Aqueous and Precipitation Reactions

Outline

Characteristics of Reactions in Aqueous Solutions
Practice with Concepts
Discussion Questions
Optional Website Reading

Characteristics of Reactions in Aqueous Solutions

Reactions in water are extremely important in living cells, electrical cells, and geological processes, so they are the focus of much chemical investigation. Most reactions depend on the extent to which the solutes compounds dissolved in the water solvent can dissociate and interact.

Solubility and Precipitation

Precipitation reactions involve ionic compounds which have completely dissociated in aqueous solution. To put it bluntly, they come apart (dissociate) in water:

NaCl → Na⁺ + Cl^-

When table salt dissolves in water at a specific concentration, there are only ions present. No solid salt is left. NaCl is soluble in water.

It is possible to put so much solute in a solution that even if the compound is considered soluble, there will still be solid compound present because no more can dissolve in the solution. For this reason, the solubility of all compounds is determined at a standard concentration (.1 moles of solute/liter of solution, or .1M, where M stands for Molar: we will go into detail about this next week). If no solid occurs at this concentration, the compound is considered soluble.

Some compounds form solids no matter what the solute/solution ratio is. Consider two ionic compounds, AB and CD, which both dissociate completely in water:

AB + H₂O → A⁺ + B^- + H₂O

CD+ H₂O → C⁺ + D^- + H₂O

Here is the question: if we combine the solutions, will AD or BC form a solid compound? We already know AB and CD will not (we have the dissociated ion solutions to prove it).

Long experience with various compounds in the lab give us a set of rules: certain ions (monatomic and polyatomic) always form precipitates when combined with specific ions; others never form compounds. The following table spells out the rules:

Always Soluble	Exceptions	Never Soluble	Exceptions
Salts of Na+, K+, NH₄+ Nitrates: NO₃^- Chlorates: ClO₃^-, ClO₄^- Acetates: CH₃CO₂^-	None	Salts of: carbonates CO₃^- phosphates PO₄^3- oxalates C₂O₄^2- chromates CrO₄^2- sulfides S^2-	Group I ions and NH₄⁺
Salts of Cl^-, Br^-, I^-	with Ag+, Hg₂²⁺, Pb²⁺	Metal hydroxides OH^- Metal oxides	Group I ions, plus Sr²⁺, Ba²⁺
Salts of F^-	with Mg²⁺, Ca²⁺, Sr²⁺, Ba²⁺, Pb²⁺
Salts of SO₄^2-	with Ca²⁺, Sr²⁺, Ba²⁺, Pb²⁺, Ag⁺

Predicting Precipitation

If we combine a solution of Ba(NO₃)₂ dissolved in water with a solution of Na₂CO₃, will a precipitate form? We need to look at the dissociation equations for each ion:

Ba(NO₃)₂ → Ba²⁺ + 2(NO₃)^-

Na₂CO₃ → 2Na⁺ + CO₃^2-

Now look at the possible compounds we could have: BaCO₃ and NaNO₃. From the chart, we find NO₃ never combines to form a solid compound, so sodium nitrate is not a possible product. CO₃ always forms a solid compound except with Group I elements, and barium is a Group II element, so barium carbonate will form and precipitate (fall out as a solid).

When we have a precipitation reaction, we write the reaction formula so that it includes only those reactants involved in making the solid products. For our barium carbonate reaction, the formula would be:

Ba²⁺ + CO₃^2- → BaCO₃

This reaction is the net ionic equation, since it includes only those ions which actually participate in the reaction, even though Na⁺ and 2(NO₃)^- are present.

In most cases, it really is that simple to predict precipitation: you write the ion dissociation equation, look at the possible combinations of ions, then check the chart to see which combinations will precipitate.

Types of Aqueous Reactions

There are three types of reactions that take place in aqueous (water) solutions):

Precipitation reactions, in which two solutions containing dissociated ions combine and produce a solid compound
Acid-Base reactions, in which hydrogen-donating compounds react with hydroxide-donating compounds
Oxidation-reduction reactions, in which electrons are transferred from one compound component to another.

We will consider each of these in turn through the rest of this chapter.

Net Ionic Equations

Let's revisit a precipitation reaction to make sure we know what is going on.

When two soluble compounds are dumped into water, they will both dissociate. The resulting ions may, however, react with each other to form a precipitate. Net ionic equations express this precipitation reaction ONLY, and ignore any uninvolved ions.

Consider the situation when iron(III) chloride and sodium hydroxide are both dissolved in water. A red precipitate of iron hydroxide forms. The overall reaction can be written:

FeCl₃ + 3 NaOH -→ 3 NaCl + Fe(OH)₃

On closer inspection, however, we realize that the NaCl in water is soluble, while the Fe(OH)₃ is not. So our reaction becomes

FeCl₃(aq) + 3 NaOH(aq) -→ 3 NaCl (aq) + Fe(OH)₃(s)

Obviously, we need to look at the individual dissociations to find out just what ions are formed.

FeCl₃(aq) → Fe⁺³ (aq)+ 3 Cl^- (aq)

NaOH(aq) → Na⁺ (aq)+ OH^-(aq)

Since the precipitation reaction occurs only between the iron and the oxide ions, we can write the combined reaction:

Fe⁺³(aq) + 3Cl^- (aq) + Na⁺(aq) + OH^- (aq) → Cl^- (aq) + Na⁺ (aq) + Fe(OH)₃ (s)

Only the red ions actually participate in the formation of the precipitate. The blue ions are spectator ions in the precipitation reaction. So the net reaction involves only the red components:

Fe⁺³ + OH^- (aq) → Fe(OH ₃ (s)

It is certainly possible for a reaction in aqueous solution to produce multiple precipitates. Combining iron(III) sulfate with barium hydroxide results in both an iron hydroxide precipitate and a barium sulfate precipitate. There overall equation is:

Fe₂(SO₄)₃ (aq) + 3 BaOH₂(aq) → FeOH₃(s) + 3 BaSO₄(s)

There are two net ionic equations, one for each precipitate:

Fe⁺³ (aq) + 3 OH^- (aq)→ Fe(OH)₃ (s)

Ba⁺² (aq) + SO₄^-2 (aq) → Ba SO₄ (s)

Notice that the net ionic equations are balanced separately from the overall equation. From the standpoint of each precipitation reaction, the other ions are spectator ions. That is, as far as the precipitation of the iron hydroxide is concerned, the barium and sulfate are spectator ions....although in this case, the spectators aren't watching the "performers" because they are busy doing their own thing.

Practice with the Concepts

Discussion Questions

Why do we care about solubility of compounds?
How does concentration affect solubility?

Optional Readings

Check out Boundless Chemistry's Precipitation Reactions web page for more help with these concepts.

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Chemistry

Chapter 3: 4-5