You will be determined according to the following:

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 \( l \) } {}

*Slightly soluble.

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.

CAUTION WEAR EYE PROTECTION

1. Copper (II) sulfate + sodium carbonate

2. Copper(II) sulfate + barium chloride

3. Copper(II) sulfate + sodium phosphate

4. Sodium carbonate + sulfuric acid

5. Sodium carbonate + hydrochloric acid

6. Cadmium chloride + sodium sulfide

7. Cadmium chloride + sodium hydroxide

8. Nickel chloride + silver nitrate

9. Nickel chloride + sodium carbonate

10. Hydrochloric acid + sodium hydroxide

11. Ammonium chloride + sodium hydroxide

12. Sodium acetate + hydrochloric acid

13. Sodium sulfide + hydrochloric acid

14. Lead nitrate + sodium sulfide

15. Lead nitrate + sulfuric acid

16. Potassium chloride + sodium nitrate

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)