Precipitation Reactions and Solubility Rules

 
This is part of Year 11 HSC Chemistry course under the topic of Chemical Reactions.

HSC Chemistry Syllabus

  • Conduct investigations to predict and identify the products of a range of reactions, for example:
– Precipitation

      Precipitation Reactions and Solubility Rules

      This video discusses precipitation reactions and the solubility rules.

       

      What is Precipitation?

      Precipitation is a chemical reaction where an insoluble ionic compound forms in solution. This usually occurs when two solutions are mixed together, resulting in a combination of ions that are insoluble in water. The solution turns cloudy (turbid), and is sometimes coloured, like for example: the formation of a yellow precipitate, lead(II) iodide

       


      Precipitation reactions follow the general formula:

      $$\text{Solution A} + \text{Solution B} \rightarrow \text{insoluble salt in Solution C}$$


      For instance:

      Solution 1: Contains `A^+` and `B^-` ions

      Solution 2: Contains `C^+` and `D^-` ions

      When solution 1 and 2 are mixed, ions `A^+` and `D^-` combine to form a solid ionic substance AD. However `C^+` and `B^-` remain dissolved and in aqueous state. 

      The reason for this can be explained by exploring solubility rules. 

      Solubility Rules:

      The solubility rules help determine which ions react to form an insoluble compound. The rules predict the outcome of precipitation reactions.

      Soluble Compounds:

      • All Group 1 metal cation (e.g. sodium, potassium) containing compounds are soluble.
      • All ammonium containing compounds are soluble.
      • All acetate (`CH_3COO^-`) and nitrate (`NO_3^-`) containing compounds are soluble.
      • Halide containing compounds are mostly soluble. Exceptions include silver, lead halides, and group 2 metal fluorides. 
      • Sulfate (`SO_4^{2-}`) containing compounds are mostly soluble. Exceptions: silver, barium, calcium, lead. 

      Sparingly Soluble and Insoluble Compounds:

      Common insoluble ionic compounds that are seen in precipitation reactions can be found in the NESA Chemistry Data Sheet. Solubility constants will be discussed in Year 12 HSC Chemistry under Module 5. 

        Precipitation Equations

        Precipitation reactions can be expressed as either neutral species, complete ionic, or net ionic equations. We will use the precipitation reaction between silver nitrate and sodium chloride to illustrate them.

        Neutral Species Equation:

        A neutral species equation provides the complete neutral chemical formulae of both reactant and products without breaking them down into individual ions. 

        $$AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)$$

         

          The combination of silver and chloride ions form an insoluble compound, silver chloride, that precipitates out of the solvent, causing the solution to turn cloudy or turbid. Sodium nitrate is a soluble compound that will remain dissolved.

          White silver chloride precipitate

          Complete Ionic Equation:

          A complete ionic equation shows all of the ions present in a solution. When writing a complete ionic equation - ionic compounds that are soluble in water (aqueous state) are broken down into their individual cations and anions. This is because the ions of an ionic compounds in aqueous state are always dissociated (separated). 

          For the same reaction:

          $$Na^+(aq) + Cl^-(aq) + Ag^+(aq) + NO_3^-(aq) \rightarrow AgCl(s) + Na^+(aq) + NO_3^-(aq)$$

           

          Notice that sodium and nitrate ions are aqueous before and after the reaction. These ions are not part of the precipitation reaction and therefore are known as spectator ions.

            Net Ionic Equations:

            The net ionic equation represents only the chemical species that are involved in the reaction. This means the spectator ions are eliminated and only chemical species which undergo chemical change are represented. 

            From complete ionic equation: 

            $$Na^+(aq) + Cl^-(aq) + Ag^+(aq) + NO_3^-(aq) \rightarrow AgCl(s) + Na^+(aq) + NO_3^-(aq)$$

            To net ionic equation:

            $$Cl^-(aq) + Ag^+(aq)  \rightarrow AgCl(s)$$

            Precipitation Reaction – Example 1

            When a solution of copper(II) sulfate is mixed with a solution of barium chloride, the following reaction occurs:

            Neutral species equation:

            $$CuSO_4(aq) + BaCl_2(aq) \rightarrow BaSO_4(s) + CuCl_2(aq)$$

            Net ionic equation:

            $$Ba^{2+}(aq) + SO_4^{2-}(aq) \rightarrow BaSO_4(s)$$

            The formation of the white barium sulfate precipitate causes the solution to turn cloudy. The presence of blue copper(II) chloride solution gives the blue background as shown.

             

            White barium sulfate solution in blue copper(II) chloride solution

              Practice Questions (answered in video):

              1. Write a balanced chemical equation demonstrating the reaction between silver nitrate and potassium chloride solutions.

              2. Write a chemical equation demonstrating what happens when a solution of barium hydroxide is added to a solution of magnesium sulfate.

              3. What solution can be added to potassium sulphate to produce a precipitate?

              Limitation of Solubility Rules

              Solubility rules provide general guidelines for predicting the solubility of ionic compounds in water, stating which salts are typically soluble and which are typically insoluble. However, these rules are simplifications and do not capture the full complexity of solubility behaviour.

              Here are some key points explaining why solubility rules are not always accurate:

              • Saturation and Supersaturation: Even salts that are considered "soluble" have a maximum concentration in solution, known as the solubility limit. When the concentration of the dissolved salt exceeds this limit, the solution becomes supersaturated, and the excess salt will precipitate out of the solution. For example, sodium chloride (table salt) is highly soluble in water, but if you keep adding salt beyond its solubility limit at a given temperature, it will start to form solid crystals and precipitate.
              • 'Insoluble' Salts and Dissolution: Salts labeled as "insoluble" typically have very low solubility, but this does not mean they cannot dissolve at all. Given a sufficient volume of solvent, even these salts will dissolve to some extent. For instance, calcium carbonate (CaCO₃) is considered insoluble in water, but a very small amount can still dissolve. If you have a large enough quantity of water, you can dissolve more of the "insoluble" salt. 
              • Formation of Complex Ions: Some salts may dissolve better than expected due to the formation of complex ions. For example, silver chloride (AgCl) is considered insoluble, but it can dissolve in the presence of ammonia due to the formation of the complex ion [Ag(NH₃)₂]⁺.

              • Temperature and Solubility: Solubility is highly dependent on temperature. For most salts, solubility increases with temperature, but there are exceptions. Solubility rules are often stated without specifying the temperature, leading to potential inaccuracies.

              Dissolution and solubility of ionic compounds is discussed in greater detail in Year 12 HSC Chemistry. Read more here

               

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