Lesson 3: Precipitation Reactions

A precipitation reaction is a reaction in which soluble ions in separate solutions are mixed together to form an insoluble compound that settles out of solution as a solid. That insoluble compound is called a precipitate.

Solubility Rules for Ionic Compounds

Solubility rules are useful summaries of information about which ionic compounds (or combinations of ions) are soluble in water and which are not. They are also important tools for making predictions about whether certain ions will react with one another to form a precipitate. In addition, they are useful for figuring out what ions might be involved when a precipitation reaction has been observed.

You will be responsible for being able to use the solubility rules to achieve the tasks listed above, although you will not be responsible for memorizing the solubility rules.

In this section you will use solubility rules to predict precipitation reactions and then write equations to represent them.

Using Solubility Rules | Prediciting Precipitation Reactions | Writing Equations for Precipitation Reactions

Using Solubility Rules

Solubility rules can be presented in several ways. A table of solubility rules is shown in your workbook as example 7-a. The more traditional version of the solubility rules is shown in example 7-b of your workbook.

In the table (Ex. 7-a), the first column lists the "type of compound" and the next two columns list which of those substances are soluble and which are insoluble or slightly soluble. Under "type of compound" the phrase "nitrates" refers collectively to all compounds which contain nitrate ions. Similarly, "chlorides" refers to all compounds that contain chloride ions. And so on.

Look over both the table and the rules to make sure they make sense to you. If you want additional clarification, read over the list below.

• In the first row of the table (Ex. 7.a), it says nitrates (i.e. ionic compounds containing nitrate ions) are all soluble.
• Chlorides, it says, are all soluble except for some that are listed as copper(I), lead(II), mercury(I) and silver. Note that copper(I) chloride is insoluble but more common copper(II) chloride is soluble.
• All sulfates are soluble except those listed under the insoluble category. Remember that sulfate is SO₄^2-.
• With hydroxides the pattern turns around. The previous pattern was that the compounds generally were soluble with some exceptions. With hydroxides it is the other way around. There is a short list of those that are soluble: NaOH, KOH, NH₄OH. The rest of them are insoluble or slightly soluble (Ca(OH)₂, Sr(OH)₂, Ba(OH)₂).
• For sulfides (S^2-), again there is a short list of those that are soluble while the rest are insoluble.(Be careful not to get sulfate mixed up with sulfide!)
• Again with carbonates and phosphates we have a short list of those ions which are soluble with these polyatomic ions and all the rest are insoluble.
• Perhaps you noticed that the group 1-A metals (Na⁺ and K⁺) and ammonium ion were always listed as being soluble. So that can be made to be another rule, and it is shown on the next line.

Practice

Now use these solubility rules to decide whether or not the compounds listed here are soluble.  After you have done that, check your answers below and then continue with the lesson.

NH4OH

CuS

FeSO4

PbCl2

Ca(NO3)2

K2CO3

 

Answers

You should have found that ammonium hydroxide is soluble, copper(II) sulfide is insoluble, iron(II) sulfate is soluble, lead(II) chloride is insoluble (or slightly soluble), calcium nitrate is soluble, and potassium carbonate is soluble.

As you have just seen, solubility rules can be used to figure out whether a certain combination of ions will come apart and dissolve in water. Some will and some won't.

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Prediciting Precipitation Reactions

Solubility rules can be used to figure out whether ions that are already in solution will come together to form an insoluble compound, that is, precipitate.

Let's briefly consider two examples. Because the rule for "hydroxides" says that sodium hydroxide is soluble, sodium ions and hydroxide ions will not come together out of solution to form a solid material.	Na+(aq) + OH-(aq) no reaction
On the other hand, the rule for "chlorides" says that lead(II) chloride is insoluble. Therefore lead(II) ions and chloride ions already in solution will come together to form a solid material that we say "precipitates out of solution."	Pb2+(aq) + 2 Cl-(aq) PbCl2(s)

 

Practice

Take a moment now to determine whether the combinations of ions listed below will result in the formation of insoluble precipitates. In each case where they do, give the name and the formula of the compound that is formed. Then check your answers below before you continue with the lesson.

NH4+ and Cl-

Pb2+ and OH-

Ba2+ and PO43-

K+ and S2-

 

Answers

• Ammonium chloride is soluble.
• Lead (II) hydroxide is insoluble and has the formula Pb(OH)₂.
• Barium phosphate is insoluble and has the formula Ba₃(PO₄)₂.
• Potassium sulfide is soluble.

 

Mixing Actual Solutions

When working with actual solutions, things are not as simple as in the examples and exercises above. The complicating factor is that each ionic solution contains both positive ions and negative ions. Consequently, even in the simplest cases, there are two combinations that must be considered and checked for solubility.

For example, when barium nitrate is mixed with copper(II) sulfate, one possible combination is copper(II) nitrate and the other is barium sulfate. According to the "nitrates" rule, all nitrates, including copper(II) nitrate, are soluble. According to the "sulfates" rule, barium sulfate is insoluble. Therefore, we can predict that this combination of ions will form a precipitate and that the precipitate will be barium sulfate. We can test that prediction by mixing the solutions and seeing what happens.

That is what you have done (or will do) in exercise 9. First use the solubility rules to determine whether a precipitate will occur for each combination. That should be done now. Then, when you are in the lab, mix the solutions on a spot plate to see which combinations actually do form precipitates. If there are discrepancies between your predictions and your observations, check with the instructor in the lab to find out why.

 

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Writing Equations for Precipitation Reactions

Precipitation reactions can be represented using several types of chemical equations: complete-formula equations (also known as "molecular" equations), complete ionic equations, and net ionic equations.  Each provides a different perspective on the chemicals involved in the reaction.

An example of how to write each type of equation is given in the reading below, using one of the combinations from exercise 9 in your workbook.

Complete-Formula Equations

Let's work with these combinations of ions from exercise 9 in your workbook.	Ba2+ NO3- Cu2+ SO42-
The compound that contains both barium and nitrate ions is called barium nitrate and has the formula Ba(NO3)2. The compound that contains both copper(II) ions and sulfate ions is called copper(II) sulfate (or cupric sulfate) and has the formula CuSO4.	barium nitrate copper(II) sulfate Ba(NO3)2 CuSO4
When these compounds react with one another we get a double displacement reaction and the products are barium sulfate and copper(II) nitrate. The formulas are BaSO4 and Cu(NO3)2.	Ba(NO3)2 + CuSO4 BaSO4 + Cu(NO3)2
To balance this equation we have to make sure that we have the same number of bariums, nitrates, coppers, and sulfates on each side of the equation. We do, so the equation is balanced.
	(Please note that it is a coincidence that each formula contains two nitrate ions.  This is because both barium and copper ions have a charge of +2.)

Precipitation reactions are often represented using complete-formula equations. Their advantage is that they show the compounds that are involved. Their disadvantage is that they do not show the forms of the chemicals that are actually involved in the reaction. (Sometimes these equations are called "molecular" equations, even though the compounds are ionic rather than molecular.)

Practice: Now please do exercise 9 in your workbook and check your answers below before continuing with the lesson.

Answers

These are the complete-formula equations for the reactions in which precipitates are predicted in Exercise 9. The precipitates are marked as (s).

CuSO4(aq) + Ba(NO3)2(aq) Cu(NO3)2(aq) + BaSO4(s) copper(II) sulfate + barium nitrate copper(II) nitrate + barium sulfate

CuSO4(aq) + 2 AgNO3(aq) Cu(NO3)2(aq) + Ag2SO4(s) copper(II) sulfate + silver nitrate copper(II) nitrate + silver sulfate

CuSO4(aq) + 2 NaOH(aq) Cu(OH)2(s) + Na2SO4(aq) copper(II) sulfate + sodium hydroxide copper(II) hydroxide + sodium sulfate

Ba(NO3)2(aq) + 2 NaOH(aq) Ba(OH)2(s) + 2 NaNO3(aq) barium nitrate + sodium hydroxide barium hydroxide + sodium nitrate

AgNO3(aq) + NaOH (aq) AgOH(s) + NaNO3(aq) silver nitrate + sodium hydroxide silver hydroxide + sodium nitrate

AgNO3(aq) + KCl (aq) AgCl(s) + KNO3(aq) silver nitrate + potassium chloride silver chloride + potassium nitrate

 

Complete Ionic Equations

Complete ionic equations represent soluble ionic compounds as separated ions. Notice that if the "molecular" formula was labeled aqueous, that compound exists in solution as aqueous ions; the precipitate remains a solid.

In the example shown above, barium nitrate exists in solution as separate barium ions and nitrate ions.  Copper(II) sulfate exists in solution as separate copper(II) ions and sulfate ions. After the barium and sulfate ions combine to form the solid barium sulfate precipitate, the copper(II) ions and the nitrate ions remain in solution. This representation is called the complete ionic equation. The barium and sulfate ions actually react to form the barium sulfate product. The copper(II) ions and the nitrate ions are not actually involved in the reaction. They are called spectator ions.

Net Ionic Equations

If we eliminate the spectator ions from the complete ionic equation we have what is called the net ionic equation. It shows only the ions that are actually involved in the reaction.

 

Practice

Take some time now to write the net ionic equations for the reactions you have been working with. Also, Identify the spectator ions in these reactions. Check your answers below before continuing with the lesson. (This is Ex. 15b.)

Answers

These are the net-ionic equations for the reactions in which precipitates are predicted in Exercise 9.

Ba²⁺(aq) + SO₄^2-(aq)   BaSO₄        copper(II) and nitrate were the spectator ions

2 Ag⁺(aq) + SO₄^2-(aq)   Ag₂SO₄      copper(II) and nitrate were the spectator ions

Cu²⁺(aq) + 2 OH^-(aq) Cu(OH)₂     sodium and sulfate were the spectator ions

Ba²⁺(aq) + 2 OH^-(aq) Ba(OH)₂      sodium and nitrate were the spectator ions

Ag⁺(aq) + OH^-(aq) AgOH      sodium and nitrate were the spectator ions

Ag^+(aq) + Cl^-(aq) AgCl       potassium and nitrate were the spectator ions

 

 

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