LAB #4- Double Replacement Reactions and Ionic Formulas
An application of differences in solubility- Ksp
Background:
Ionic compounds are formed when metals (positive ions- cations) and nonmetals (negative ions- anions) combine.
The ions combine in ratios that make the ionic compound neutral. Ionic compounds, that dissolve in an aqueous solution,
dissociate (separate) and are surrounded by the water molecules (ion-dipole attraction) as they move freely in solution.
The ions are hydrated - surrounded by the oppositely-charged ends of the water molecules. The ions will remain in
solution, until the maximum concentration is reached. When this maximum is reached, the solution is saturated and at
equilibrium. The strength of the ionic bond will determine the extent of dissolving prior to reaching saturation. If the ions
are more attracted to the water molecules than each other, they tend to dissolve to a greater extent. The calculated
equilibrium constant, Ksp, indicates the extent of dissolving. A compound is considered soluble in water with a Ksp
greater than 0.10. If the ions in the ionic compound are more attracted to each to each other than to the water molecules,
the compound will be insoluble. Very little of the compound dissolves before the solution is saturated; the Ksp is less
than 0.1 (that is very few ions are in solution- most will be in the precipitate- undissolved)
Double replacement reactions may occur when two ionic compounds are mixed. The metal ions “switch
partners” to form a new combination. This is however misleading, since the ions were in solution and thus dissociated
and attracted to water molecules rather than to each other. At least one set of ions is removed from the solution in a
double replacement reaction: either forming a solid precipitate, water molecule or a gas that leaves the solution. If one of
the resulting new compounds is insoluble in water, a precipitate forms. The solution becomes cloudy (a suspension has
formed) indicating the formation of a precipitate - solid particles may settle to the bottom or float on the surface or the
solution may remain cloudy.
Purpose:
To observe a series of double replacement reactions
To practice writing formulas and balance chemical equations
To apply Ksp to explain the formation of precipitates
Pre-Lab Questions:
1. Ionic Compounds- Identify the combinations of ions in the reactants.
a. Complete the following table for the reactants.
Name of
Compound
silver nitrate
Formula
AgNO3
Pb(NO3)2
CuSO4
MgSO4
FeCl3
KI
NaOH
Na2CO3
NaCH3COO
Cation(Metal)
Ag+
Name of
Cation
silver
Anion(Nonmetal)
NO3-
Name of Anion
nitrate
2. Select one of the reactants and draw visual representation of the ions in solution.
Procedures:
A. Obtain the colored grid titled Aqueous Reactions of Ionic Compounds. The grid shows the
combinations of solutions to be reacted. Copy this grid in your lab report.
B. Place a clear transparency over the grid to show the combinations.
C. With droppers, place 1 drop of each solution in the indicated places.
D. Record your observations of the mix on the grid in your lab report.
E. Rinse and dry the transparency.
F. Return both the colored grid and transparency.
Analysis:
3. Write equations for the reactions that formed precipitates.
Skip two lines between each equation.
a. Balance the equations to show conservation of mass.
b. Use the solubility rules below to identify the precipitate. Add symbols to show precipitate (s).
c. All other combinations were soluble- so add (aq) to the equations where appropriate.
Example:
AgNO3(aq)
silver nitrate
+
NaCl (aq)
sodium chloride

NaNO3 (aq)
+
AgCl (s)
sodium nitrate silver chloride
4. Consider the reactions that formed precipitates.
a. Write ionic equations below the molecular equations from #3. Show the ions as dissociated in
solution or together in the precipitate.
ex. Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) ----> AgCl(s) + NO3-(aq) + Na+(aq)
The ionic equation shows the ions dissociated in solution. See Holt pages 329 - 330 in the text for a
discussion of complete ionic equations.
b. In order to simplify the reaction, we can cross out the ions that appear in the same form on both
sides of the equation- the spectator ions. Without the spectator ions, the net ionic equation remains.
Write the net ionic equations for all the reactions that formed precipitates. (See the example on the
top of Holt page 330 in the textbook)
5. How does the Ksp of the soluble compounds compare to the Ksp of the insoluble precipitates? Explain
to show your understanding of dissolving, Ksp, and equilibrium. Explain why ionic compounds differ in
their solubilities in water. Use Ksp to explain the formation of the precipitate as well as the strength of
the ionic bond.
6. Select one of the reactions that formed a precipitate.
a. Draw a visual representation of each of the reactants and the products in the mix.
b. Use a collision model to explain the reaction. Include relative attractions between the ions and the
water molecules.
c. What ions were unchanged in the reaction? Explain. These ions are the spectator ions.
d. Write an equation to show the ions that have combined to form a precipitate. This is a net ionic
equation. The net ionic equation omits the spectator ion.
e. Compare the Ksp of the precipitate to the Ksp of the ionic compounds that are soluble.
7. HP. Use the value the concept of reaction quotient (Q) to explain the formation of the precipitate or
saturated solution. Your explanation should address the concentration of ions in solution immediately
upon mixing. Explain which reaction is faster immediately after mixing and how this leads to a
saturated solution. See Kotz chapter 18.
SOLUBILITY RULES
A. Soluble Salts
1. Alkali metals (column IA - Na+, K+...) and ammonium NH4+ form soluble salts
2. Nitrate form soluble salts (NO3-)
3. Halides: Chloride(Cl-), Bromide(Br-), and Iodide(I-) generally soluble
exceptions- with Pb2+, Hg22+, Cu+, Ag+
4. Sulfate generally soluble- except insoluble BaSO4, SrSO4, PbSO4
only slightly soluble- Ag2SO4, CaSO4, Hg2SO4
B. Insoluble Salts - with cations other than alkali (ex. Na+, K+) and ammonium NH4+,
1. Sulfides usually insoluble except with Mg2+, Ca2+ , Sr2+, Ba2+
2. Oxides usually insoluble except with Sr2+, Ba2+ soluble and Ca2+ only slightly
3. Hydroxides usually insoluble except withSr2+, Ba2+ and Ca2+ only slightly
4. Chromates usually insoluble except with Mg2+
5. Phosphates and Carbonates usually insoluble
Solubility rules that apply to aqueous solution:
These rules must be applied in order.
(1) All alkali metal and ammonium compounds are soluble.
(2) All acetate, perchlorate, chlorate, and nitrate compounds are soluble.
(3) Silver, lead, and mercury(I) compounds are insoluble.
(4) Chlorides, bromides, and iodides are soluble
(5) Carbonates, hydroxides, oxides, phosphates, silicates, and sulfides are insoluble.
(6) Sulfates are soluble except for calcium and barium.
Solubility rules that apply to aqueous solution:
These rules must be applied in order.
(1) All alkali metal and ammonium compounds are soluble.
(2) All acetate, perchlorate, chlorate, and nitrate compounds are soluble.
(3) Silver, lead, and mercury(I) compounds are insoluble.
(4) Chlorides, bromides, and iodides are soluble
(5) Carbonates, hydroxides, oxides, phosphates, silicates, and sulfides are insoluble.
(6) Sulfates are soluble except for calcium and barium.
Solubility rules that apply to aqueous solution:
These rules must be applied in order.
(1) All alkali metal and ammonium compounds are soluble.
(2) All acetate, perchlorate, chlorate, and nitrate compounds are soluble.
(3) Silver, lead, and mercury(I) compounds are insoluble.
(4) Chlorides, bromides, and iodides are soluble
(5) Carbonates, hydroxides, oxides, phosphates, silicates, and sulfides are insoluble.
(6) Sulfates are soluble except for calcium and barium.
Solubility rules that apply to aqueous solution:
These rules must be applied in order.
(1) All alkali metal and ammonium compounds are soluble.
(2) All acetate, perchlorate, chlorate, and nitrate compounds are soluble.
(3) Silver, lead, and mercury(I) compounds are insoluble.
(4) Chlorides, bromides, and iodides are soluble
(5) Carbonates, hydroxides, oxides, phosphates, silicates, and sulfides are insoluble.
(6) Sulfates are soluble except for calcium and barium.
Solubility rules that apply to aqueous solution:
These rules must be applied in order.
(1) All alkali metal) and ammonium compounds are soluble.
(2) All acetate, perchlorate, chlorate, and nitrate compounds are soluble.
(3) Silver, lead, and mercury(I) compounds are insoluble.
(4) Chlorides, bromides, and iodides are soluble
(5) Carbonates, hydroxides, oxides, phosphates, silicates, and sulfides are insoluble.
(6) Sulfates are soluble except for calcium and barium.