======Equilibrium of Iron (III) Thiocyanate======


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**Alternative titles:** Ferric Thiocyanate Equilibrium
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====Summary====                   
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====Procedure====  
  - Prepare the ferric thiocyanate solution by mixing equal volumes of dilute ferric chloride solution and potassium thiocyanate solution.  
  - Pour a thin layer of this reddish-brown solution into a glass petri dish.  
  - Add a small amount of iron(III) chloride solid to one area of the dish and observe the color darkening to red-brown.  
  - Add potassium thiocyanate solid to another area and observe a similar red-brown darkening.  
  - Add potassium phosphate solid to a different area, which removes Fe³⁺ ions by precipitation, shifting the equilibrium back and causing a more yellow color.  
  - Add silver nitrate solid to another area, which removes SCN⁻ ions by precipitation, also shifting the equilibrium left and producing more yellow.  
  - Compare all four effects side by side in the same dish for a striking visual display.


====Links==== 
Le Chatelier's Principle - Iron (III) Thiocyanate - Chemistry and Biochemistry Demo lab at OSU: 
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Iron(III) Thiocyanate Equilibrium - Science Ready: 
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📄 Equilibrium of Iron (III) Thiocyanate - Harvard Natural Sciences Lecture Demonstrations: [[https://sciencedemonstrations.fas.harvard.edu/presentations/equilibrium-iron-iii-thiocyanate]]\\

====Variations====
  * Warm the solution to observe temperature effects on the equilibrium position.  
  * Use a spectrophotometer to measure absorbance changes and determine the equilibrium constant.  
  * Demonstrate the reaction in a cuvette under a document camera for larger audiences.



====Safety Precautions====
  * Wear safety glasses, gloves, and a lab coat when handling chemicals.  
  * Handle silver nitrate with care; it stains skin and clothing.  
  * Work with small amounts of solid chemicals to minimize waste and exposure.  
  * Dispose of all solutions and solids in an inorganic chemical waste container.




====Questions to Consider====
  * Why does adding FeCl₃ or KSCN make the solution more red-brown? (Adding reactants shifts equilibrium to the right, forming more FeSCN²⁺ complex.)  
  * Why does adding potassium phosphate or silver nitrate make the solution more yellow? (They remove Fe³⁺ or SCN⁻, shifting equilibrium left toward reactants.)  
  * How does this experiment illustrate Le Chatelier’s principle? (The system responds to stresses such as added or removed species by shifting equilibrium to counteract the change.)  
  * How could you determine the equilibrium constant experimentally? (By measuring absorbance of FeSCN²⁺ at a known wavelength and applying Beer’s Law.)