4.3 Precipitation Reactions

1. Chemical Reactions Types

There are 3 major reactions some of which have sub types of chemical reactions:

All these types are considered to be in an aqueous medium where water is used as a solvent.

1. Precipitation Reactions: Also known as Double Replacement/Displacement Reactions
2. Acid – Base Reactions: Also known as Double Replacement/Displacement Reactions
3. Redox Reactions: There are several sub types for Redox Reactions:

1. 1. Single Replacement/Displacement Reactions
   2. Decomposition Reactions
   3. Combustion Reactions
   4. Combination (Synthesis) Reactions

Type 1. Precipitation Reactions:

Precipitation reactions are reactions by which the cations of one ionic compound will combine itself with another anion of another ionic compound. The intermolecular ionic forces of the anions and the cation yield a chemical compound which is settling at the bottom of the reaction as a solid that is called precipitate. This precipitate is water insoluble.

Example:

AgNO₃(aq)    +   K₂Cr₂O₇(aq)      →   Ag₂ Cr₂O₇(s)   +  KNO₃(aq)

Ag₂ Cr₂O₇(s) is a ruby red precipitate (solid) insoluble in water. The image below illustrates the Ag₂ Cr₂O₇ precipitate:

Figure 4.17 Precipitation Reaction

Reference: http://dept.harpercollege.edu/chemistry/chm/100/dgodambe/thedisk/chemrxn/r19.htm

One can predict the precipitation process by which the precipitate is produced. The solubility table is an excellent tool that predicts the formation of such precipitates.

Solubility Table:

The following are the solubility rules for common ionic solids. If there two rules appear to contradict each other, the preceding rule takes precedence.

1. Salts containing Group I elements (Li⁺, Na⁺, K⁺, Cs⁺, Rb⁺) are soluble . There are few exceptions to this rule. Salts containing the ammonium ion (NH₄⁺) are also soluble.
2. Salts containing nitrate ion (NO₃^–) are generally soluble.
3. Salts containing Cl ^–, Br ^–, or I ^– are generally soluble. Important exceptions to this rule are halide salts of Ag⁺, Pb²⁺, and (Hg₂)²⁺. Thus, AgCl, PbBr₂, and Hg₂Cl₂ are insoluble.
4. Most silver salts are insoluble. AgNO₃ and Ag(C₂H₃O₂) are common soluble salts of silver; virtually all others are insoluble.
5. Most sulfate salts are soluble. Important exceptions to this rule include CaSO₄, BaSO₄, PbSO₄, Ag₂SO₄ and SrSO₄ .
6. Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of transition metals and Al³⁺ are insoluble. Thus, Fe(OH)₃, Al(OH)₃, Co(OH)₂ are not soluble.
7. Most sulfides of transition metals are highly insoluble, including CdS, FeS, ZnS, and Ag₂S. Arsenic, antimony, bismuth, and lead sulfides are also insoluble.
8. Carbonates are frequently insoluble. Group II carbonates (CaCO₃, SrCO₃, and BaCO_​3) are insoluble, as are FeCO₃ and PbCO₃.
9. Chromates are frequently insoluble. Examples include PbCrO₄ and BaCrO₄.
10. Phosphates such as Ca₃(PO₄)₂ and Ag₃PO₄ are frequently insoluble.
11. Fluorides such as BaF₂, MgF₂, and PbF₂ are frequently insoluble.

Balancing Equations using Solubility Rule:

General approach:

1. Balance the molecular ionic chemical equation, the complete ionic equation and the net ionic equation.
2. With the help of the net ionic equation and the combination of the cations and the anions, one has to check ionic compound again the solubility table and determine whether this formed ionic compound is soluble and insoluble.

For very detailed solubility chart, one can use the periodic table solubility chart:

Net ionic equations

Ionic compounds are usually dissociated in aqueous solution; thus if we combine solutions of silver nitrate AgNO₃ and sodium chloride NaCl we are really combining four different species: the cations (positive ions) Ag+ and Na+ and the anions (negative ions) NO₃^– and Cl^–. It happens that when the ions Ag⁺ and Cl^– are brought together, they will combine to form an insoluble precipitate of silver chloride. The net equation for this reaction is

Ag⁺(aq) + Cl^–(aq)→ AgCl(s)

Note that the ions NO₃^– and Cl^– are not directly involved in this reaction; the equation expresses only the net change, which is the removal of the silver and chloride ions from the solution to form an insoluble solid.

the symbol (aq) signifies that the ions are in aqueous solution, and thus are hydrated, or attached to water molecules. The symbol (s) indicates that the substance AgCl exists as a solid. When a solid is formed in a reaction that takes place in solution, it is known as a precipitate. The formation of a precipitate is often indicated by underscoring.

Predicting the outcome when dissolved salts are mixed

From the above example involving silver chloride, it is clear that a meaningful net ionic equation can be written only if two ions combine to form an insoluble compound. In order to make this determination, it helps to know the solubility rules— which all students of chemistry were at one time required to memorize, but are nowadays usually obtained from tables such as the one shown below.

Balancing this molecular ionic chemical equation:

2 NaI(aq)     +   Pb(NO₃)₂(aq)    →    PbI₂(s)     +   2 NaNO₃(aq)

Complete ionic chemical equation:

Each aqueous ionic compound in the ionic chemical equation will be taken apart to its original ions, keeping all solids, liquids and gaseous compounds untouched:

2 NaI(aq)     +   Pb(NO₃)₂(aq)    →    PbI₂(s)     +   2 NaNO₃(aq)

The complete ionic chemical equation:

2 Na⁺(aq)  +  2 I^–(aq)    +  Pb²⁺(aq)    +   2 NO₃^–(aq)    →    PbI₂(s)   +   2 Na⁺(aq)   +   2 NO₃^–(aq)

Net ionic chemical equation

The net ionic chemical equation will consider only ions react and produce the product. It eliminates the ions that appear at the same time at both reactants’ as well as products’ sides. Such ions are called: Spectator ions. They present in the reaction but they do not participate in actual chemical reaction which ends up producing the products.

2 Na⁺(aq)  +  2 I^–(aq)    +  Pb²⁺(aq)    +   2 NO₃^–(aq)    →    PbI₂(s)   +   2 Na⁺(aq)   +   2 NO₃^–(aq)

In the above complete ionic chemical equation, the spectator ions are: Na⁺(aq) and NO₃^–(aq).

By eliminating these spectator ions from both the reactants’ and products’ sides:

Net Ionic Chemical Equation:

2 I^–(aq)    +     Pb²⁺(aq)      →     PbI₂(s)

Example:

Complete molecular ionic chemical equation:

KI(aq)     +   Pb(NO₃)₂(aq)    →   PbI₂(s)     +    KNO₃(aq)

Complete ionic chemical equation:

2 K⁺(aq)  +  2 I^–(aq)    +  Pb²⁺(aq)    +   2 NO₃^–(aq)     →    PbI₂(s)   +   2 K⁺(aq)   +   2 NO₃^–(aq)

With K⁺(aq) and NO₃^–(aq) are being the spectator ions.

Net ionic chemical equation:

2 I^–(aq)    +     Pb²⁺(aq)      →     PbI₂(s)

Problem Example 3: net ionic equations

Write net ionic equations for what happens when aqueous solutions of the following salts are combined:

a) PbCl₂ + K₂SO₄
b) K₂CO₃ + Sr(NO₃)₂
c) AlCl₃ + CaSO₄
d) Na₃PO₄ + CaCl₂

Solution: Use the solubility rules table(above) to find the insoluble combinations:

a)Pb²⁺(aq) + SO₄^2–(aq) → PbSO4(s)
b) Sr²⁺₍aq) + CO₃^2–(aq) → SrCO₃(s)
c) no net reaction
d) 3 Ca²⁺(aq) + 2 PO₄^3–(aq) → 3 Ca₃(PO₄)₂(s)
(Note the need to balance the electric charges)

Precipitation Reaction:

AgNO₃(aq)     +    BaCl₂(aq)        à  AgCl(s)      +     Ba(NO₃)₂(aq)

AgCl is a solid and called a precipitate. It settles at the bottom of the reaction container.

Now try the combinations given in the simulation and fill in the table below:

Determine which one of these combination will make a reaction (Double Displacement/Replacement Reaction also called Precipitation Reaction). Just type Reaction in case there is a reaction. Use the Solubility table and information given above.