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The Fe(SCN)²⁺/ Fe(SCN)²⁺ Equilibrium

  • Page ID

    Chemical Concepts Demonstrated

    • Equilibrium
    • LeChatelier's Principle


    Four beakers are filled with DI water. HNO3, FeCl3, and NH4SCN are added to each of the beakers.
    1. FeCl3 is added to the first beaker.
    2. NH4SCN is added to the second beaker.
    3. NH4Cl is added to the third.
    4. KNO3 is added to the fourth.


    At first, the beakers have a light red color.

    1.  When the FeCl3 is added to the first beaker, the solution turns a deep red. 
    2.  When the NH4SCN is added to the second beaker, the solution also turns a deep red.
    3.  When the NH4Cl is added to the third beaker, the color of the solution lightens.
    4.  When the KNO3 is added to the fourth beaker,  the color of the solution will also lighten , but not as much as with the NH4Cl.

    Explanations (including important chemical equations)

    The initial light red color indicates the presence of the Fe(SCN)2+/Fe(SCN)2+ complexes.  The addition of more Fe3+and SCN- causes the equilibrium to shift in favor of the products and more of the complex is formed, turning the solution to a deeper red (this indicates that free SCN- and Fe3+ ions were present in the solution).  The addition of more NH4Cl shifts the equilibrium away from the complex towards the reactants to lighten the color of the solution.  This is because the Cl - ions sequester some of the Fe3+ ions.  The increased ionic strength also helps to shift the equilibrium to the left.  

    The KNO3 also lightens the color, but not as much as the the NH4Cl.  The ionic strength increase helps to lighten the color, but the lack of Cl - ions hinders the process.