Cobalt(II) Chloride Equilibrium

This is part of the HSC Chemistry course under the topic Factors that Affect Equilibrium

HSC Chemistry Syllabus

  • Conduct practical investigations to analyse the reversibility of chemical reactions, for example:

– cobalt (II) chloride hydrated and dehydrate

  • Investigate the effects of temperature, concentration, volume and/or pressure on a system at equilibrium and explain how Le Chatelier’s principle can be used to predict such effects, for example:

– heating cobalt (II) chloride hydrate

    How does the Cobalt(II) Chloride equilibrium work?

    In this video, the mechanisms of the cobalt(II) chloride hydrate and dehydrate equilibrium as well as the ways that the equilibrium position can be manipulated will be explored. 

    Cobalt(II) Chloride Equilibrium

    • Cobalt (II) chloride, when dissolved in water, forms an equilibrium between a hydrated form and dehydrated form. The hydrated form involves the cobalt ion being compounded to 6 water molecules while the dehydrated form does not.

      $${{Co(H_2O)_6}^{2+}}_{(aq)} + \; 4{Cl^-}_{(aq)}\; \rightleftharpoons \;{{CoCl_{4}}^{2-}}_{(aq)} + \; 6H_2O_{(l)} \;\;\; \Delta{H}>0$$

      Figure: Two solutions of cobalt chloride containing different concentrations of hydrated and dehydrated forms. Left (pink) contains higher concentration of hydrated form while the solution on the right (blue) contains higher concentration of the dehydrated form.


      • Hydrated cobalt ion displays a pink colour whilst cobalt (II) chloride (dehydrated) displays a blue colour. This chemical equilibrium can be easily monitored by the colour change between pink and blue.
      • Pressure and volume do not affect this reaction as there are no gaseous reactants or products.



      • Adding HCl to the solution will initially dilute the solution and make the colour paler. The addition also increases [Cl-] which in turn increases the rate of forward reaction and shifts the chemical equilibrium to the right. The solution turns blue.

      $${Ag^+}_{(aq)} + {Cl^-}_{(aq)} \rightarrow AgCl_{(s)}$$


      • Adding silver nitrate (AgNO3): Ag+ forms a white AgCl precipitate with Cl- which decreases [Cl-]. The solution quickly turns cloudy white. This change increases the rate of reverse reaction and shifts the chemical equilibrium to the left. Solution turns pink. The pink colour will only be visible after the white precipitate settles to the bottom.


      • Adding distilled water (de-ionised water) increases the volume of the solution, and more importantly increases the amount of H2O in the system. This initially dilutes the solution, making the colour paler. As a result of this, a dilution effect will occur where the addition of solvent will increase the effective volume of the solution. By Le Chatelier's principle, an increase in the volume of the system will shift the chemical equilibrium to the side with more moles (the left side of the equation since H2O does not have a concentration). Solution turns light pink.



      • Heating the solution in a water bath drives the forward reaction (endothermic). This increases the concentration of hydrated form of cobalt ion. Solution turns blue.
      • Cooling the solution in ice drives the reverse reaction (exothermic). This increases the concentration of dehydrated form of cobalt ion. Solution turns pink.


      Previous section:  Practical Investigations of Le Chatelier's Principle (Cobalt (II) Chloride Equilibrium)

      Next section: Practical Investigations of Le Chatelier's Principle (Iron (III) Thiocyanate Equilibrium)