- To give practice writing equations for metathesis reactions, including net ionic equations
- To illustrate the concept of solubility and the effect of temperature and crystallization
You will be determined according to the following:
- Pre-lab (10%)
- Lab Report Form (80%)
- Must include detailed observations for each reaction
- TA Evaluation of lab procedure (10%)
- Complete the pre-lab exercise, including the plot (due at the beginning of lab)
- Read the instructions for the lab and refresh your memory on anything that isn’t clear by reading your textbook
In molecular equations for many aqueous reactions, cations and anions appear to exchange partners. These reactions conform to the following general equation:
Equation 1:
AX+BY→AY+BXAX+BY→AY+BX size 12{"AX"+"BY" rightarrow "AY"+"BX"} {}
These reactions are known as metathesis reactions. For a metathesis reaction to lead to a net change in solution, ions must be removed from the solution. In general, three chemical processes can lead to the removal of ions from solution, comcomitantly serving as a driving force for metathesis to occur:
1. The formation of a precipitate2. The formation of a weak electrolyte or nonelectrolyte3. The formation of a gas that escapes from solution
The reaction of barium chloride with silver nitrate is a typical example:
Equation 2:
BaCl2(aq)+2AgNO3(aq)→Ba(NO3)2(aq)+2AgCl(s)BaCl2(aq)+2AgNO3(aq)→Ba(NO3)2(aq)+2AgCl(s) size 12{"BaCl" rSub { size 8{2} } \( "aq" \) +"2AgNO" rSub { size 8{3} } \( "aq" \) rightarrow "Ba" \( "NO" rSub { size 8{3} } \) rSub { size 8{2} } \( "aq" \) +"2AgCl" \( s \) } {}
This form of the equation for this reaction is referred to as the molecular equations. Since we know that the salts
BaCl2BaCl2 size 12{"BaCl" rSub { size 8{2} } } {},
AgNO3AgNO3 size 12{"AgNO" rSub { size 8{3} } } {}, and
Ba(NO3)2Ba(NO3)2 size 12{"Ba" \( "NO" rSub { size 8{3} } \) rSub { size 8{2} } } {} are strong electrolytes and are completely dissociated in solution, we can more realistically write the equation as follows:
Equation 3:
Ba2+(aq)+2Cl−(aq)+2Ag+(aq)+2NO3−(aq)→Ba2+(aq)+2NO3−(aq)+2AgCl(s)Ba2+(aq)+2Cl−(aq)+2Ag+(aq)+2NO3−(aq)→Ba2+(aq)+2NO3−(aq)+2AgCl(s) size 12{"Ba" rSup { size 8{2+{}} } \( "aq" \) +"2Cl" rSup { size 8{ - {}} } \( "aq" \) +"2Ag" rSup { size 8{+{}} } \( "aq" \) +"2NO" rSub { size 8{3} rSup { size 8{-{}} } } \( "aq" \) rightarrow "Ba" rSup { size 8{2+{}} } \( "aq" \) +"2NO" rSub { size 8{3} rSup { size 8{-{}} } } \( "aq" \) +"2AgCl" \( s \) } {}
This form, in which all ions are shown, is known as the complete ionic equation. Reaction occurs because the insoluble substance AgCl precipitates out of solution. The other product, barium nitrate, is soluble in water and remains in solution. We see that
Ba2+Ba2+ size 12{"Ba" rSup { size 8{2+{}} } } {} and
NO3−NO3− size 12{"NO" rSub { size 8{3} rSup { size 8{ - {}} } } } {} ions appear on both sides of the equation and thus do not enter into the reaction. Such ions are called spectator ions. If we eliminate or omit them from both sides, we obtain the net ionic equation:
Equation 4:
Ag+(aq)+Cl−(aq)→AgCl(s)Ag+(aq)+Cl−(aq)→AgCl(s) size 12{"Ag" rSup { size 8{+{}} } \( "aq" \) +"Cl" rSup { size 8{ - {}} } \( "aq" \) rightarrow "AgCl" \( s \) } {}
This equation focuses our attention on the salient feature of the reaction: the formation of the precipitate AgCl. It tells us that solutions of any soluble
Ag+saltAg+salt size 12{"Ag" rSup { size 8{+{}} } "salt"} {}and any soluble
Cl−saltCl−salt size 12{"Cl" rSup { size 8{ - {}} } "salt"} {}, when mixed, will form insoluble AgCl. When writing net ionic equations, remember that only strong electrolytes are written in the ionic form. Solids, gases, nonelectrolytes, and weak electrolytes are written in the molecular form. Frequently the symbol (aq) is omitted from ionic equations. The symbols (g) for gas and (s) for solid should not be omitted. Thus, Equation 4 can be written as
Equation 5:
Ag++Cl−→AgCl(s)Ag++Cl−→AgCl(s) size 12{"Ag" rSup { size 8{+{}} } +"Cl" rSup { size 8{ - {}} } rightarrow "AgCl" \( s \) } {}
Consider mixing solutions of KCl and
NaNO3NaNO3 size 12{"NaNO" rSub { size 8{3} } } {}. The ionic equation for the reaction is
Equation 6:
K+(aq)+Cl−(aq)+Na+(aq)+NO3−(aq)→K+(aq)+NO3−(aq)+Na+(aq)+Cl−(aq)K+(aq)+Cl−(aq)+Na+(aq)+NO3−(aq)→K+(aq)+NO3−(aq)+Na+(aq)+Cl−(aq) size 12{K rSup { size 8{+{}} } \( "aq" \) +"Cl" rSup { size 8{ - {}} } \( "aq" \) +"Na" rSup { size 8{+{}} } \( "aq" \) +"NO" rSub { size 8{3} rSup { size 8{-{}} } } \( "aq" \) rightarrow K rSup { size 8{+{}} } \( "aq" \) +"NO" rSub { size 8{3} rSup { size 8{-{}} } } \( "aq" \) +"Na" rSup { size 8{+{}} } \( "aq" \) +"Cl" rSup { size 8{ - {}} } \( "aq" \) } {}
Because all the compounds are water-soluble and are strong electrolytes, they have been written in the ionic form. They completely dissolve in water. If we eliminate spectator ions from the equation, nothing remains. Hence, there is no reaction: Equation 7:
K+(aq)+Cl−(aq)+Na+(aq)+NO3−(aq)→no reactionK+(aq)+Cl−(aq)+Na+(aq)+NO3−(aq)→no reaction size 12{K rSup { size 8{+{}} } \( "aq" \) +"Cl" rSup { size 8{ - {}} } \( "aq" \) +"Na" rSup { size 8{+{}} } \( "aq" \) +"NO" rSub { size 8{3} rSup { size 8{-{}} } } \( "aq" \) rightarrow "no reaction"} {}
Metathesis reactions occur when a precipitate, a gas, a weak electrolyte, or a nonelectrolyte is formed. The following equations are further illustrations of such processes.
Molecular equation: Equation 8:
2HCl(aq)+Na2S(aq)→2NaCl(aq)+H2S(g)2HCl(aq)+Na2S(aq)→2NaCl(aq)+H2S(g) size 12{"2HCl" \( "aq" \) +"Na" rSub { size 8{2} } S \( "aq" \) rightarrow "2NaCl" \( "aq" \) +H rSub { size 8{2} } S \( g \) } {}
Complete ionic equation:
2H+(aq)+2Cl−(aq)+2Na+(aq)+S2−(aq)→2Na+(aq)+2Cl−(aq)+H2S(g)2H+(aq)+2Cl−(aq)+2Na+(aq)+S2−(aq)→2Na+(aq)+2Cl−(aq)+H2S(g) size 12{"2H" rSup { size 8{+{}} } \( "aq" \) +"2Cl" rSup { size 8{ - {}} } \( "aq" \) +"2Na" rSup { size 8{+{}} } \( "aq" \) +S rSup { size 8{2 - {}} } \( "aq" \) rightarrow "2Na" rSup { size 8{+{}} } \( "aq" \) +"2Cl" rSup { size 8{ - {}} } \( "aq" \) +H rSub { size 8{2} } S \( g \) } {}
Net ionic equation:
2H+(aq)+S2−(aq)→H2S(g)2H+(aq)+S2−(aq)→H2S(g) size 12{"2H" rSup { size 8{+{}} } \( "aq" \) +S rSup { size 8{2 - {}} } \( "aq" \) rightarrow H rSub { size 8{2} } S \( g \) } {}
or
2H
+
+
S
2
−
→
H
2
S
(
g
)
2H
+
+
S
2
−
→
H
2
S
(
g
)
size 12{"2H" rSup { size 8{+{}} } +S rSup { size 8{2 - {}} } rightarrow H rSub { size 8{2} } S \( g \) } {}
Molecular equation:
HNO
3
(
aq
)
+
NaOH
(
aq
)
→
H
2
O
(
l
)
+
NaNO
3
(
aq
)
HNO
3
(
aq
)
+
NaOH
(
aq
)
→
H
2
O
(
l
)
+
NaNO
3
(
aq
)
size 12{"HNO" rSub { size 8{3} } \( "aq" \) +"NaOH" \( "aq" \) rightarrow H rSub { size 8{2} } O \( l \) +"NaNO" rSub { size 8{3} } \( "aq" \) } {}
Complete ionic equation:
H
+
(
aq
)
+
NO
3
−
(
aq
)
+
Na
+
(
aq
)
+
OH
−
(
aq
)
→
H
2
O
(
l
)
+
Na
+
(
aq
)
NO
3
−
(
aq
)
H
+
(
aq
)
+
NO
3
−
(
aq
)
+
Na
+
(
aq
)
+
OH
−
(
aq
)
→
H
2
O
(
l
)
+
Na
+
(
aq
)
NO
3
−
(
aq
)
size 12{H rSup { size 8{+{}} } \( "aq" \) +"NO" rSub { size 8{3} rSup { size 8{-{}} } } \( "aq" \) +"Na" rSup { size 8{+{}} } \( "aq" \) +"OH" rSup { size 8{ - {}} } \( "aq" \) rightarrow H rSub { size 8{2} } O \( l \) +"Na" rSup { size 8{+{}} } \( "aq" \) " NO" rSub { size 8{3} rSup { size 8{-{}} } } \( "aq" \) } {}
Net ionic equation:
H
+
(
aq
)
+
OH
−
(
aq
)
→
H
2
O
(
l
)
H
+
(
aq
)
+
OH
−
(
aq
)
→
H
2
O
(
l
)
size 12{H rSup { size 8{+{}} } \( "aq" \) +"OH" rSup { size 8{ - {}} } \( "aq" \) rightarrow H rSub { size 8{2} } O \( l \) } {}
In order to decide if a reaction occurs, we need to be able to determine whether or not a precipitate, a gas, a nonelectrolyte, or a weak electrolyte will be formed. The following brief discussion is intended to aid you in this regard. Table 1 summarizes solubility rules and should be consulted while performing this experiment.
The common gases are
CO2CO2 size 12{"CO" rSub { size 8{2} } } {},
SO2SO2 size 12{"SO" rSub { size 8{2} } } {},
H2SH2S size 12{H rSub { size 8{2} } S} {}, and
NH3NH3 size 12{"NH" rSub { size 8{3} } } {}. Carbon dioxide and sulfur dioxide may be regarded as resulting form the decomposition of their corresponding weak acids, which are initially formed when carbonate and sulfite salts are treated with acid:
H
2
CO
3
(
aq
)
→
H
2
O
(
l
)
+
CO
2
(
g
)
H
2
CO
3
(
aq
)
→
H
2
O
(
l
)
+
CO
2
(
g
)
size 12{H rSub { size 8{2} } "CO" rSub { size 8{3} } \( "aq" \) rightarrow H rSub { size 8{2} } O \( l \) +"CO" rSub { size 8{2} } \( g \) } {}
and
H
2
SO
3
(
aq
)
→
H
2
O
(
l
)
+
SO
2
(
g
)
H
2
SO
3
(
aq
)
→
H
2
O
(
l
)
+
SO
2
(
g
)
size 12{H rSub { size 8{2} } "SO" rSub { size 8{3} } \( "aq" \) rightarrow H rSub { size 8{2} } O \( l \) +"SO" rSub { size 8{2} } \( g \) } {}
Ammonium salts form
NH3NH3 size 12{"NH" rSub { size 8{3} } } {} when they are treated with strong bases:
NH
4
+
(
aq
)
+
OH
−
→
NH
3
(
g
)
+
H
2
O
(
l
)
NH
4
+
(
aq
)
+
OH
−
→
NH
3
(
g
)
+
H
2
O
(
l
)
size 12{"NH" rSub { size 8{4} rSup { size 8{+{}} } } \( "aq" \) +"OH" rSup { size 8{ - {}} } rightarrow "NH" rSub { size 8{3} } \( g \) +H rSub { size 8{2} } O \( l \) } {}
Table 1
| Water-soluble salts |
|
|
Na
+
,K
+
,NH
4
+
Na
+
,K
+
,NH
4
+
size 12{"Na" rSup { size 8{+{}} } ",K" rSup { size 8{+{}} } ",NH" rSub { size 8{4} rSup { size 8{+{}} } } } {}
|
All sodium, potassium, and ammonium salts are soluble. |
|
NO
3
−
,CIO
3
−
,C
2
H
3
O
2
−
NO
3
−
,CIO
3
−
,C
2
H
3
O
2
−
size 12{"NO" rSub { size 8{3} rSup { size 8{-{}} } } ",CIO" rSub { size 8{3} rSup { size 8{-{}} } } ",C" rSub { size 8{2} } H rSub { size 8{3} } O rSup { size 8{2 - {}} } } {}
|
All nitrates, chlorates, and acetate are soluble. |
|
Cl
−
Cl
−
size 12{"Cl" rSup { size 8{ - {}} } } {}
|
All chlorides are soluble except AgCl,
Hg2Cl2Hg2Cl2 size 12{"Hg" rSub { size 8{2} } "Cl" rSub { size 8{2} } } {}, and
PbCl2∗PbCl2∗ size 12{"PbCl" rSub { size 8{2} } * } {}. |
|
Br
−
Br
−
size 12{"Br" rSup { size 8{ - {}} } } {}
|
All bromides are soluble except AgBr,
Hg2Br2Hg2Br2 size 12{"Hg" rSub { size 8{2} } "Br" rSub { size 8{2} } } {},
PbBr2∗PbBr2∗ size 12{"PbBr" rSub { size 8{2} } * } {}, and
HgBr2∗HgBr2∗ size 12{"HgBr" rSub { size 8{2} } * } {}. |
|
I
−
I
−
size 12{I rSup { size 8{ - {}} } } {}
|
All iodides are soluble except AgI,
Hg2I2Hg2I2 size 12{"Hg" rSub { size 8{2} } I rSub { size 8{2} } } {},
PbI2PbI2 size 12{"PbI" rSub { size 8{2} } } {}, and
HgI2.0HgI2.0 size 12{"HgI" rSub { size 8{2} } "." 0} {} |
|
SO
4
2
−
SO
4
2
−
size 12{"SO" rSub { size 8{4} rSup { size 8{2 - {}} } } } {}
|
All sulfates are soluble except
CaSO4∗CaSO4∗ size 12{"CaSO" rSub { size 8{4} } * } {},
SrSO4SrSO4 size 12{"SrSO" rSub { size 8{4} } } {},
BaSO4BaSO4 size 12{"BaSO" rSub { size 8{4} } } {},
Hg2SO4Hg2SO4 size 12{"Hg" rSub { size 8{2} } "SO" rSub { size 8{4} } } {},
PbSO4PbSO4 size 12{"PbSO" rSub { size 8{4} } } {}, and
Ag2SO4Ag2SO4 size 12{"Ag" rSub { size 8{2} } "SO" rSub { size 8{4} } } {}. |
Table 2
| Water-insoluble salts |
|
| CO32−CO32− size 12{"CO" rSub { size 8{3} rSup { size 8{2 - {}} } } } {},
SO32−SO32− size 12{"SO" rSub { size 8{3} rSup { size 8{2 - {}} } } } {},
PO43−PO43− size 12{"PO" rSub { size 8{4} rSup { size 8{3 - {}} } } } {} |
All carbonates, sulfites, phosphates, and chromates are insoluble except those of
CrO42−CrO42− size 12{"CrO" rSub { size 8{4} rSup { size 8{2 - {}} } } } {} alkali metals and
NH4+NH4+ size 12{"NH" rSub { size 8{4} rSup { size 8{+{}} } } } {}. |
|
OH
−
OH
−
size 12{"OH" rSup { size 8{ - {}} } } {}
|
All hydroxides are insoluble except those of alkali metals and
Ca(OH)2∗Ca(OH)2∗ size 12{"Ca" \( "OH" \) rSub { size 8{2} } * } {},
Sr(OH)2∗Sr(OH)2∗ size 12{"Sr" \( "OH" \) rSub { size 8{2} } * } {}, and
Ba(OH)2Ba(OH)2 size 12{"Ba" \( "OH" \) rSub { size 8{2} } } {}. |
|
S
2
−
S
2
−
size 12{S rSup { size 8{2 - {}} } } {}
|
All sulfides are insoluble except those of the alkali metals, alkaline earths, and
NH4+NH4+ size 12{"NH" rSub { size 8{4} rSup { size 8{+{}} } } } {}. |
*Slightly soluble.
Table 3
| Salts |
All common soluble salts |
| Acids |
HClO4HClO4 size 12{"HClO" rSub { size 8{4} } } {}, HCl, HBr, HI,
HNO3HNO3 size 12{"HNO" rSub { size 8{3} } } {}, and
H2SO4H2SO4 size 12{H rSub { size 8{2} } "SO" rSub { size 8{4} } } {} are strong electrolytes; all others are weak. |
| Bases |
Alkali metal hydroxides,
Ca(OH)2Ca(OH)2 size 12{"Ca" \( "OH" \) rSub { size 8{2} } } {},
Sr(OH)2Sr(OH)2 size 12{"Sr" \( "OH" \) rSub { size 8{2} } } {}, and
Ba(OH)2Ba(OH)2 size 12{"Ba" \( "OH" \) rSub { size 8{2} } } {} are strong electrolytes; all others are weak. |
Which are the weak electrolytes? The easiest way of answering this question is to identify all of the strong electrolytes, and if the substance does not fall in that category then it is a weak electrolyte. Note, water is a nonelectrolyte. Strong electrolytes are summarized in Table.2.
In the first part of this experiment, you will study some metathesis reactions. In some instances it will be very evident that a reaction has occurred, whereas in others it will not be so apparent. In the doubtful case, use the guidelines above to decide whether or not a reaction has taken place. You will be given the names of the compounds to use but not their formulas. This is being done deliberately to give practice in writing formulas from names.
In the second part of this experiment, you will study the effect of temperature on solubility. The effect that temperature has on solubility varies from salt to salt. We conclude that mixing solutions of KCl and
NaNO3NaNO3 size 12{"NaNO" rSub { size 8{3} } } {} resulted in no reaction (see Equations 6 and 7). What would happen if we cooled such a mixture? The solution would eventually become saturated with respect to one of the salts, and crystals of that salt would begin to appear as its solubility was exceeded. Examination of Equation 6 reveals that crystals of any of the following salts could appear initially:
KNO3KNO3 size 12{"KNO" rSub { size 8{3} } } {}, KCl,
NaNO3NaNO3 size 12{"NaNO" rSub { size 8{3} } } {}, or NaCl.Consequently, if a solution containing
Na+Na+ size 12{"Na" rSup { size 8{+{}} } } {},
K+K+ size 12{K rSup { size 8{+{}} } } {},
Cl−Cl− size 12{"Cl" rSup { size 8{ - {}} } } {}, and
NO3−NO3− size 12{"NO" rSub { size 8{3} rSup { size 8{ - {}} } } } {} ions is evaporated at a given temperature, the solution becomes more and more concentrated and will eventually become saturated with respect to one of the four compounds. If a evaporation is continued, that compound will crystallize out, removing its' ions from solution. The other ions will remain in solution and increase in concentration. Before beginning this laboratory exercise you are to plot a graph of the solubilities of the four salts given in Table 3 on your report sheet.
Hopefully here for the Pre-Lab
Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
This assignment must be completed individually and turned in to your TA at the beginning of lab. You will not be allowed to begin the lab until you have completed this assignment.
1. Write molecular, complete ionic, and net ionic equations for the reactions that occur, if any, when solutions of the following substances are mixed: (a) nitric acid and barium carbonate
(b) zinc chloride and lead nitrate
(c) acetic acid and sodium hydroxide
(d) calcium nitrate and sodium carbonate
(e) ammonium chloride and potassium hydroxide
2. Which of the following are not water soluble:
Ba(NO3)2Ba(NO3)2 size 12{"Ba" \( "NO" rSub { size 8{3} } \) rSub { size 8{2} } } {},
FeCl3FeCl3 size 12{"FeCl" rSub { size 8{3} } } {},
CuCO3CuCO3 size 12{"CuCO" rSub { size 8{3} } } {},
CuSO4CuSO4 size 12{"CuSO" rSub { size 8{4} } } {}, ZnS,
ZnSO4ZnSO4 size 12{"ZnSO" rSub { size 8{4} } } {}?
3. Write equations for the decomposition of
H2CO3H2CO3 size 12{H rSub { size 8{2} } "CO" rSub { size 8{3} } } {} and
H2SO3H2SO3 size 12{H rSub { size 8{2} } "SO" rSub { size 8{3} } } {}.
4. At what temperature (from your graph) do
KNO3KNO3 size 12{"KNO" rSub { size 8{3} } } {} and NaCl have the same molar solubility?
5. Which of the following are strong electrolytes:
BaCl2BaCl2 size 12{"BaCl" rSub { size 8{2} } } {},
AgNO3AgNO3 size 12{"AgNO" rSub { size 8{3} } } {}, HCl,
HNO3HNO3 size 12{"HNO" rSub { size 8{3} } } {},
HC2H3O2HC2H3O2 size 12{"HC" rSub { size 8{2} } H rSub { size 8{3} } O rSub { size 8{2} } } {}?
6. Which of the following are weak electrolytes:
HNO3HNO3 size 12{"HNO" rSub { size 8{3} } } {}, HF, HCl,
NH3(aq)NH3(aq) size 12{"NH" rSub { size 8{3} } \( "aq" \) } {}, NaOH?
7. For each of the following water-soluble compounds, indicate the ions present in an aqueous solution: NaI,
K2SO4K2SO4 size 12{K rSub { size 8{2} } "SO" rSub { size 8{4} } } {}, NaCN,
Ba(OH)2Ba(OH)2 size 12{"Ba" \( "OH" \) rSub { size 8{2} } } {},
(NH4)2SO4(NH4)2SO4 size 12{ \( "NH" rSub { size 8{4} } \) rSub { size 8{2} } "SO" rSub { size 8{4} } } {}.
Hopefully here for the Report Form
Note: In preparing this report you are free to use references and consult with others. However, you may not copy from other students’ work (including your laboratory partner) or misrepresent your own data (see honor code).
Name: ___________________________________________
Section: __________________________________________
1. Copper (II) sulfate + sodium carbonate
Table 4
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
2. Copper(II) sulfate + barium chloride
Table 5
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
3. Copper(II) sulfate + sodium phosphate
Table 6
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
4. Sodium carbonate + sulfuric acid
Table 7
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
5. Sodium carbonate + hydrochloric acid
Table 8
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
6. Cadmium chloride + sodium sulfide
Table 9
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
7. Cadmium chloride + sodium hydroxide
Table 10
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
8. Nickel chloride + silver nitrate
Table 11
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
9. Nickel chloride + sodium carbonate
Table 12
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
10. Hydrochloric acid + sodium hydroxide
Table 13
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
11. Ammonium chloride + sodium hydroxide
Table 14
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
12. Sodium acetate + hydrochloric acid
Table 15
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
13. Sodium sulfide + hydrochloric acid
Table 16
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
14. Lead nitrate + sodium sulfide
Table 17
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
15. Lead nitrate + sulfuric acid
Table 18
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
16. Potassium chloride + sodium nitrate
Table 19
| Observations |
|
| Molecular equation |
|
| Complete ionic equation |
|
| Net ionic equation |
|
Table 3 Molar Solubilities of NaCl,
NaNO3NaNO3 size 12{"NaNO" rSub { size 8{3} } } {}, KCl, and
KNO3KNO3 size 12{"KNO" rSub { size 8{3} } } {} (mol/L)
Table 20
| Compound |
0°C |
20°C |
40°C |
60°C |
80°C |
100°C |
| NaCl |
5.4 |
5.4 |
5.5 |
5.5 |
5.5 |
5.6 |
|
NaNO
3
NaNO
3
size 12{"NaNO" rSub { size 8{3} } } {}
|
6.7 |
7.6 |
8.5 |
9.4 |
10.4 |
11.3 |
| KCl |
3.4 |
4.0 |
4.6 |
5.1 |
5.5 |
5.8 |
|
KNO
3
KNO
3
size 12{"KNO" rSub { size 8{3} } } {}
|
1.3 |
3.2 |
5.2 |
7.0 |
9.0 |
11.0 |
- Calculate the initial molarities of
NaNO3NaNO3 size 12{"NaNO" rSub { size 8{3} } } {}, KCl, NaCl,
KNO3KNO3 size 12{"KNO" rSub { size 8{3} } } {}.
- The first batch of crystals is formed at 10°C. Describe these crystals in terms of shape, color, consistency, etc.
- Use the solubility chart to determine the identity of the crystals. Explain your answer.
- What was the temperature of the filtrate when the second batch of crystals formed?
- Describe these crystals in terms of shape, color, consistency, etc.
- Assuming that the first batch of crystals crystallized completely, removing all of it’s ions from solution, what is the identity of the second batch of crystals? (use your solubility chart and explain your answer)
- The reason that we see a third batch of crystals is that we don’t get complete crystallization of the salts in the 1st and 2nd batches of crystals. Assuming there is a small concentration of each ion left in solution after the second batch of crystals is formed, what is the identity of the third batch of crystals according to your solubility chart?
- Describe these crystals in terms of shape, color, consistency, etc.
BONUS QUESTION:
Describe two different tests that we could use to determine the identity of each set of crystals. To receive full credit you must give the name of the test, a description of the theory, and describe the results that you would for your proposed identity of each set of crystals. (5 points max.)
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