GENERAL CHEMISTRY TOPICS

Precipitation reactions

Dissociation of ionic compounds in aqueous solution. Solubility rules. Using solubility rules: Predicting when a precipitation reaction will occur. Writing molecular, complete ionic, and net ionic equations for a precipitation reaction.

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A precipitation reaction occurs upon the mixing of two solutions of ionic compounds when the ions present together in the mixture can form an insoluble compound. In such cases, the solution turns visibly cloudy, a phenomenon known as precipitation. The cloudiness is due to the formation of small aggregations of solid substance (the precipitate). The precipitate can be separated from the remaining solution by filtration.

Dissociation of ionic compounds in aqueous solution

When an ionic compound dissolves in water to form a solution, the compound dissociates into separated ions. For most purposes, we can consider this dissociation as a separation of pre-existing ions from a solid crystal lattice into individual ions that are free to move about in solution. This is why pure water does not conduct electricity well, but a solution of NaCl (for example) can. We call an ionic compound that dissociates in solution to give rise to mobile ions an electrolyte. In this view, dissociation seems more like a physical change than a chemical change, but we can still represent the process as a chemical equation. Of course, just like any other chemical equation, it must be balanced: the number of atoms of each type and the net charge on each side of the equation must be the same. So for example, the dissociation equations for the soluble ionic compounds ammonium sulfate and sodium nitrate are

(NH₄)₂SO₄(s)  →  2 NH₄⁺(aq) + SO₄²⁻(aq)

NaNO₃(s)  →  Na⁺(aq) + NO₃⁻(aq)

Notice that dissociation equations express stoichiometric relationships. Note in particular that each formula unit of (NH₄)₂SO₄ gives rise to two ammonium (NH₄⁺) ions in solution. All the other stoichiometric relationships are one-to-one. The solvent molecules interact very strongly with the ions, which are said to be solvated.

Solubility rules

A simple set of rules, known as solubility rules, allows us to predict when a precipitation reaction will occur. These vary in detail, according to different sources, but broadly speaking the rules provide a quick qualitative check on the solubility of combinations of ions in water. The solubility rules presented below are from Oxtoby, et al. 4ed, p.149.

Table 1: Solubility of ionic compounds in aqueous solution

	Anion	Soluble	Slightly soluble		Insoluble
	nitrate, NO3–	All	—		—
	acetate, CH3COO–	Most	—		Be(CH3COO)2
	chlorate, ClO3–	All	—		—
	perchlorate, ClO4–	Most	KClO4		—
	fluoride, F –	Group I, AgF, BeF2	SrF2, BaF2, PbF2		MgF2, CaF2
	chloride, Cl –	Most	PbCl2		AgCl, Hg2Cl2
	bromide, Br –	Most	PbBr2 , HgBr2		AgBr, Hg2Br2,
	iodide, I –	Most	—		AgI, Hg2I2, PbI2, HgI2
	sulfate, SO42–	Most	CaSO4, Ag2SO4, Hg2SO4		SrSO4,BaSO4, PbSO4
	sulfide, S2–	Groups I and II, (NH4)2S	—		Most
	carbonate, CO32–	Group I, (NH4)2CO3	—		Most
	sulfite, SO32–	Group I, (NH4)2CO3	—		Most
	phosphate, PO43–	Group I, (NH4)3PO4	Li3PO4		Most
	hydroxide, OH –	Group I, Ba(OH)2	Sr(OH)2, Ca(OH)2		Most
	Soluble: solubility > 1 g / 100 g water ; Slightly soluble: 1 g / 100 g water ≥ solubility ≥ 0.01 g / 100 g water ; Insoluble: solubility < 0.01 g / 100 g water.

 

Using solubility rules

The above table (or one similar to it) can be used to predict whether precipitation will occur when two solutions of different soluble ionic compounds are mixed.

Example 1. For the two soluble compounds above, ammonium sulfate and sodium nitrate, will a precipitation occur upon mixing their aqueous solutions?

The combinations we must check are ammonium nitrate and sodium sulfate. All nitrates and most sulfates are soluble, and sodium sulfate is not listed among the insoluble or slightly soluble sulfates. No precipitation reaction should occur.

Example 2. Solutions of lead acetate and sodium iodide are mixed. Will a precipitation reaction occur?

The answer is yes. Lead iodide is insoluble and will precipitate. Lead iodide actually forms a bright yellow solid. Sodium acetate is soluble and remains in solution.

Writing net ionic equations

When a precipitation reaction occurs, a chemical equation for the precipitation reaction that involves only the ionic species that react (the reactants) and the ionic compound that forms the precipitate (the products) is known as a net ionic equation. The ion types that do not undergo reaction, and remain in solution after precipitation, are known as spectator ions. Spectator ions are unchanged in the precipitation reaction, and therefore do not appear in a net ionic equation.

What follows is a procedure for writing net ionic equations, with an example.

(1) Write the dissolution equation for each ionic compound in solution.

Example: Suppose we want to write chemical equations correctly describing what will occur when we mix solutions of ammonium sulfate and barium chloride.

(NH₄)₂SO₄(s)  →  2 NH₄⁺(aq) + SO₄²⁻(aq)

BaCl₂(s)  →  Ba²⁺(aq) + 2 Cl⁻(aq)

are the two dissolution equations. The products of both these equations should all be written as reactants in a new equation:

2NH₄⁺(aq)  +  SO₄²⁻(aq)  +  Ba²⁺(aq)  +  2 Cl⁻(aq)  → 

We next fill in the products.

(2) Check the solubility rules for insoluble combinations. If all combinations are soluble, there is no reaction, and thus no net ionic equation to write. For an insoluble combination, write the formula for that compound as a product (solid state). The other ions remain in solution, and should also be written on the product side, but as aqueous ions.

In this example, we know ammonium sulfate and barium chloride are soluble. If we we switch partners, we would have barium sulfate and ammonium chloride. Only a few chlorides are insoluble according to our table, and ammonium chloride is not among them. Barium sulfate is insoluble, so we would have a precipitation reaction upon mixing solutions of ammonium sulfate and barium chloride

2 NH₄⁺(aq)  +  SO₄²⁻(aq)  +  Ba²⁺(aq)  +  2 Cl⁻(aq)  →  BaSO₄(s)  +  2 NH₄⁺(aq)  +  2 Cl⁻(aq)

(3) This is the point to make sure the equation is balanced. Be sure you have the correct formula for the precipitating ionic solid (there should be no net charge for the ion combination), and that the component ions balance on the product and reactant side. Note that even if you have written correct, balanced equations for the dissociation reactions, the equation at this point is not guaranteed to be balanced, although in our example the equation happens to be balanced.

(4) Cancel ions appearing on both sides of the equation. These are not reacting, and are therefore spectator ions.

Ba²⁺(aq)  +  SO₄²⁻(aq)   →   BaSO₄(s)  

This is the net ionic equation for the example precipitation reaction. By convention the cation is usually written first.