bChemistryH /A
Name_________________Date__________
Lab – Double replacement reactions and Net ionic equations
Double replacement reactions involve reactions between ions in aqueous solutions. The polar nature of water makes it an
effective solvent for many ionic solids; the dissolving process is referred to as solvation. It is important to remember that the
solution remains electrically neutral because the ions’ charges cancel. The individual ions in the solution can move about
independently, and can interact with other ions that are added to the solution. Of course, if the water were to evaporate, these
ions would return to a crystalline solid. For example, in a common salt water solution there are an equal number of sodium ions
(Na+) and chloride ions (Cl-1) dispersed throughout the water molecules. If the water evaporates, the Na+ ions and Cl-1 ions will
be attracted to each other and solid NaCl will form.
In this experiment, you will mix six ionic solutions, two at a time, to determine which combinations of ions form new
compounds that are not soluble. If the “soup” of ions in the mixture contains a positive ion and a negative ion that attract each
other so strongly that water will not hold them individually dispersed, you will notice a precipitate. These insoluble products
are recognized by a cloudy or thickened appearance when two transparent solutions are mixed.
This activity will give you practice at recognizing and writing chemical formulas, balancing double replacement equations,
identifying precipitates, and writing net ionic equations. Your textbook supports these skills in Chapter 11 (in particular,
section 11.3). Refer to the section as your analyze your lab results.
Write formulas for the ions in each of the following compounds (incl. charge), and then write the formula for each
compound. Two examples are given:
sodium chloride
calcium hydroxide
Name
Na 1+
Ca 2+
Cl 1OH-
Ions present (incl. charges)
NaCl
Ca(OH)2
Formula
Silver nitrate
Sodium chromate
Aluminum sulfate
Barium chloride
Potassium chromate
Strontium nitrate
Procedure: Use clean, dry glass plates for the mixings – only a drop or two of each solution is needed. Be sure to allow
sufficient time for each combination of ions to react.
Report your observations in the data chart. For each combination of two solutions, record either “NR” (if no reaction occurs)
or a brief description of the precipitate. You will need to write a balanced equation for any combinations that form
precipitates.
Use loose-leaf paper to complete the questions that accompany the lab.
Data: Observation Chart
Silver
nitrate
Aluminum
sulfate
Barium
chloride
Potassium
chromate
Strontium
nitrate
Sodium
chromate
Silver
nitrate
Aluminum
sulfate
Barium
chloride
Potassium
Chromate
Strontium
nitrate
Sodium
chromate
Analysis: Follow these steps to write equations for each of the reactions you observed.
Q1. Write a skeleton equation for each combination of solutions that formed a precipitate. Be sure to write individual
formulas carefully – check charge, subscript agreement for each reactant and product. This is hard! You need to make it a
habit to write products thoughtfully. Then add whole-number coefficients to make a balanced equation.
Q2. Find the precipitate in the product: . Use a Solubility Table to identify which one of the two possible products in
each case is insoluble. Mark the precipitate in each equation you have written for Q1.
Honors only below this point. Text reference, 11.3.
Q3. Now write an ionic equation expand each balanced equation.. Here is an example of what to do: Consider mixing
two solutions, CaBr2 & K2CO3. This reaction forms a cloudy white precipitate. The products would be KBr and CaCO 3.
So, your balanced equation would be:
CaBr2(aq) + K2CO3(aq)
 CaCO3 (s) + 2 KBr(aq)
After checking solubility rules, you find out that CaCO3 is insoluble. Therefore, the KBr remains dissolved as
separate ions. Notice how the subscripts (aq) and (s) have been used. An ionic equation can be written for this
reaction. Write all soluble species as separate ions. Write any precipitate a “new” compound. Again, use symbols
to identify physical states of reactants and products: (aq) and (s) for solid.
Ca2+(aq) + 2Br 1- (aq) + 2K1+ (aq) + CO32-(aq)  CaCO3 (s) + 2 K 1+(aq) + 2 Br 1- (aq)
Q4. Simplify each ionic equation by leaving out any ions that appear on both sides of the equation. These are called
spectator ions – they do not undergo any change in the chemical reaction. In the reaction above, Br-1 and K+1 ions are the
spectator ions. (At this point in your work you should also reduce coefficients if possible.) The net ionic equation would
be:
Ca2+(aq) + CO32-(aq)  CaCO3 (s)
Solubility Rules: The following rules are useful for predicting solubility in water.
1.
Salts containing alkali metal cations or ammonium are soluble.
2.
Salts containing nitrate, acetate, chlorate, or perchlorate anions are soluble.
3.
Salts containing chloride, bromide or iodide are soluble. (Except for Cu +, Ag+, Hg2+2 and Pb+2)
4.
Salts containing sulfate or sulfite are soluble. (Except Ca+2, Ba+2, Sr+2 and Pb+2)
5.
Oxide and hydroxide salts are insoluble. (Except with group IA cations and ammonium.)
6.
Sulfide and hydrogen-sulfide salts are insoluble. (Except with group IA cations and ammonium.)
7.
Carbonate and chromate salts are insoluble. (Except with group IA and ammonium cations and ammonium.)