Identification of Unknown Inorganic Ions - · PDF file5 Qualitative Analysis ... • Use a flowchart to separate and identify mixtures of common unknown ions. ... drop some salt in

121

CHAPTER

5 Qualitative AnalysisIdentification of Unknown Inorganic Ions

ObjectivesWhen you finish this lab you should be able to:

• Perform“spottest”precipitationreactions.• Writeandbalanceprecipitationreactionequations.• Applysolubilityrulestoidentifyunknownionsusing“spottest”reactions.• Useaflowcharttoseparateandidentifymixturesofcommonunknownions.

BackgroundIn a later experiment you will perform a quantitative analysis to find out how much nickel is in an unknown sample; in this experiment you will be doing qualitative analysis. When performing a qualitative analysis, the scientist does not know which elements, ions, or compounds are present in an unknown sample.

The key to finding which unknown ions are present in a solution will be to make them come out of the solution in the form of precipitates. You will add a few drops of a known reagent to your unknown and then observe the result. This procedure is called a “spot test.” But how do you know which ions are coming out of solution? The answer to that question is by using solubility rules. The solubility rules you will use in this experiment can be summarized as follows:

Solubility Rules

1. All ammonium (NH4) and alkali metal (Group 1A) compounds are soluble.

2. All compounds containing nitrate (NO3–) are soluble.

3. Most hydroxides (OH–) are not soluble. Exceptions are hydroxides of Group 1A and ammonium (see rule #1), and barium hydroxide [Ba(OH)2] which are soluble.

4. Most chlorides (Cl–), bromides (Br–), and iodides (I–) are soluble. Exceptions are chloride, bromide, and iodide compounds containing silver (Ag), mercury (Hg2

2), and lead (Pb2) which are not soluble.

continued

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Experimental Chemistry

5. Most carbonates (CO32–), phosphates (PO4

3–), sulfides (S2–), and chromates (CrO42–)

are not soluble. Exceptions are carbonate, phosphate, and sulfide compounds containing group 1A metals and ammonium, which are soluble.

6. Most sulfates (SO42–) are soluble. Exceptions include the following sulfate com-

pounds, which are not soluble:

barium sulfate (BaSO4) *calcium sulfate (CaSO4)

mercury(II) sulfate (HgSO4) *silver sulfate (Ag2SO4)

lead sulfate (PbSO4)

*These compounds are classified as “slightly soluble,” but for the purposes of this experiment, they can be considered insoluble.

Ions in olutionLet’s discuss for a moment what happens when an ionic compound such as sodium chloride “dissolves” in water. NaCl, or table salt, exists as a solid crystal, but when you drop some salt in water, the crystals seem to disappear. The water tastes salty, so you know the salt is still there, but why can’t we see it? The answer to this question lies in the fact that water, H2O, is a polar molecule. The electrons involved in the O–H bonds in water spend more time around the oxygen than around the two hydrogens. Thus the oxygen side of the water molecule has a slight negative charge while the hydrogen side has a slight positive charge. See Figure 1.

When water comes in contact with many ionic compounds, the water molecules sur-round the individual ions. The electrically charged ions get separated from the solid crystal and become completely surrounded by the polar water molecules. This process causes the ionic crystal to dissociate or “dissolve.” See Figure 2.

O

H H

slight negative charge (�–)

slight positive charge (�+)

Figure 1. The polarity of water.

Some combinations of positive and negative ions have such a strong attraction for each other that water does not break up the ionic crystal lattice. These compounds do not dissolve in water. When combinations of positive and negative ions are created by mixing two solutions, the ions combine to form solid crystalline precipitates which “fall out” of solution. Which compounds break up and dissolve in water, and which do not dissolve in water and form precipitates when solutions are mixed? Check the solubility rules!

Qualitative Analysis: Identification of Unknown Inorganic Ions

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Chapter 5

In this experiment you will test solutions for the presence of five metal ions. The solu-tions were prepared by dissolving nitrate compounds (rule #2) in water. You will use test reagents that were prepared by dissolving other ionic compounds in water. The dissociation reactions in Table 1 describe what happens when these ionic compounds dissolve.

Positive and negative ionsin a solid crystal

Positive ion surrounded by water

Negative ion surrounded by water

+H H

Many water molecules

O

H H

O

H H

O

O

O

H

H

O

H

H

HH

H

H

H H

H

H

O

O

O

©Hayden-McNeil, LLC

Figure 2. An ionic crystal dissolves in water.

Another reagent that you will be using is dimethylglyoxime, DMG. DMG is an organic compound that reacts with nickel ion to form a large, nickel-containing molecule that we will refer to as Ni(DMG)2. Ni(DMG)2 is not soluble in water, and it forms as a bright pink precipitate when DMG is added to a solution that contains nickel. The reaction of DMG with nickel ion will be discussed in the background material for the quantitative analysis experiment (Experiment 9).

Reactions of Ions in olutionIf you were to dissolve two different ionic compounds in water and mix the solutions, the ions would have a chance to react with one another. Let’s use silver nitrate, AgNO3 and potassium chromate, K2CrO4, as an example. According to the solubility rules, both of these compounds dissolve in water. The dissociation reactions to form the two solu-tions are given in Table 1. If the two solutions are mixed, you would initially have two positive ions, K and Ag and two negative ions, NO3

– and CrO42– in the mixture.

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Table 1. issociation reactions for ionic compounds and reagents used in this eperiment.

Metal Ions

Lead (Pb2) Pb(NO3)2 → Pb2 2 NO3–

Iron (Fe3) Fe(NO3)3 → Fe3 3 NO3–

Nickel (Ni2) Ni(NO3)2 → Ni2 2 NO3–

Silver (Ag) AgNO3 → Ag NO3–

Barium (Ba) Ba(NO3)2 → Ba2 2 NO3–

Reagents

Potassium chromate K2CrO4 → 2 K CrO42–

Potassium thiocyanate KSCN → K SCN–

Ammonium hydroxide NH4OH → NH4 OH–

Hydrochloric acid HCl → H Cl–

Sulfuric acid H2SO4 → 2 H SO42–

An important point to remember is when an ionic compound dissociates in water, the hydrated ions (ions surrounded by water molecules) are constantly moving through the solution. Any positive ion in the solution has a chance of reacting with any negative ion.

Let’s say you mixed some potassium chromate solution and some silver chloride solution together, and a brownish-red precipitate formed. What compound would the precipi-tate be? There are four possible ionic compounds that could be made from the four ions in the mixed solution: potassium nitrate; potassium chromate; silver nitrate; silver chromate. Based on the solubility rules, the first three should be soluble in water, so the precipitate must be silver chromate, Ag2CrO4. If you wanted to test for the presence of chromate ion in a solution, adding silver nitrate solution might be a good test.

What is the balanced equation that represents the formation of this precipitate? We know that the reactants are solutions of silver nitrate and potassium chromate, and that one of the products is silver chromate. The other product must be potassium nitrate, which remains in solution. Our first representation of the reaction equation, giving reactants and products with correct chemical formulas, would be:

AgNO3 K2CrO4 → Ag2CrO4 KNO3 (1)

Is this equation correct? The Law of Conservation of Mass states that matter can neither be created nor destroyed. To us that means that there must be the same numbers and types of atoms on the reactants side of a chemical equation as there are on the products side. This is not true for Equation 1, so we need to “balance” the equation.


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In this course we balance the equation by adding coefficients in front of the formulas for products or reactants. We can never balance an equation by adding or changing subscripts, since that would change the chemical identity of a product or reactant. (Look on page 132 for a method of balancing equations describing double replacement reac-tions.) We can balance Equation 1 as follows:

2 AgNO3 K2CrO4 → Ag2CrO4 2 KNO3 (2)

Chemical equations can also indicate the formation of a precipitate. Often subscript letters, (s), (l), (g), and (aq), are used to describe the physical states of products and reactants. A subscript (s) means that the chemical in question is in solid form; (l) stands for liquid, (g) for gas, and (aq) means in aqueous solution or dissolved in water. Let’s apply the physical states to Equation 2.

2 AgNO3 (aq) K2CrO4 (aq) → Ag2CrO4 (s) 2 KNO3 (aq) (3)

The subscript (s) after the formula for silver chromate tells us that the silver chromate comes out of solution as a precipitate.

eparation of Ions in Mixed olutions The reaction we have discussed is an example of a spot test. Spot tests work well when there is only one ionic compound in solution. The problem becomes more complicated when there are a mixture of compounds in solution. When trying to isolate and test for an individual ion in a solution that contains a mixture of ions, it is best to use a qualita-tive analysis flowchart. A flowchart is an outline of a step-by-step procedure which allows you to separate and test for individual cations or anions from a mixed solution. In this experiment we will only be testing for cations (positively charged ions). When the cations have been separated, you can perform spot tests for the presence of specific cations using the solubility rules.

Safety PrecautionsObserve all normal lab safety precautions. Some of the solutions you will be using are toxic and/or corrosive. Clean up spills. If any chemicals or solutions come in contact with your skin, wash the affected areas thoroughly. Wash your hands before you leave the laboratory. Silver solutions can stain skin and clothing and these stains are very difficult to remove. Use designated containers for chemical wastes and for broken or disposable glassware. YOU MUST WEAR EYE PROTECTION AT ALL TIMES. Additional safety precautions are given in the notes for the flowchart procedures. Read them before doing that part of the experiment.

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MaterialsReagents• Solutionsofleadnitrate,iron(ferric)nitrate,nickelnitrate,silvernitrate,andbarium

nitrate• Spottestreagents:solutionsofpotassiumchromate,potassiumthiocyanate,dimeth-

ylglyoxime (DMG), ammonium hydroxide, hydrochloric acid, and sulfuric acid• Flowchartseparationreagents:6Mhydrochloricacid,9Mammoniumhydroxide,

and 4 M nitric acid (HNO3)• Litmuspaper

Euipment• Hotplate,spotplate,andpipettebulb(checkoutfromstockroom)• Centrifuge,Pasteurpipettes,andglasswool(providedinlab)• Testtubes,testtubeholder,beakers,medicinedropper,andstirringrod(fromyour

locker)

Experimental ProceduresI. Performing pot TestsSmall dropper bottles containing metal ion solutions (as nitrates) and spot test reagent solutions will be provided on the lab benches near your workstation.

1. In the wells of your spot plate, test one drop of each metal ion solution with one drop of each spot test reagent. Use the Spot Test Matrix table at the end of this experiment to keep track of the test combinations and to record the results you observe for each test.

2. Check the results of your tests against the solubility rules and write balanced equa-tions for all reactions that produced precipitates. You will need to include these reaction equations in your report.

3. At this point, you are ready to identify a solution that contains only one of the metal ions. One lab partner will select a metal ion solution and put a drop in each of six spot plate wells. The other lab partner will test the drops of this solution with a drop of each spot test reagent and use the observations recorded on your Spot Test Matrix table to identify the metal ion. Then reverse roles and identify another metal ion.

II. eparating Multiple Ion MixturesIt would be nice if analytical chemists only had to test unknowns containing just one compound. Unfortunately, most unknown solutions contain mixtures of many different ions. Simple spot tests on unknown mixtures usually cannot be interpreted unambigu-ously because some spot test reagents can give precipitates with several different ions. Also, the presence of one ion can often interfere with a test for another ion. If your unknown is a mixture, you must first separate the ions before you can successfully use spot tests.


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On page 129 you will find a flowchart outline of procedures to use to separate the metal ions in a mixed unknown solution. The outline is followed by a description of each of the separation steps and some procedural and safety notes that are referenced in the outline and the descriptive text. You should read through all of this material before proceeding as instructed below.

The reagents used in the flowchart separations can be found in the laboratory fume hood. They may be different in concentration from similar solutions in the spot test reagent bottles on your lab bench, so be sure that you are using the right reagents for the flowchart separation procedures. Spot test reagents should only be used when spot tests are called for.

1. Start heating a water bath to almost boiling (see Note C) so it will be ready when you need it.

2. Prepare a sample of approximately 1 mL that contains all five of the metal cations you used in the spot tests. This will take about 5 drops of each metal ion solution.

3. Starting with this as your practice mixture, follow the steps of the flowchart pro-cedures. Refer to detailed description of flowchart separation steps below. Take careful notes on what happens at each step. As you isolate each metal ion, refer to your Spot Test Matrix and the solubility rules to determine an appropriate spot test for the ion you have separated. Perform the spot test(s) to confirm the presence of each ion.

4. Write balanced chemical equations describing the reactions that took place each time something precipitated or dissolved.

III. Analyzing an nknown MixtureWhen you are confident of your ability to separate and identify the five metal ions, ask your lab instructor for an unknown. This solution will contain at least one and not more than four of the five metal ion solutions. Go through the flowchart procedure with the unknown mixture, taking notes on everything you observe. In your lab report you will need to include your observations and provide balanced equations for all of the reactions you observe during the separation steps. Be sure to write down and report your unknown number or letter and identify the metal ions that you believe are present in the unknown. Clean the glassware you have used and wash your hands before you leave the lab.

Detailed Description of Flowchart Separation StepsStep 1In a clean 3″ test tube, add 1–2 drops of 6 M HCl solution to 1 mL of your mixed ion unknown solution. If a precipitate forms, stop. If a precipitate doesn’t form, then add another 2–3 drops. (See Note A.) Stir thoroughly, then cool in an ice bath for 3–4 minutes to allow any solid that forms to precipitate completely out of solution. Centrifuge for one minute or until the liquid above the precipitate (if there is any precipitate) is clear. (See Note B.) Pour the liquid into another clean test tube for step 3.

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Step 2Add 30 drops of deionized water to the solid (if there is any) recovered from step 1. Heat in a water bath for 2–3 minutes, stirring frequently (Note C). If any solid remains, centrifuge it for one minute (Note B), and spot test the liquid.

Step 3Add 9 M NH4OH, a few drops at a time, to the liquid from step 1 (Note A) until the solution is basic. To know when you have added enough NH4OH, test the solution frequently with litmus paper. Litmus is an organic compound that turns blue when basic and red when acidic. Litmus paper is filter paper that has been dipped in a litmus solution. To test with litmus paper, stir the solution with a glass stirring rod and touch the end of the stirring rod to a piece of red litmus paper placed on a clean paper towel. If the wet spot on the litmus paper turns blue, you have added enough NH4OH; if not, continue adding NH4OH and testing. It may take 20 or more drops to make the solu-tion basic. If a precipitate has formed, centrifuge to separate the solid and liquid (Note B). Use the solid (if there is any) in step 4, and the liquid in step 5.

Step 4Dissolve the solid from step 3 (if there is any) in 4 M HNO3. Slowly add the acid, one drop at a time (Note A), until all of the solid dissolves. Perform the appropriate spot test(s).

Step 5To the liquid from step 3 add 8–10 drops of dimethylglyoxime (DMG) solution and stir. If any nickel is present in the solution, it will react with the DMG to form a pink precipitate. If no precipitate forms, spot test the liquid for the only possible remaining metal ion. If there is a pink precipitate (Ni(DMG)2) you will need to separate it from the liquid. First centrifuge for one minute; then filter out any remaining Ni(DMG)2 from the liquid as follows. Transfer supernatant to a clean test tube and centrifuge for one minute. Repeat transfer of supernatant and centrifuge until a clear liquid can be extracted. (Ex-traction of liquid without disruption of the pellet will allow for a clear supernatant.)

Notes on Safety and ProceduresThe following procedural notes contain safety precautions, along with tips on equip-ment and techniques that you will use in this experiment. These notes are referred to in the flowchart and accompanying text.

Note A. Adding Acid or BaseIn some of the steps outlined in the flowchart, you are instructed to add an acid (HCl or HNO3) or base (NH4OH) to a solution or a precipitate. Acid–base reactions often give off heat, so be very careful that your sample doesn’t heat up suddenly and boil out of the test tube. Add the acid or base slowly and be very careful that the open end of the test tube is not pointed at yourself or someone near you.


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Chapter 5

Mixture of Pb2+, Ag+, Fe3+, Ni2+, and Ba2+

PbCl2AgCl

Mixture of Fe3+,Ni2+, and Ba2+

Add HCl (Step 1, Notes A & B)

Add NH4OH (Step 3, Notes A & B)

Add H2O(Step 2, Notes C & B)

liquid

solid

QUALITATIVE ANALYSIS FLOWCHART

AgCl

Pb2+liquid

Spot Test

remaining solid

Ni(DMG)2

Ba2+liquid

Spot Test

solid

Fe(OH)3

Fe3+

Mixture ofNi2+ and Ba2+

liquid Add DMG(Step 5, Note B)

solid

Dissolve withHNO3 (Step 4,Note A)

Spot Test

Note B. CentrifugingAn important aspect of using a centrifuge is to keep it balanced. A good way to do this is to place a test tube containing water at about the same level as your sample test tube directly across from your sample. This will keep the centrifuge from vibrating like a washing machine on spin cycle with an unbalanced load of wet clothes.

Your mixed ion unknown solution will be supplied in a disposable “soft glass” test tube. Do not put a disposable test tube in a centrifuge, since it could shatter and injure some-one with flying glass! When you are finished with your unknown sample, you should either return the disposable test tube to your lab instructor, or empty it into a waste container and dispose of it in the glass waste box. Do not put the empty disposable test tube in your locker where it might be mistaken for a regular “hard glass” test tube.

Do not try to slow or stop the centrifuge rotor with your hands! Wait for the centrifuge to slow down on its own. After your sample has finished spinning, pour off the liquid into a clean test tube. It is important that this be done slowly and smoothly so that the solid at the bottom of the test tube is not disturbed. To do this, touch the outside of

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the lip of the test tube you are pouring from to the inside of the lip of the test tube you are pouring into. Then pour slowly down the side of the clean test tube. This should allow you to transfer the liquid smoothly, without transferring any of the solid. A small amount of liquid will remain with the solid. If you need to use the solid in another step, you can clean it up by adding some deionized water, stirring, centrifuging again, and discarding the rinse water.

Note C. Heating Test TubesThe safest way to heat solutions in test tubes is to set up a hot water bath. Heat a half full beaker (150 or 250 mL) of tap water on a hot plate. When the water is hot enough (as close to boiling as possible), immerse the test tube in the water bath using a test tube holder to support it. This reduces the likelihood that the contents of the test tube might “bump” (boil over), spilling on the lab bench, your notes, your clothing, or you. When heating a test tube do not point the open end toward yourself or another person.

For Your Lab ReportAll scientific reporting must clearly explain what the experimenter was doing in the laboratory, and what he or she observed. The ability to communicate in writing is not an easy skill to learn; it takes much practice. The bottom line for lab reports is that your instructor wants to know a) what you did and b) what you saw. Here are items specific to this experiment that you must include in your report.

1. Turn in your completed Spot Test Observation Matrix.

2. Write balanced equations for all the spot test reactions that produced precipitates. Be sure to include the physical states [(s), or (aq)] for products and reactants. (Hope-fully you did this while you were in lab.)

3. Write a detailed account of what you did and saw when you separated known and unknown ion mixtures. Write balanced equations for all reactions where a precipi-tate formed or was dissolved into solution. Include physical state designations.

4. Also include reaction equations for spot tests you performed after separating ions from mixtures.

5. Report your unknown number or letter and the ions you believe were present.


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Pre-lab QuestionsMake sure you can answer these questions before you come to lab!

1. Write the ions formed from the following compounds when they dissolve in water:

a. AgNO3 →

b. K2CrO4 →

c. NH4OH →

d. KSCN →

2. Predict the results of the reactions between the following pairs of compounds. Balance the reaction equations and include the physical states of products [(l), (g), (s), or (aq)].

a. AgNO3 (aq) HCl (aq) →

b. Pb(NO3)2 (aq) K2CrO4 (aq) → (Hint: Forms a precipitate)

c. Fe(NO3)3 (aq) NH4OH (aq) →

d. Ba(NO3)2 (aq) H2SO4 (aq) →

Ammonium hydroxide, marked on its bottle as NH4OH, is actually a water solution of ammonia, NH3. The ammonia molecule reacts with adjacent water molecules to form the ammonium ion and the hydroxide ion:

NH3 + H2O ➞ NH4+ + OH–

This is a reversible reaction that favors the reverse direction; at equilibrium there is more ammonia and water in the solution than ammonium and hydroxide ions. Ammonia is a “weak base”; it does not do a good job of forming hydroxide ions.

Because the hydroxide ion (OH–) is the reacting ion in most experiments, the reagent’s formula is usually indicated as NH4OH, or ammonium hydroxide, in chemical equations and on reagent bottles in the lab.

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Balancing Equations Describing Double Replacement ReactionsDouble replacement reactions, such as are involved in this qualitative analysis lab, can be easily balanced by a simple four-step process. To illustrate this process, we will develop a balanced equation to describe the reaction between silver nitrate and potas-sium chromate:

Part I. Write formulas with a balanced charge for the reactants.

Step 1: Write the names of the reactants:

Silver nitrate Potassium chromate

Step 2: Write symbols and oxidation numbers for the ions involved:

Ag NO3– K CrO4

2–

Step 3: Using parentheses and subscript numbers, balance the charges involved. Put the charges in boxes below the ions to make the charge stand out.

Ag NO3– K2

CrO42–

1

1 1–

1 1

2 2–

1

Part II. Predict the identity of the reaction products. Positive ions will react only with negative ions. If a reaction takes place, the ion partners will change, hence the term “double replacement.”

Positive Ions

Ag+

K+

NO3–

CrO4–

Negative Ions

The products will be silver chromate and potassium nitrate.

Part III. Using parentheses and subscript numbers, write formulas with balanced charge for the predicted products.

Ag2 CrO4

2– K NO3–

1 2

2– 1

1 1

1– 1

Part IV. Now write and balance the entire equation, using large numbers (called coef-ficients) to balance the molecules. Do not use subscript numbers anymore—you will upset the charge balance.


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Chapter 5

Silver nitrate Potassium chromate Silver chromate Potassium nitrate

2 Ag NO3– K2

CrO42– → Ag2

CrO42– 2 K NO3

–

Now look at the most complex product compound—in this case it is silver chromate. How many silver ions are required to make this molecule? Two. Where do these silver ions come from? The silver nitrate. So let’s use two silver nitrate molecules to get these two silver ions. Put a large two in front of the silver nitrate reactant.

Now—how many chromate ions are required to make this compound? Just one. Where does it come from? The potassium chromate. So we can use the implied 1 in front of this compound.

Now—how many parts are going in that don’t yet have a home? Two nitrate ions, one from each of the silver nitrates, and two potassium ions from the potassium chromate. This is enough to make two potassium nitrates. Put a 2 in front of the potassium nitrate. Do not use subscript numbers—you have already written the simplest charge–balance formula for this compound.

Part V. Check to make sure as many atoms come out as go into the reaction:

Ion In ut

Ag 2 2

NO3– 2 2

K 2 2

CrO4 1 1

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CHAPTER

5

Worksheet

Name I No.

Instructor Course/Section

artner’s Name if applicable ate of Lab Meeting

Qualitative Analysis: Identification of Unknown Inorganic Ions Chapter 5

pot Test bservation Matrix

Reagent

K2CrO4 KSCN NH4OH DMG HCl H2SO4

Pb(NO3)2

Fe(NO3)3

Ni(NO3)2

AgNO3

Ba(NO3)2

Under each reagent indicate what happened (if anything) when mixed with the metal ion shown on the left. Write balanced reactions to represent what happened in each spot test.

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Identification of Unknown Inorganic Ions - · PDF file5 Qualitative Analysis ... • Use a flowchart to separate and identify mixtures of common unknown ions. ... drop some salt in