A precipitation reaction is a reaction where aqueous solutions of soluble ionic compounds (salts), when mixed, produce an insoluble product called a precipitate. In this guide, we will learn how to predict the products of a precipitation reaction. Other names for precipitation reactions include:
1. double displacement reactions
2. double replacement reactions
3. metathesis reactions.
Before we discuss precipitation reactions, we will talk about solubility of ionic compounds in water. Many ionic compounds are soluble in water and others are not soluble in water. For example, 36 grams of sodium chloride can be dissolved in 100 grams of water at 20oC. The solubility of sodium chloride is 36.0 g NaCl/100 g H2O. Sodium chloride is readily soluble in water. The solubility of AgCl, silver chloride, in water at 20oC is 0.018 g AgCl/100 mL H2O. It is obvious that silver chloride is not very soluble in water. In fact, we consider it insoluble in water. Below is a table of solubility rules for ionic compounds in water. Please become familiar with them as they are used in the lecture and lab quite often.
From the table, we see that Na3PO4 is soluble while Ca3(PO4)2 is not soluble. Both K2S and CaS are soluble, but MgS and FeS is insoluble in water. Please get familiar with the solubility rules as they are important when writing equations.
Next, we look at precipitation reactions. These are also called double replacement reactions. Two aqueous solutions, AgNO3 and KCl are mixed together (please see figure below).
Note, the AgNO3 and KCl are both dissolved in water and are present as ions. Soluble ionic compounds dissociate in water as ions. When the two solutions are mixed together, the silver, Ag+, and chloride ions, Cl–, form a precipitate of solid AgCl. You can think of the Ag+ and Cl– ions as having a greater affinity for one another than for water molecules. When the precipitate forms, there are NO3– and K+ ions dissociated in the solution. These are the spectator ions.
To write the chemical reaction, we include the reactants and the products. First we write the reactants:
Include the phase labels for the reactants. Both reactants are aqueous which means they are dissolved in water.
Next, we determine the products. The easiest way to do this is to replace the Ag in AgNO3 with the potassium from KCl. Then replace the potassium in KCl with silver.
It is important to determine if the products are aqueous or solids. KNO3 is soluble in water, therefore, it is aqueous. AgCl is insoluble in water and is the precipitate. Add the phase labels to the products.
Finally, balance the equation. This equation is balanced, and our molecular equation is:
The balanced molecular equation is our overall chemical equation. We now will write an ionic equation from our molecular equation. We see that AgNO3, KCl, and KNO3 are aqueous–they are in solution as dissolved ions. The AgCl is a precipitate (solid). The net ionic equation can be written as:
The potassium ions and nitrate ions cancel from both sides of the equation. What is left is the net ionic equation.
The net ionic equation shows only the reaction between the ions to form the precipitate. The K+ and NO3– ions are spectator ions. In fact, the net ionic equation and the spectator ions can be determined from the molecular equation. The net ionic equation can be written from the precipitate, AgCl(s), and the spectator ions from the aqueous KNO3 (aq).
Would a precipitate form if aqueous solutions of NaCl and KCH3COO were mixed together? Let’s write the molecular equation:
After checking the solubility rules, we see that NaCH3COOH and KCl are both soluble. Therefore, our molecular equation is:
All of the reactants and products are in solution as dissolved ions. The ionic equation is:
The ions all cancel out which means no reaction has occurred–NR.
Click Here for the Solubility and Precipitation Reactions Worksheet
Before attempting the exercises, you might be interested in watching the following video:
Video: Precipitation Reactions
Exercises
Exercise 1. An aqueous silver nitrate solution is mixed with an aqueous solution of potassium carbonate.
a) Write the chemical formulas for silver nitrate and potassium carbonate
b) Write the balanced molecular (overall) equation for the reaction.
c) Determine if a precipitate is formed and if so, label it.
d) Write the ionic equation
e) Write the net ionic equation
f) Name the spectator ions
Exercise 2. An aqueous solution of sodium oxalate is mixed with an aqueous solution of calcium chloride.
a) Write the balanced molecular (overall) equation for the reaction.
b) Determine if a precipitate is formed and if so, label it.
c) Write the ionic equation
d) Write the net ionic equation
e) What are the spectator ions?
Exercise 3. An aqueous solution of iron(III) chloride is mixed with an aqueous solution of lead(II) nitrate.
a) Write the balanced molecular (overall) equation for the reaction.
b) Determine if a precipitate is formed and if so, label it.
c) Write the ionic equation
d) Write the net ionic equation
e) What are the spectator ions?
Exercise 4. Write the net ionic equation for the following. Identify any spectator ions.
Fe(ClO4)2 (aq) + (NH4)3(PO4)2 (aq) → ??
Exercise 5. If aqueous solutions of KBr and ZnCrO4 are mixed, will a precipitate form? If so, write the net ionic equation, and identify any spectator ions.
Exercise 6. A chemist mixes 50.00 mL of aqueous 0.100 M K2SO4 with 45.00 mL of aqueous 0.150 M BaCl2. What is the mass, in g, of the precipitate formed?
Exercise 7. Determine if a precipitation reaction will occur if aqueous solutions of the substances in each pair are mixed. If a precipitation reaction does occur, write a net ionic equation and identify the spectator ions.
a) KOH and HClO4
b) Hg(NO3)2 and Na3PO4
c) CoCl2 and Pb(NO3)2
d) Sr(OH)2 and BaCl2
Exercise 8. A chemist mixes 150.0 mL of 0.100 M Na2SO4, 75.0 mL 0.350 M CoCl2, and 150.0 mL of 0.250 M SrS.
a) Which salts (ionic compounds) will precipitate out of solution?
b) Assuming complete precipitation of insoluble salts, what is the concentration, in M, of each ion that remains in solution?
Exercise 9. A chemist has 4 aqueous solutions, each solution containing the following pairs of ions.
Solution 1. Cl– and NO3–
Solution 2. Ca2+ and NH4+
Solution 3. Fe2+ and Ba2+
Solution 4. SO42- and CO32-
The chemist has the following salts available: Pb(NO3)2, NaOH, BaS, Na2SO4, K2C2O4, and CsNO3. For each solution, indicate the salt that would separate the ions using a precipitation reaction. Write the net ionic equation for each.
Exercise 10. Exercise 10. A chemist mixed 250.00 mL of aqueous 0.450 M Na3PO4 and with 300.00 mL of aqueous 0.355 M CaCl2. She recovered 10.39 g of a precipitate. What is the percent yield?
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thanks. very helpful.
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