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 ionsanions) 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 loose in solution. The ions are hydrated - surrounded by the
oppositely-charged ends of the water molecules. If the ions are more attracted to the water molecules
than each other, than the ions will remain in solution, until the maximum concentration is reached.
These compounds are 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- not
dissolve. Very little of the compound dissolves before the solution is saturated; the Ksp is less than
0.01(that is very little ions are in solution- most will remain as a precipitate- undissolved)
Double replacement reactions may occur when two ionic compounds are mixed. The metal ions
switch partners to form a new combination. 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. A gas or water molecules may also be formed by the new combination of ions in a double
replacement reaction.
Purpose:
To observe a series of double replacement reactions
To predict and write formulas
To practice writing balanced chemical equations
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
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.
Procedure:
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 a table of solubilities to identify which ion combination is the precipitate.
Or use the solubility rules
c. Add symbols to show precipitate (s).
d. All other combinations were soluble- so add (aq) to the equations where appropriate.
Example:
AgNO3(aq)
+ NaCl (aq)
----> NaNO3 (aq)
+
AgCl (s)
silver nitrate
sodium chloride
sodium nitrate silver chloride
4. Consider the reactions that formed precipitates.
a. Write ionic equations below the molecular equations from #3 for all the reactions that formed
precipitates. 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 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 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.
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. Show the
ions in each.
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.
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. 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