, the solutio n looks pink and when products predominate, the solution looks blue. The reactants are of a different color (pink) than the products (blue). With a K eq of 1.7 x 10 -3 the equilibrium favors the reactants but some products are also present. The instructor will have to interpret and narrate the demonstration as the changing equilibrium of cobalt ion complexes in solution is difficult to follow. Is the reaction, as written left to right, endothermic or exothermic Cooling will shift the products towards the hydrated complex, which is more pink. If heat is added, the equilibrium will shift towards the cobalt chloride complex, which is blue in color. A change in temperature or concentration of the ions will shift the equilibrium. The tube placed in cold water will turn more pink. The tube placed in hot water will turn blue. Test tubes containing a pink solution of cobalt and chloride ions are placed in hot water and cold water.
Add water, however, and the equilibrium will shift back towards the pink hydrated species. For example, when hydrochloric acid is added, the added chloride ions shift the equilibrium position in favour of blue 2- ions and water. If the chloride or cobalt concentrations increase, the equilibrium will also shift towards the blue anhydrous cobalt chloride. Initially, a beaker contains a red-pink solution of cobalt (II) chloride, present as 2+ ions and chloride ions. The equilibrium equation representing the system is An equilibrium exists between a hydrated cobalt species and anhydrous cobalt chloride, both Co ions have an oxidation state of 2+.
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