CHAPTER 10DOUBLE DISPLACEMENT REACTIONS
Background
Double displacement reactions are among the most common simple
chemical reactions and comparatively easy to study. In each part of this experiment two water solutions, each
containing positive and negative ions, will be mixed in a test tube. Consider
the hypothetical reaction:
AB (aq) + CD (aq) ¨ AD (?) + CB (?)
Reaction 1
where AB (aq) exists as A+ (aq)
and B- (aq) ions in solution and CD (aq)
exists as C+ (aq) and D– (aq)
ions in solution. As the ions come
in contact with each other, there are six possible combinations that might
conceivably cause a chemical reaction. Two of these combinations are the
meetings of ions of like charge; that is A+ (aq) + C+
(aq) and B-
(aq) + D– (aq). Since like charges repel, no
reaction will occur. Two other possible combinations are those of the original
two compounds; that is, A+ (aq) + B–
(aq) and C+ (aq) + D– (aq).
Since we originally had a solution containing each of these pairs of ions, they
can mutually exist in the same solution; therefore they do not recombine. Thus the two possibilities for the
combination of each of the positive ions with the negative ion of the other
compound; that is, A+ (aq) + D–
(aq) and C+ (aq) + B– (aq). Let us look at some examples.
In order for a double displacement
reaction to occur one of the following must happen.
1.
A precipitate is formed.
2.
A gas is produced.
3.
Water or another
slightly ionized substance is formed.
The previous example, example 1, met none of
the criteria. While in the
following example, example 2, the formation of the insoluble salt, silver
chloride (AgCl), satisfy criteria 1.
Example 3 meets criteria 2, a gas is
evolved.
Finally, example 4, the reaction of an acid
and a base is really just another form of a double replacement reaction. It is
also called a neutralization reaction. All acid-base reactions result in the
formation of a salt and water.
Figure 1 -
Solubility Rules
Table 1 - Slightly Ionized Substances
|
HC2H3O2 (acetic acid) |
|
H2SO3 (sulfurous acid) |
|
HF (hydrofluoric acid) |
|
H2C2O4 (oxalic acid) |
|
H3PO4 (phosphoric acid) |
|
NH3 or NH4OH (ammonia) |
|
CH3NH2 (methyl amine) |
Procedure
Each
part of this experiment consists of mixing equal volumes of two solutions in a
test tube. Use approximately 3 mL of each solution. There is no need to measure
this quantity in a graduated cylinder. The important part is to use
equal volumes of solution. To determine where the 3 mL mark is on you test
tube, fill your 10 mL graduated cylinder with 3 mL of water. Pour this water
into the test tube and make a mental note where the fluid level is. This will
serve as your 3 mL mark.
Look
for evidence of chemical reaction. This may be the formation of a precipitate,
the formation of a gas, or the evolution of heat. Make sure that you give the
reaction mixtures sufficient time to react.
Formation of
a precipitate — Look for the
formation of an insoluble compound, the solution may appear cloudy.
Formation of
a gas — Look for the formation
of NH4OH
(NH3),
H2CO3, or H2SO3.
These
compounds decompose into gases. You will see bubbles, or the solution
effervesce.
Formation of
a slightly-ionized substance — Heat
usually accompanies the formation of H2O, HC2H3O2, or any other slightly-ionized compound.
In
each instance where a reaction occurred, write the complete, balanced reaction.
Where there is no evidence of reaction write ̉No reactionÓ.
Table 2: Solutions
|
1 Mix 0.1M NaCl and 0.1M KNO3 solutions. |
|
2 Mix 0.1M NaCl and 0.1M AgNO3 solutions. |
|
3 Mix 0.1M Na2CO3 and 6M HCl
solutions |
|
4 Mix 10% NaOH and dilute (6M) HCl solutions. |
|
5 Mix 0.1M BaCl2 and dilute(3M) H2SO4 solutions. |
|
6 Mix dilute (6M) NH4OH and dilute (3M) H2SO4 solutions. |
|
7 Mix 0.1M CuSO4 and 0.1M Zn(NO3)2 solutions. |
|
8 Mix 0.1M Na2CO3 and 0.1M CaCl2 |
|
9 Mix 0.1M CuSO4 and 0.1M NH4Cl solutions. |
|
10 Mix 10% NaOH and dilute
(6M) HNO3 solutions. |
|
11 Mix 0.1M FeCl3 and dilute (6M) NH4OH solutions. |
|
12 Do this part under the
hood. Add 1g of solid Na2SO3 to 3mL of
water and shake to dissolve. Add approximately 1mL of concentrated (12M) HCl solution, drop wise. |
Data and Observations
Interpretation of observations: In the space provided describe your
observations. If there is evidence
of reaction occurring give the balanced conventional, total and net ionic
equations; otherwise write NO RXN.
Make sure you include all
phases and charges where necessary.
Table 3: Reaction Data
|
1 Observation NaCl (aq) & KNO3 (aq) |
|
Conventional |
|
Total Ionic |
|
Net Ionic |
|
2 Observation NaCl (aq) & AgNO3 (aq) |
|
Conventional |
|
Total Ionic |
|
Net Ionic |
|
3 Observation Na2CO3(aq) & HCl (aq) |
|
Conventional |
|
Total Ionic |
|
Net Ionic |
|
4 Observation NaOH (aq)
& HCl (aq) |
|
Conventional |
|
Total Ionic |
|
Net Ionic |
|
5 Observation BaCl2
(aq) & H2SO4 (aq) |
|
Conventional |
|
Total Ionic |
|
Net Ionic |
|
6 Observation NH3
(aq) & H2SO4
(aq) |
|
Conventional |
|
Total Ionic |
|
Net Ionic |
|
7 Observation CuSO4 (aq) & Zn(NO3)2
(aq) |
|
Conventional |
|
Total Ionic |
|
Net Ionic |
|
8 Observation Na2CO3
(aq) & CaCl2 (aq) |
|
Conventional |
|
Total Ionic |
|
Net Ionic |
|
9 Observation CuSO4
(aq) & NH4Cl (aq) |
|
Conventional |
|
Total Ionic |
|
Net Ionic |
|
10 Observation NaOH (aq)
& HNO3 (aq) |
|
Conventional |
|
Total Ionic |
|
Net Ionic |
|
11 Observation FeCl3
(aq) & NH3 (aq) |
|
Conventional |
|
Total Ionic |
|
Net Ionic |
|
12 Observation Na2SO3
(aq) & HCl (aq) |
|
Conventional |
|
Total Ionic |
|
Net Ionic |
Post Laboratory Questions & Problems
1. Give three examples which
provide evidence for a double displacement reaction.
2. Write the equation for the
decomposition of sulfurous acid.
3. Using the three criteria
for double displacement reactions, and the solubility table,
predict whether a double displacement reaction will
occur for each of the following. If
the reaction will occur, balance the reaction and
include the proper phase labels. If no
reaction will occur, write ̉no reactionÓ. All
reactants are aqueous solutions.
(a) Potassium sulfide and copper(II) sulfate →
(b) Ammonia and oxalic acid ˆ
(c) Potassium hydroxide and ammonium chloride ˆ
(d) Sodium acetate and hydrochloric acid ˆ
(e) Sodium chromate and lead(II) acetateˆ
(f) Ammonium sulfate and sodium chloride ˆ
(g) Bismuth(III) chloride and sodium hydroxide ˆ
(h) Potassium acetate and cobalt(II) sulfate ˆ
(i) Sodium carbonate and nitric acid ˆ
(j) Zinc bromide and potassium phosphate ˆ
(k) Iron(III) chloride and ammonium nitrate ˆ
(l)
Calcium chlorate and
sodium bromide ˆ
(m) Sodium chromite and barium nitrate ˆ