Determining the Solubility Rules of Ionic Compounds

General Chemistry

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Overview

Source: Laboratory of Dr. Neal Abrams — SUNY College of Environmental Science and Forestry

An ionic compound's solubility can be determined via qualitative analysis. Qualitative analysis is a branch of analytical chemistry that uses chemical properties and reactions to identify the cation or anion present in a chemical compound. While the chemical reactions rely on known solubility rules, those same rules can be determined by identifying the products that form. Qualitative analysis is not typically done in modern industrial chemistry labs, but it can be used easily in the field without the need of sophisticated instrumentation. Qualitative analysis also focuses on understanding ionic and net ionic reactions as well as organizing data into a flow chart to explain observations and make definitive conclusions.

Many cations have similar chemical properties, as do the anion counterparts. Correct identification requires careful separation and analysis to systematically identify the ions present in a solution. It is important to understand acid/base properties, ionic equilibria, redox reactions, and pH properties to identify ions successfully.

While there is a qualitative test for virtually every elemental and polyatomic ion, the identification process typically begins with knowing a "class" of ions being analyzed; cations or anions, elemental or polyatomic, groups or periods, transition or main group. In this experiment, both types of ions, cations and anions, are identified. The cations include polyatomic ions as well.

Cite this Video

JoVE Science Education Database. General Chemistry. Determining the Solubility Rules of Ionic Compounds. JoVE, Cambridge, MA, (2017).

Principles

Identifying cations and anions is based on known chemical reactions between the unknown ion and given reactant. Sometimes, it may be the lack of a reaction that positively identifies the ion as well. All ionic compounds are composed of a cation and an anion, and when a reaction occurs between two different ionic compounds, the cation of one compound is electrostatically attracted to the anion of another, forming a new ionic compound. (NOTE: Some unique ionic compounds have one or more cations or ions. An example would be KNaC4H4O6 or (NH4)2Fe(SO4)2. The overall charge of the ionic compound must still sum to zero.) This type of reaction is known as a metathesis, or double displacement, reaction and is shown below:

wAB(aq) + xCD(aq)yAD(s) + zCB(aq)

molecular reaction

where A and C are cation reactants, B and D are anion reactants, and the compounds are in molar proportions w and x, respectively. The same follows for products AD(s) and CB(aq) with molar ratios of y and z. When a reaction takes place in aqueous solution, the molecular reaction can be written as a combination of free ions and insoluble products known as an ionic reaction:

A+(aq) + B-(aq) + C+(aq) + D-(aq) →AD(s) + B-(aq) + C+(aq)

ionic reaction

An ionic reaction shows both the ions involved in the reaction as well as those that do not participate, known as spectator ions. The formation of the insoluble product AD(s) identifies the reacting ions or could be used to determine a solubility rule for those ions. In all cases, a net ionic reaction underlies all observations, which is a simplified form of the ionic reaction and shows only the ions involved in the reaction.

A+(aq) + D-(aq)→ AD(s)

net ionic reaction

Observing a chemical reaction producing an insoluble product, or precipitate, is a marker for the participants of a net ionic reaction.

Reactions may be unique to a certain cation or anion, or common to all ions within a group or class of reagents. For example, all transition metal ions react with the sulfide ion, S2-, to form insoluble precipitates. Many alkaline earth metals form white precipitates in the presence of carbonate or phosphate ions. More selective identification analyses can be performed with mixed solutions through a combination of solubility rules and chemical reactivity. For example, a solution containing zinc, silver, nickel, and iron could be separated according to the flowchart in Figure 1. Chloride is first added to the solution, precipitating out silver chloride, AgCl. The remaining metals are all precipitated in hydroxide, with excess hydroxide re-dissolving the zinc. The zinc is confirmed in the presence of potassium hexacyanoferrate, forming a green precipitate. The remaining iron and nickel precipitates are collected and excess ammonia is added to dissolve the nickel and the solid iron complex is collected. The iron is re-dissolved in the presence of acid and confirmed with thiocyanate ion. Nickel is positively identified by adding dimethylglyoxime, forming a solid reddish precipitate.

Figure 1
Figure 1. Example flowchart of solution separation.

Procedure

1. General Methods

  1. Preparing for Qualitative Analysis
    1. Reactions are generally done in small test tubes with volume of 5 mL or less.
    2. Solutions need to be fully soluble and should be relatively dilute, typically ~0.1 M.
    3. Reagents should be slowly added drop-wise and observed carefully.
    4. Several common "test solutions" are required to establish solubility rules or identify an unknown ion. These contain ions known to react specifically with certain chemical species (cations or anions).
      1. Common solutions include CaNO3, BaCl2, (NH4)2MoO4, HCl, AgNO3, and NaOH, and other solutions as needed.
  2. Mixing
    1. Mix solutions by tapping or swirling the test tube in a vertical direction. Use a cork or stopper to prevent splashing the solution.
    2. Remove the cork or stopper, then gently heat the solutions with a water bath or cool flame to induce a reaction. Point the test tube away from any individuals in the lab.
  3. Observation and Recovery
    1. Separate the supernatant (non-reacting solution) and precipitate using centrifugation. If more precipitate forms when additional test ion is added, the reaction is incomplete. Continue adding test ion until no more precipitate forms.
    2. Wash the precipitate using centrifugation and pouring or decanting off the supernatant. Add more water and repeat the process for a total of three washings.
    3. Wash large quantities of precipitate by vacuum filtration and recover the dried precipitate from the filter paper.
    4. Note the formation of a precipitate as well as the properties of the precipitate such as color, thickness (gelatinous, cloudy, fine), and crystal formation.
  4. Safety and Waste
    1. Always wear safety eyewear while performing qualitative analysis experiments. Gloves may also be necessary based on the reagents used and products formed.
    2. Proper waste disposal methods must be followed closely. Harmful products can be formed when multiple reactants are combined in one container.

2. Anion Analysis

  1. Identifying phosphate, carbonate, chloride, and sulfide ions; PO43-, CO32-, Cl-, S2-
    1. Phosphate
      1. Add a solution containing phosphate, PO43-, to another solution containing calcium ions, Ca2+. The formation of a white precipitate indicates the formation of calcium phosphate, Ca3(PO4)2.
      2. Since many cations form insoluble products with calcium, a more specific reaction is possible. Add H+ (acid) to Ca3(PO4)2 to dissolve the solid and form HPO42-. Then combine the HPO42- with ammonium molybdate, (NH4)2MoO4. A positive test yields the yellow precipitate ammonium phosphomolybdate, NH4)3PO4(MoO3)12(s). The net ionic reactions are as follows:
        3 Ca2+(aq) + 2 PO43-(aq) Ca3PO4(s)
        Ca3PO4(s) + 2 H+(aq) → 3 Ca2+ + 2 HPO42-(aq)
        HPO42-(aq) + 12 (NH4)2MoO4(aq) + 23 H+(aq)
        (NH4)3PO4 (MoO3)12(s) + 21 NH4+(aq) + 12 H2O(l) 
    2. Carbonate
      1. Carbonate salts are generally insoluble except in the presence of Group 1 and ammonium cations. Add a few drops of calcium chloride, CaCl2, to the carbonate-containing solution. In solutions with high carbonate concentrations, a white precipitate forms and indicates the possible formation of calcium phosphate, CaCO3. The reaction has many interferences, including other anions like phosphate.
        Ca2+(aq) + CO32-(aq)CaCO3(s) 
      2. Add H+ (acid) to a solution containing carbonate, CO32-. The formation of bubbles indicates presence of CO2, signifying CO32- as a reactant. Carbonate ion behaves as a base in the presence of strong acid to form carbon dioxide gas and water.
        CO32-(aq) + H+(aq) → CO2(g) + H2O(l)
    3. Chloride
      1. Add silver nitrate to a chloride-containing solution. The formation of a white precipitate indicates the formation of AgCl(s):
        Ag+(aq) + Cl-(aq) → AgCl(s)
    4. Sulfide
      1. Add a copper chloride solution to a solution containing sulfide. The formation of a black precipitate indicates the formation of copper sulfide, CuS. In general, solutions containing sulfide ions, S2-, react with metal ions to yield an insoluble metal sulfide.
        S2- + Cu2+ → CuS(s). 
        The value of the solubility product, Ksp = 6.3 x 10-36, indicates the high degree of insolubility of the product.

3. Cation Analysis

  1. All alkali metals (group 1) and some alkaline earth metals (group 2) are soluble except under specific conditions.
  2. Nearly all Group 3–13 metals are considered insoluble in the presence of sulfide, carbonate, phosphate, and hydroxide. The color and type of precipitate will vary.
    1. Place a chromium solution in a hydroxide solution. A green precipitate will be observed. The general reaction of a +2 metal with a hydroxide is shown below:
      M2+ + OH- → M(OH)2(s)
    2. It is not possible to differentiate most metal ions based on solubility alone with some notable exceptions:
      1. The addition of silver, Ag+, mercury, Hg22+, or lead, Pb2+ to chloride, bromide, or iodide results in precipitate formation.
      2. The addition of strontium, Sr2+, barium, Ba2+, mercury, Hg22+, or lead, Pb2+ results in a precipitate in the presence of sulfate.
      3. Ba2+ forms a yellow solid in the presence of CrO42-, BaCrO4(s). This is pigment used in oil-based paint commonly known as "barium yellow".
  3. Limited insolubility of metal ions allows for other qualitative tests to positively identify each metal. While some form precipitates, others undergo unique color changes in the presence of chelating ions or molecules. Cation identifications include nickel, iron, aluminum, and zinc; Ni2+, Fe3+, Al3+, Zn2+.
    1. Add nickel (II) in the presence of dimethylglyoxime (H2dmg) to form the rose-red precipitate Ni(H2dmg):
      Ni2+(aq) + 2 H2dmg(aq) → Ni(Hdmg)2(s) + 2 H+(aq) 
    2. Add Iron (III) to thiocyanate ion, SCN- to form the blood-red [FeNCS]2+] complex:
      Fe3+(aq) + SCN-(aq) → [FeNCS]2+(aq)
    3. Aluminum ion
      1. Combine aluminum (III) with pyrocatechol violet in a pH 6 ammonium acetate buffer solution to form a blue solution.
      2. Aluminum (III) is also precipitated in the presence of weak base to form the gelatinous-white Al(OH)3(s) compound. Addition of more base causes the compound to form the clear and colorless [Al(OH)4]-(aq) soluble complex.
    4. Zinc ion
      1. Add zinc (II) to a small amount of base to form a white precipitate. Then add more base to re-dissolve the precipitate and form the soluble [Zn(OH)4]2- complex.
      2. Add zinc (II) to potassium hexacyanoferrate, K4[Fe(CN)6] to form the light green precipitate K2Zn3[Fe(CN)6]2(s):
        3 Zn2+(aq) + 2 K4[Fe(CN)6](aq)K2Zn3[Fe(CN)6]2(s) + 6 K+(aq) 

Trends in the solubility properties of ionic compounds can be used for the qualitative analysis of ionic solutions. When a compound is added to a mixture of ionic solutions, many products can form, each with different solubility properties. If only one product is insoluble, then it alone will leave the solution. By performing sequential reactions, ions in a solution can be systematically identified and isolated.

While a variety of analytical instruments exist for elemental analysis, the techniques are often time-consuming or require transporting samples between laboratories. Qualitative analytical techniques such as examining solubility properties are fast, accessible pre-screening methods for analysis.

This video will introduce the solubility properties of ionic compounds, demonstrate procedures for selectively precipitating ionic compounds, and introduce a few applications of qualitative analysis using solubility trends in industrial settings.

Ionic compounds are composed of a cation and an anion. When a reaction occurs between two different ionic compounds, the cation of one compound is electrostatically attracted to the anion of another, forming a new compound. The ions that do not participate in the reaction are called spectator ions, and are omitted from the net ionic reaction. When an ionic compound dissolves, they reversibly interact with solvent molecules, and the ions dissociate. If the interaction between an ion and the new counter-ion is stronger than between the ion and the solvent molecules, it will be more favorable for the product to be in the solid phase. The formation of solid product from solution is known as precipitation, and the solid is called the precipitate.

Ions can be selectively isolated from solution by inducing reactions with insoluble precipitates. To design these reactions, cations and anions are assigned to broad categories based on solubility trends. Cations are grouped by identifying the anion common to their insoluble reaction products, and anions are likewise grouped by common cations. Solutions of these common ions are used to test for these groups.

When separation is desired for ions belonging to the same group, specialized reagents or concentrated solutions can be used to induce selective reactions once the ions in that group have been isolated. These specialized reagents can also be used to confirm the identity of an isolated ion. Now that you understand the principles behind the qualitative analysis of ions, let's go through a technique for analyzing a solution for phosphate, followed by a procedure for separating a mixture of metal cations.

To analyze a solution for phosphate, first prepare dilute test solutions of aqueous calcium, ammonium orthomolybdate, and concentrated nitric acid. Then, place 5 mL of the unknown solution in a test tube.

Add the calcium solution dropwise to the unknown solution. The formation of a white precipitate could indicate the presence of calcium phosphate, or calcium carbonate. To verify the presence of phosphate, slowly add nitric acid to the test tube. Dissolution of the precipitate indicates that hydrogen phosphate has formed. The lack of gas bubbles indicates that no carbonate is present, as carbonate would have reacted with the acid to form carbon dioxide and water.

Finally, slowly add the ammonium orthomolybdate to the test tube. Ammonium phosphomolybdate forms as a yellow precipitate, confirming the presence of phosphate in the solution.

First, prepare dilute test solutions as listed in the text protocol. Obtain four test tubes and caps suitable for use in a centrifuge. Place a mixture of aqueous zinc, nickel, silver, and iron nitrates in one test tube. To begin separation, first slowly add dilute hydrochloric acid to the mixture, swirling gently. The white precipitate that forms is silver chloride. Continue adding chloride solution until no more precipitate forms.

Separate the liquid, or supernatant, and the solid silver chloride by centrifugation. Decant the supernatant into the second test tube. Wash the silver chloride three times with water and decant each wash into the second test tube. Next, add the sodium hydroxide solution dropwise to the second test tube. Three precipitates will form: white zinc hydroxide, yellow iron hydroxide, and green nickel hydroxide. Continue adding sodium hydroxide until the solid white zinc compound dissolves, forming the soluble zincate ion. Separate the zinc solution and the solid nickel and iron compounds by centrifugation, and then decant the solution into the third test tube. Wash the solids with water three times and decant each into the zinc solution.

Slowly add hydrochloric acid to the zinc solution in the third test tube until zinc hydroxide precipitates and then dissolves.

Then, add potassium hexacyanoferrate dropwise to the zinc solution to form potassium zinc hexacyanoferrate as a white precipitate. Now, to the test tube containing solid nickel hydroxide and iron hydroxide, slowly add ammonia to form the soluble blue nickel hexammine ion. Separate the nickel solution from the solid iron hydroxide by centrifugation and decant the nickel solution into the fourth test tube. Wash the iron hydroxide three times with water and decant the washes into the nickel solution. Then, slowly add dimethylglyoxime to the nickel solution to form nickel dimethylglyoxime as a red precipitate. To the solid iron hydroxide, carefully add concentrated hydrochloric acid to form a solution of ferric chloride. To confirm the presence of iron, add thiocyanate to form the deep red thiocyanatoiron cation.

The simplicity and speed of performing qualitative analysis of ions in solution makes this technique widely used in environmental chemistry and industry.

When water contains a high concentration of metal cations such as calcium or magnesium, it is called hard water. These metal cations can react with anions in the water such as carbonate to form chalky deposits that clog pipes or hot water heaters. Water hardness can be assessed by adding a carbonate solution to a water sample. White precipitate indicates high levels of calcium.

Phosphate is an important nutrient for many forms of life and is therefore used in both industrial and garden fertilizers, but an excess of phosphate can be detrimental, particularly in freshwater environments. Wastewater in residential and commercial areas can be tested for phosphates by adding nitric acid and ammonium orthomolybdate. Yellow precipitate indicates high levels of phosphates.

You've just watched JoVE's introduction to solubility rules for ions. You should now be familiar with the principles of ionic reactions, a few procedures for qualitative analysis of solutions, and some applications of qualitative analysis using solubility.

Thanks for watching!

Applications and Summary

The reactions shown here can be used to identify the presence of a class of cations or anions or be used very specifically for a certain ion. Because two reagents are used in the analyses, either reagent can be typically detected using the other. For example, instead of analyzing for the presence of chloride using silver ion, silver ion can be identified using chloride. A combination of common rules of precipitation followed by specific colorimetric or precipitation tests can be used to positively identify nearly every ion, atomic or polyatomic, available. At the same time, most of those same rules can be established by reacting anions and cations together systematically to generate a set of rules for cation and anion solubility.

Qualitative analysis and rules related to solubility are common experiments in the general chemistry laboratory. This is due, in part, to the ease, speed, and inexpensive nature of the tests. It is for these reasons that qualitative tests are also used in field-based analyses and confirmatory lab tests. For example, a geology firm may wish to know if significant quantities of nickel exist in stream runoff from a mine. A simple test by adding the water to dimethylgloxime is selective for nickel ion. Similarly, water-quality authorities can use barium (or some other group 2 metals) to detect carbonate in water, thus detecting the level of water hardness. Advanced instrumentation is used, however, where quantitative results are required or multiple ions need to be identified at very low levels. This includes various forms of mass spectroscopy as well as ion chromatography and light spectroscopy.

References

  1. Eaton, A. Standard Methods for the Examination of Water & Wastewater. Centennial ed. Washington, DC: American Public Health Association (2005).

1. General Methods

  1. Preparing for Qualitative Analysis
    1. Reactions are generally done in small test tubes with volume of 5 mL or less.
    2. Solutions need to be fully soluble and should be relatively dilute, typically ~0.1 M.
    3. Reagents should be slowly added drop-wise and observed carefully.
    4. Several common "test solutions" are required to establish solubility rules or identify an unknown ion. These contain ions known to react specifically with certain chemical species (cations or anions).
      1. Common solutions include CaNO3, BaCl2, (NH4)2MoO4, HCl, AgNO3, and NaOH, and other solutions as needed.
  2. Mixing
    1. Mix solutions by tapping or swirling the test tube in a vertical direction. Use a cork or stopper to prevent splashing the solution.
    2. Remove the cork or stopper, then gently heat the solutions with a water bath or cool flame to induce a reaction. Point the test tube away from any individuals in the lab.
  3. Observation and Recovery
    1. Separate the supernatant (non-reacting solution) and precipitate using centrifugation. If more precipitate forms when additional test ion is added, the reaction is incomplete. Continue adding test ion until no more precipitate forms.
    2. Wash the precipitate using centrifugation and pouring or decanting off the supernatant. Add more water and repeat the process for a total of three washings.
    3. Wash large quantities of precipitate by vacuum filtration and recover the dried precipitate from the filter paper.
    4. Note the formation of a precipitate as well as the properties of the precipitate such as color, thickness (gelatinous, cloudy, fine), and crystal formation.
  4. Safety and Waste
    1. Always wear safety eyewear while performing qualitative analysis experiments. Gloves may also be necessary based on the reagents used and products formed.
    2. Proper waste disposal methods must be followed closely. Harmful products can be formed when multiple reactants are combined in one container.

2. Anion Analysis

  1. Identifying phosphate, carbonate, chloride, and sulfide ions; PO43-, CO32-, Cl-, S2-
    1. Phosphate
      1. Add a solution containing phosphate, PO43-, to another solution containing calcium ions, Ca2+. The formation of a white precipitate indicates the formation of calcium phosphate, Ca3(PO4)2.
      2. Since many cations form insoluble products with calcium, a more specific reaction is possible. Add H+ (acid) to Ca3(PO4)2 to dissolve the solid and form HPO42-. Then combine the HPO42- with ammonium molybdate, (NH4)2MoO4. A positive test yields the yellow precipitate ammonium phosphomolybdate, NH4)3PO4(MoO3)12(s). The net ionic reactions are as follows:
        3 Ca2+(aq) + 2 PO43-(aq) Ca3PO4(s)
        Ca3PO4(s) + 2 H+(aq) → 3 Ca2+ + 2 HPO42-(aq)
        HPO42-(aq) + 12 (NH4)2MoO4(aq) + 23 H+(aq)
        (NH4)3PO4 (MoO3)12(s) + 21 NH4+(aq) + 12 H2O(l) 
    2. Carbonate
      1. Carbonate salts are generally insoluble except in the presence of Group 1 and ammonium cations. Add a few drops of calcium chloride, CaCl2, to the carbonate-containing solution. In solutions with high carbonate concentrations, a white precipitate forms and indicates the possible formation of calcium phosphate, CaCO3. The reaction has many interferences, including other anions like phosphate.
        Ca2+(aq) + CO32-(aq)CaCO3(s) 
      2. Add H+ (acid) to a solution containing carbonate, CO32-. The formation of bubbles indicates presence of CO2, signifying CO32- as a reactant. Carbonate ion behaves as a base in the presence of strong acid to form carbon dioxide gas and water.
        CO32-(aq) + H+(aq) → CO2(g) + H2O(l)
    3. Chloride
      1. Add silver nitrate to a chloride-containing solution. The formation of a white precipitate indicates the formation of AgCl(s):
        Ag+(aq) + Cl-(aq) → AgCl(s)
    4. Sulfide
      1. Add a copper chloride solution to a solution containing sulfide. The formation of a black precipitate indicates the formation of copper sulfide, CuS. In general, solutions containing sulfide ions, S2-, react with metal ions to yield an insoluble metal sulfide.
        S2- + Cu2+ → CuS(s). 
        The value of the solubility product, Ksp = 6.3 x 10-36, indicates the high degree of insolubility of the product.

3. Cation Analysis

  1. All alkali metals (group 1) and some alkaline earth metals (group 2) are soluble except under specific conditions.
  2. Nearly all Group 3–13 metals are considered insoluble in the presence of sulfide, carbonate, phosphate, and hydroxide. The color and type of precipitate will vary.
    1. Place a chromium solution in a hydroxide solution. A green precipitate will be observed. The general reaction of a +2 metal with a hydroxide is shown below:
      M2+ + OH- → M(OH)2(s)
    2. It is not possible to differentiate most metal ions based on solubility alone with some notable exceptions:
      1. The addition of silver, Ag+, mercury, Hg22+, or lead, Pb2+ to chloride, bromide, or iodide results in precipitate formation.
      2. The addition of strontium, Sr2+, barium, Ba2+, mercury, Hg22+, or lead, Pb2+ results in a precipitate in the presence of sulfate.
      3. Ba2+ forms a yellow solid in the presence of CrO42-, BaCrO4(s). This is pigment used in oil-based paint commonly known as "barium yellow".
  3. Limited insolubility of metal ions allows for other qualitative tests to positively identify each metal. While some form precipitates, others undergo unique color changes in the presence of chelating ions or molecules. Cation identifications include nickel, iron, aluminum, and zinc; Ni2+, Fe3+, Al3+, Zn2+.
    1. Add nickel (II) in the presence of dimethylglyoxime (H2dmg) to form the rose-red precipitate Ni(H2dmg):
      Ni2+(aq) + 2 H2dmg(aq) → Ni(Hdmg)2(s) + 2 H+(aq) 
    2. Add Iron (III) to thiocyanate ion, SCN- to form the blood-red [FeNCS]2+] complex:
      Fe3+(aq) + SCN-(aq) → [FeNCS]2+(aq)
    3. Aluminum ion
      1. Combine aluminum (III) with pyrocatechol violet in a pH 6 ammonium acetate buffer solution to form a blue solution.
      2. Aluminum (III) is also precipitated in the presence of weak base to form the gelatinous-white Al(OH)3(s) compound. Addition of more base causes the compound to form the clear and colorless [Al(OH)4]-(aq) soluble complex.
    4. Zinc ion
      1. Add zinc (II) to a small amount of base to form a white precipitate. Then add more base to re-dissolve the precipitate and form the soluble [Zn(OH)4]2- complex.
      2. Add zinc (II) to potassium hexacyanoferrate, K4[Fe(CN)6] to form the light green precipitate K2Zn3[Fe(CN)6]2(s):
        3 Zn2+(aq) + 2 K4[Fe(CN)6](aq)K2Zn3[Fe(CN)6]2(s) + 6 K+(aq) 

Trends in the solubility properties of ionic compounds can be used for the qualitative analysis of ionic solutions. When a compound is added to a mixture of ionic solutions, many products can form, each with different solubility properties. If only one product is insoluble, then it alone will leave the solution. By performing sequential reactions, ions in a solution can be systematically identified and isolated.

While a variety of analytical instruments exist for elemental analysis, the techniques are often time-consuming or require transporting samples between laboratories. Qualitative analytical techniques such as examining solubility properties are fast, accessible pre-screening methods for analysis.

This video will introduce the solubility properties of ionic compounds, demonstrate procedures for selectively precipitating ionic compounds, and introduce a few applications of qualitative analysis using solubility trends in industrial settings.

Ionic compounds are composed of a cation and an anion. When a reaction occurs between two different ionic compounds, the cation of one compound is electrostatically attracted to the anion of another, forming a new compound. The ions that do not participate in the reaction are called spectator ions, and are omitted from the net ionic reaction. When an ionic compound dissolves, they reversibly interact with solvent molecules, and the ions dissociate. If the interaction between an ion and the new counter-ion is stronger than between the ion and the solvent molecules, it will be more favorable for the product to be in the solid phase. The formation of solid product from solution is known as precipitation, and the solid is called the precipitate.

Ions can be selectively isolated from solution by inducing reactions with insoluble precipitates. To design these reactions, cations and anions are assigned to broad categories based on solubility trends. Cations are grouped by identifying the anion common to their insoluble reaction products, and anions are likewise grouped by common cations. Solutions of these common ions are used to test for these groups.

When separation is desired for ions belonging to the same group, specialized reagents or concentrated solutions can be used to induce selective reactions once the ions in that group have been isolated. These specialized reagents can also be used to confirm the identity of an isolated ion. Now that you understand the principles behind the qualitative analysis of ions, let's go through a technique for analyzing a solution for phosphate, followed by a procedure for separating a mixture of metal cations.

To analyze a solution for phosphate, first prepare dilute test solutions of aqueous calcium, ammonium orthomolybdate, and concentrated nitric acid. Then, place 5 mL of the unknown solution in a test tube.

Add the calcium solution dropwise to the unknown solution. The formation of a white precipitate could indicate the presence of calcium phosphate, or calcium carbonate. To verify the presence of phosphate, slowly add nitric acid to the test tube. Dissolution of the precipitate indicates that hydrogen phosphate has formed. The lack of gas bubbles indicates that no carbonate is present, as carbonate would have reacted with the acid to form carbon dioxide and water.

Finally, slowly add the ammonium orthomolybdate to the test tube. Ammonium phosphomolybdate forms as a yellow precipitate, confirming the presence of phosphate in the solution.

First, prepare dilute test solutions as listed in the text protocol. Obtain four test tubes and caps suitable for use in a centrifuge. Place a mixture of aqueous zinc, nickel, silver, and iron nitrates in one test tube. To begin separation, first slowly add dilute hydrochloric acid to the mixture, swirling gently. The white precipitate that forms is silver chloride. Continue adding chloride solution until no more precipitate forms.

Separate the liquid, or supernatant, and the solid silver chloride by centrifugation. Decant the supernatant into the second test tube. Wash the silver chloride three times with water and decant each wash into the second test tube. Next, add the sodium hydroxide solution dropwise to the second test tube. Three precipitates will form: white zinc hydroxide, yellow iron hydroxide, and green nickel hydroxide. Continue adding sodium hydroxide until the solid white zinc compound dissolves, forming the soluble zincate ion. Separate the zinc solution and the solid nickel and iron compounds by centrifugation, and then decant the solution into the third test tube. Wash the solids with water three times and decant each into the zinc solution.

Slowly add hydrochloric acid to the zinc solution in the third test tube until zinc hydroxide precipitates and then dissolves.

Then, add potassium hexacyanoferrate dropwise to the zinc solution to form potassium zinc hexacyanoferrate as a white precipitate. Now, to the test tube containing solid nickel hydroxide and iron hydroxide, slowly add ammonia to form the soluble blue nickel hexammine ion. Separate the nickel solution from the solid iron hydroxide by centrifugation and decant the nickel solution into the fourth test tube. Wash the iron hydroxide three times with water and decant the washes into the nickel solution. Then, slowly add dimethylglyoxime to the nickel solution to form nickel dimethylglyoxime as a red precipitate. To the solid iron hydroxide, carefully add concentrated hydrochloric acid to form a solution of ferric chloride. To confirm the presence of iron, add thiocyanate to form the deep red thiocyanatoiron cation.

The simplicity and speed of performing qualitative analysis of ions in solution makes this technique widely used in environmental chemistry and industry.

When water contains a high concentration of metal cations such as calcium or magnesium, it is called hard water. These metal cations can react with anions in the water such as carbonate to form chalky deposits that clog pipes or hot water heaters. Water hardness can be assessed by adding a carbonate solution to a water sample. White precipitate indicates high levels of calcium.

Phosphate is an important nutrient for many forms of life and is therefore used in both industrial and garden fertilizers, but an excess of phosphate can be detrimental, particularly in freshwater environments. Wastewater in residential and commercial areas can be tested for phosphates by adding nitric acid and ammonium orthomolybdate. Yellow precipitate indicates high levels of phosphates.

You've just watched JoVE's introduction to solubility rules for ions. You should now be familiar with the principles of ionic reactions, a few procedures for qualitative analysis of solutions, and some applications of qualitative analysis using solubility.

Thanks for watching!

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