Chem 1B Saddleback College Dr. White 1 Introduction to Qualitative Analysis Chemical analysis can be either quantitative or qualitative in nature. Quantitative analysis seeks to answer questions like "What is the concentration of my acid solution?" or "What is the percentage of calcium in my marble sample?". Experiments using the techniques of quantitative analysis must be performed carefully, with the goal of determining the result as accurately as possible. On the other hand, qualitative analysis seeks to answers questions like "Does chalk contain carbonate ion or sulfate ion?" or "What cations are present in my unknown sample?". Although your goal is identification, rather than quantitation, experiments still must be performed with meticulous care. You must make careful observations, keep a careful record in your notebook, and perform all laboratory operations correctly. If your technique is sloppy, your cation identification is likely to be ambiguous or in error. The techniques used in qualitative analysis depend upon the material being analyzed. For inorganic compounds, qualitative analysis often involves the identification of ions present in a sample. This is the type of analysis that will be done during the qualitative analysis labs in Chemistry 1B; you will be given an aqueous solution containing a mixture of several metal cations that you must identify. The techniques you will learn can be used to identify ions occurring in other types of samples such as minerals, ground water and industrial waste streams. The procedures used will provide you with an opportunity to apply many general chemistry principles including those involved in solubility, acid- base chemistry, oxidation- reduction reactions, ionic equilibria, precipitation reactions and complex ion formation. Focusing on the goal to correctly identify the metal cations present in an unknown, the simplest scheme possible would involve one that has a specific reagent to test for each different cation. In such a scheme, each reagent would be required to give an easily recognized confirmation test, such as color change or precipitate formation, for only one of the cations in the mixture, regardless of the other cations present. However, different metal cations can sometimes exhibit similar behavior and a specific reagent for each separate cation is not possible. In other words, individual components in our unknowns would most likely interfere with one another. Therefore, in the scheme that we will employ, reagents will be used to separate the ions in our samples into groups. Each group will then be analyzed for the presence or absence of individual metal cations. The most common way to subdivide into smaller groups is by selective precipitation, in which a small group of cations is chemically precipitated. The ions in the precipitate can then be physically separated from those remaining in solution by centrifuging. The precipitate (solid) settles out and the solution (supernatant) is transferred into another container. In this way, the initial large group is separated into smaller and smaller groups until definitive tests can be run to confirm the presence or absence of each specific cation. The cations in our qualitative analysis scheme will be organized into groups called I- IV. The groups are summarized in Table 1 below. Table 1: The cation groups in the quantitative analysis scheme. Group Cations In Group I Ag +, Pb 2+, Hg 2+ 2 II Cu 2+, Bi 3+, Sn 4+, Sb 2+ III Ni 2+, Co 2+, Fe 3+, Al 3+, Zn 2+ IV & V Ba 2+, Ca 2+, Mg 2+, Na +, K +, NH + 4 The groups in qualitative analysis are based solely upon the solubility behavior of the cations under specific conditions. For example, Group I consists of cations that form insoluble chlorides in acidic solution. Within each of these groups, the analysis may require that there be further separations into subgroups. A specific or confirmatory test will be carried out for each ion when separations have ensured that interfering ions have been removed. Sometimes this will mean isolation of a given ion from all other cations. In other cases, it will be possible to carry out confirmatory tests in the presence of one or more other cations of the same group. To be successful, care must be taken to follow the procedures carefully; components that are not separated correctly may interfere with later tests. As you work to identify the cations present in your unknown, you should follow directions carefully so that your analysis will be more likely to produce clear results. However, keep in mind that many different variables are at play in determining how a given sample will behave in the presence of each reagent used. You must be prepared for unexpected results, be observant and keep careful, accurate notes of your observations in your notebook. When confronted with ambiguity or uncertainly, do not despair! Rather, try to apply knowledge already gained about your sample to limit the possibilities. You may need to repeat experiments, possibly making modifications to the procedure or trying
Chem 1B Saddleback College Dr. White 2 alternative tests - WITH YOUR INSTRUCTOR S APPROVAL. View this analysis as a type of puzzle that you must solve. You will be successful in solving the puzzle if you understand the background chemistry with which you are working, employ good techniques, carefully follow procedures, and carefully observe and keep thorough notes. The following sections discuss laboratory techniques, flow charts to summarize results, chemical concepts, and confirmation tests. You should study these sections carefully in order to be successful in qualitative analysis. LABORATORY TECHNIQUES USED IN INORGANIC QUALITATIVE ANALYSIS Cleanliness Make sure that all test tubes and stirring rods are clean. Rinse the test tubes with deionized water and shake out as much of the liquid as possible before use. Rinse stirring rods before using them. Rinse droppers before reusing them for a different solution. You will use the qualitative analysis reagents from the back of the room. Make sure they don t get contaminated. Don t set caps down or mix them up. Take any empty, old, cap- less, or suspicious looking bottles to the stockroom. Return the bottles to the proper place at the end of the period. For every bottle I find on your lab bench, I will take off a point from your lab report. Adding Reagents Don t use a graduated cylinder to measure volumes estimate volumes by counting drops. (1 ml is about 18 drops). Or, make marks on your test tubes for 1 ml, 2 ml, etc. Never place the tip of a dispensing dropper into your test solution in the test tube. Insert the tip about 0.5 cm below the top of the test tube, and release the indicated number of drops. Mixing If a small amount of liquid is present in a test tube, it may be mixed by flicking the base of the test tube with a finger while holding the test tube lightly by the top. Never shake a test tube that is capped with a finger or cork. Getting chemicals on fingers is an excellent means of introducing them into your body. Even if gloves are used, using a finger to cap a test tube is an easy means of contaminating other solutions. Also, if a test tube is capped with a finger or a cork, pressure may build up due to the evolution of heat or a gas in the test tube. Pressure build up can cause chemicals to spray out of the test tube. If the flicking technique is unsuccessful, or if the test tube is more than one- third full, a clean glass stirring rod should be used to mix the contents. Unless otherwise directed, always mix thoroughly after adding each reagent before making observations, checking ph or proceeding to the next step. Observations Always describe the color and clarity of mixtures and reagents before mixing and what the mixture looks like after mixing, heating, centrifuging, etc. Give precipitates a chance to form add the reagent, mix, let settle and then add one more drop to make sure the precipitation is complete before recording observations or moving on. The following might be recorded in your notebook for step 1 of the Group I qualitative analysis procedure: Starting unknown solution - clear and * 6 M HCl - clear and Added 10 drops of the HCl to the unknown solution and stirred mixture. A cloudy white, ppt** formed. After centrifuging, white precipitate settled to bottom, clear and snt***. *A note on terminology: clear means not cloudy, means uncolored (like water). The two terms are not synonymous. **ppt is shorthand for precipitate ***snt is shorthand for supernatant Centrifuging Separate the precipitate and supernatant cleanly using the centrifuge. Make sure to pour off all the supernatant, wash the precipitate with DI water, mix, let settle or centrifuge, and pour off the wash liquid. Repeat if necessary. Be sure that the test tubes in use are the appropriate size for the centrifuge. A tube of approximately the same mass in the opposite slot of the centrifuge must be used to balance the centrifuge. This is easily accomplished by using a test tube of the same size, which is filled to approximately the same height with water. If you are simultaneously testing a known and an unknown, they can usually be used to balance each other. Other test tubes in the centrifuge may be of different masses, but each opposite pair should be matched. If the centrifuge is too unbalanced, it may "walk" around the countertop while it is spinning. Be sure that test tubes being centrifuged are neither cracked nor chipped. The stress applied by the centrifuge can cause damaged test tubes to shatter, resulting in chemicals and pieces of glass being scattered inside the centrifuge. A centrifuge without a top is dangerous; always close the top. Do not slow centrifuges down with your hands. They are spinning at a high rate of speed, and if there is any imperfection on the spinning surface, it
Chem 1B Saddleback College Dr. White 3 can catch the flesh and do a great deal of damage in only an instant. Red and Blue Litmus Paper Litmus paper is used to identify acidic and basic solutions because it changes color when exposed to an acid or base. Blue litmus paper turns red under acidic conditions and red litmus paper turns blue under basic conditions. To test a solution with litmus paper, place a clean glass stirring rod into the test solution to be tested and touch it to the litmus paper. FLOW CHARTS It is possible to summarize your results for your qualitative analysis experiments in a flow chart. In the diagram, successive steps in the procedure are linked with lines. Throughout the flow charts, reagent additions (along with the concentration) and other procedures (such as heating or boiling) are indicated along the connecting lines. The formula for each species, along with any identifying physical characteristics (such as color), is given in the box. Each branch indicates the separation of precipitate from the solution. Flow charts showing the behavior of your unknown should be included in your notebook and all your qualitative analysis lab reports. The symbols and formalism used in the flow charts are given below. polyatomic ions (e.g. SO 4 2 ). Some may also involve the reaction of neutral molecules in solution. For writing NET IONIC chemical equations for reactions, the general guidelines are SUMMARIZED as follows: 1. Dissolved ionic compounds, the strong acids (HCI, HBr, HI, H 2SO 4, HNO 3 and HClO 4) and the strong soluble bases (Group 1A metal hydroxides) are completely dissociated/ionized in solution into the corresponding ions. 2. Water, gases, insoluble compounds, weak acids such as acetic acid (HC 2H 3O 2) and weak bases such as ammonia (NH 3) will be present as neutral molecules. The extent of chemical reactions and chemical equilibrium concepts It is important to remember that many chemical reactions do not go to completion. The extent to which a reaction occurs depends on the magnitude of the equilibrium constant (K) for the reaction, and the relative amounts of reagents present. Equilibria can be shifted by adding or removing reactants or products, by adding other reagents and/or by changing the temperature, in accordance with Le Chatelier's principle. By applying this principle, we can force precipitation to occur, cause some sparingly soluble compounds to dissolve, or complex particular ions so that they will not interfere with tests for other ions of interest. An example from the Group II cation group involves the following equilibrium reaction: SnS 2 (s) + 6OH - Sn(OH) 6 2- (aq) + 2S 2- (aq) The reaction can be shifted to the right by adding excess NaOH. This causes the SnS 2 precipitate to dissolve and allows the Sn 4+ ion to be separated from other cations in solution. REACTIONS INVOLVED IN INORGANIC QUALITATIVE ANALYSIS Several types of reactions encountered in qualitative analysis are discussed below along with examples of each. IMPORTANT CHEMICAL CONCEPTS INVOLVED IN QUALITATIVE ANALYSIS Net-ionic equations The reactions we will encounter take place in aqueous solution. Many of the reactions will occur between ions, including monatomic (e.g., Ag +, Cl ) and Acid Base Reactions Acid- base reactions can be classified as Arrehinius, Bronsted- Lowry or Lewis. The following reactions are other examples of this diverse class. 2H 3O + (aq) + CO 3 2 (aq) H 2O (l) + CO 2 (g) Al(OH) 3 (s) + OH - (aq) Al(OH) 4 (aq) Al(OH) 3 (s) + 3H 3O + (aq) Al 3+ (aq) + 6H 2O (l) Compounds such as Al(OH) 3, that can react with either acid or base, are said to be amphoteric. Buffers are a special case of acid- base reactions. A
Chem 1B Saddleback College Dr. White 4 solution is said to be a buffer when it has approximately equal amounts of a weak acid or base and its conjugate salt. Common examples are HC 2H 3O 2/NaC 2H 3O 2 and NH 4Cl/NH 3. A buffered solution resists changes in ph when small amounts of a strong acid or base are added. The weak acid (or base) or the conjugate salt will react to absorb the added OH or H 3O +. Buffers are used as a means of controlling the ph, sometimes in order to regulate the nature of the species in solution. Consider the equilibrium between chromate (CrO 4 ) and dichromate (Cr 2O 7 2 ) 2H 3O + (aq) + 2CrO 4 2 (aq) Cr 2O 7 2 (aq) + 3H 2O(l) Dichromate salts tend to be soluble, whereas some chromate salts are not. Careful ph control using a buffer allows one to control the concentration of chromate ion in solution. Le Chatelier s Principle tells us that adding acid (H 3O + ) shifts the reaction to the right. This increases the amount of dichromate present and decreases the amount of chromate present. We can therefore selectively precipitate specific metal cations as their insoluble chromates, leaving more soluble cations in solution. Groups II and III will employ this concept for the sulfide precipitates. Hydrolysis Reactions Many ions or molecules react with water, resulting in the formation of either H 3O + or OH ions, thus affecting the ph of the solution. This type of reaction, known as hydrolysis, is classified as acid- base. A common and important example is the hydrolysis of ammonia to form ammonium and hydroxide ions: NH 3 (aq) + H 2O (l) NH 4 +(aq) + OH - (aq) When acid is added to an ammonia solution, it reacts with ammonia, forming ammonium ion, shifting the equilibrium to the right, thus increasing the amount of NH 4 + in solution. The equilibrium can be driven to the left by heating to remove NH 3 as gas. Adding a strong base will also shift the equilibrium to the left. The concentration of the species undergoing hydrolysis has an effect on the ph. A solution of 0.1 M NH 3 in water has a ph of 8.9, whereas a concentrated ammonia solution (14.5 M) has a ph of 10.0; an increase of hydroxide ion concentration by a factor of approximately 10. The concentration of ammonia can determine whether an ion precipitates as a hydroxide salt, forms a soluble ammonia complex, or does not react at all. Precipitation Reactions The solubility product constant is related directly to the solubility of a sparingly soluble compound. Tables of K sp values are available and are valuable resources for qualitative analysis. Familiarity with the general solubility rules is also helpful. A table of K sp values (Table 4) and one of general solubility rules (Table 3) is provided. Study these and become familiar with them. When the ion product (Q sp) of ions in solution exceeds the solubility product constant (K sp) of a particular salt, precipitation will occur. Determining the presence of a precipitate is not always trivial. Some solutions become cloudy, which indicates a precipitate has formed. In some cases, you may be able to see grains of solid failing from solution. In other cases, the solution may become milky in appearance. The solution and the precipitate may or may not change color. The presence of a precipitate may be difficult to detect when a solution is dark. In such cases, centrifuging the solution to determine if a precipitate is present may help. The solubility of a particular species may be affected by adding reagents that cause a competing reaction. For instance, silver chloride, AgCl, is insoluble in deionized water, but can be made to dissolve by adding aqueous ammonia. This is due to the extremely favorable formation of the soluble silver ammonia complex, Ag(NH 3) 2 +: AgCl(s) Ag + (aq) + Cl (aq) Ag + (aq) + 2NH 3(aq) Ag(NH 3) 2 +(aq) Formation of the complex ion lowers the concentration of silver ion in solution, causing the first reaction to proceed further to the right, thus allowing more AgCl solid to dissolve. Decomposition Reactions: A decomposition reaction occurs when one chemical species decomposes into one or more different products. An example is the fizzing that results when acid is added to a carbonate: 2H 3O + (aq) + CO 3 2 (aq) H 2CO 3(aq) + 2H 2O(l) H 2CO 3(aq) CO 2(g) + H 2O(l) The carbonate ion undergoes an acid- base reaction, producing carbonic acid. Carbonic acid is very unstable, and spontaneously decomposes into water and carbon dioxide gas. The evolution of carbon dioxide from solution causes the fizzing, and shifts both equilibria further to the right. Oxidation-Reduction Reactions Oxidation- reduction, or redox, reactions are often accompanied by color changes and are used
Chem 1B Saddleback College Dr. White 5 frequently in qualitative analysis to confirm the presence or absence of an ion. Redox reactions are also used as a means of dissolving very insoluble compounds and for converting an ion to a different oxidation state, in which case it may be more easily separated or identified. In the following example, pale green chromium (III) hydroxide, Cr(OH) 3, can be separated from a mixture of hydroxide solids, where its pale color is often masked, by oxidizing it using hydrogen peroxide, H 2O 2, in basic solution to form the chromate ion, CrO 4 2. In solution, the chromate ion is bright yellow. Cr(OH) 3(s) + 4OH (aq) + 3H 2O 2(aq) 2CrO 4 2- (aq) + 8H 2O(l) The most common oxidizing agents are nitric acid, HNO 3, and basic hydrogen peroxide solution. Reducing agents include the iron (II), tin (II), thiosulfate (S 2O 3 2- ), oxalate, and iodide ions. You must be able to balance complicated redox reactions. See your textbook for a review of this concept. Disproportionation These reactions are a special case of redox reactions wherein part of the reagent is oxidized, and an equivalent part, according to the stoichiometry, is reduced. An example is the spontaneous decomposition of hydrogen peroxide where oxygen, in the 1 oxidation state in peroxide, disproportionates into oxygen gas and water: 2H 2O 2(aq) O 2(g) + 2H 2O(l) Hydrogen peroxide solutions will degrade over time due to this reaction. To slow the process and increase the shelf life of the solution, hydrogen peroxide solutions may be stored in a refrigerator. number of water molecules acting as ligands, although the water molecules are often not included when writing reactions. Square brackets, [ ], are often used to indicate the complex ion. For example, the following: CuSO 4(s) Cu 2+ (aq) + SO 4 2 (aq) should more correctly be written as 4H 2O + CuSO 4(s) [Cu(H 2O) 4] 2+ (aq) + SO 4 2 (aq) The presence of water, or other ligands, causes many of the transition metal ions to be colored in solution. When a new ligand that has a higher affinity for the metal ion is added to a solution containing a complex ion, ligand replacement often adding a particular ligand is often a diagnostic test for the presence of a given metal cation. This is a common way of confirming the presence or absence certain cations in solution. Formation of a complex ions can also be used to dissolve a precipitate or to prevent a reaction of a certain cation with a reagent being used to test for or separate a different cation. For example, in the qualitative analysis scheme for the Group I cations, AgCl can be dissolved by adding 6 M NH 3: AgCl (s) + 2NH 3 (aq) Ag(NH 3) 2 + (aq) + Cl - (aq). A table of various complex ions (Table 4) is provided for your reference. It is valuable to become familiar with the complex ions that are listed in the table. COMMON REAGENTS Some common reagents and their uses in qualitative analysis are listed Table 2. You should become familiar with these. Formation of Complex Ions Many common anions and neutral molecules can donate one or more lone pairs of electrons, thus acting as Lewis bases to a Lewis acid. A coordinate covalent bond (one where both electrons are provided by one atom) is formed between the Lewis acid and the Lewis base. Certain metal cations act as Lewis acids. This type of reaction can result in the formation of a complex ion, an ion formed by a central metal cation bonded to from two up to six lone pairs of electrons on surrounding Lewis base species. In a complex ion, the Lewis bases are known as ligands. Water and ammonia are examples of neutral ligands. Anions that readily act as ligands include F, Cl, Br, I, SCN, C 2O 4 2, OH, and CN. Metal cations in aqueous solution tend to have a fixed
Chem 1B Saddleback College Dr. White 6 Table 2: Common Reagents in Qualitative Analysis Reagent Effect and Use 6 M HCl Increases [H + ]; Increases [Cl - ]; Decreases [OH - ];, Dissolves insoluble carbonates, chromates, hydroxides, some sulfates; Destroys hydroxo and ammonia complexes; Precipitates insoluble chlorides 6M HNO 3 Increases [H + ]; Decreases [OH - ]; Dissolves insoluble carbonates, chromates, and hydroxides; Dissolves insoluble sulfides by oxidizing sulfide ion; Destroys hydroxo and ammonia complexes; Good oxidizing agent when hot 6M NaOH Increases [OH - ]; Decreases [H + ]; Forms hydroxo complexes; Precipitates insoluble hydroxides 6M NH 3 Increases [NH 3]; Increases [OH - ]; Decreases [H + ]; Precipitates insoluble hydroxides; Forms NH 3 complexes; Forms a basic buffer with NH + 4 Table 3: General Aqueous Solubility Rules for Selected Cations NH 4 +, Na +, K +, Li +, All common salts of ammonium, sodium, potassium, potassium, nitrate and acetate ions are NO 3 -, C 2H 3O - 2 soluble Cl -, Br -, and I - Most halides are soluble except those of Ag +, Pb 2+, Cu +, and Hg 2 2+. PbCl 2 and PbBr 2 are slightly soluble in hot water. SO 2-4 All sulfates are soluble except those with Pb 2+, Ca 2+, Sr 2+, Ba 2+, Hg 2 2+, and Ag + OH - All common hydroxides are insoluble except those of Group 1A and the larger members of Group IIA (beginning with Ca 2+ ). Some metals are soluble in excess hydroxide due to complex ion formation. CO 3 2-, PO 4 3-, C 2O 4 2-, All carbonates, phosphates, oxalates, and chromates are insoluble except those of Group 1A and CrO 2-4 and NH + 4 S 2- All sulfides are insoluble except those of Group 1A, Group 2A and NH 4 +. Cations Group I Ag + Pb 2+ Hg 2 2+ Table 4: Selected K sp Values at 25 C Anions Cl - S 2- OH - CO 2-3 SO 2-4 C 2O 2-4 CrO 2-4 PO 3-4 1.6 x 10-10 1.6 x 10-5 1.1 x 10-18 8 x 10-51 1.5 x 10-8 8.1 x 10-12 1.5 x 10-5 1.3 x 10-11 1.12x10-12 8.9 x 10-17 3 x 10-28 1.2 x 10-5 7.4 x 10-14 6.3 x 10-7 2.7 x 10-11 3 x 10-13 2.53 x 10-8 3.6 10-17 6.5 10-7 1.75 10-13 2.0x10-9 Group II Cu 2+ + (blue) S 8.5 x 10-45 4.8 x 10-20 2.3 x 10-10 S 4.43 x 10-10 1.4 x 10-37 Bi 3+ S 1.6 x 10-72 S Sn 2+ S 1 x 10-45 5.45 x 10-27 S Sb 2+ S 1.6 x 10-93 S Group III Ni 2+ + (green) S 3 x 10-21 5.48 x 10-16 1.3x10 7 S Co 2+ + (pink) S 1 x 10-21 2.0x10-16 1.0 x 10-10 S Fe 3+ (yellow) S 3.7 x 10-19 4x10 38 4.0x10-10 S Al 3+ S 4.0 x 10-14 1.3x10 33 S 9.8 x 10-21 Zn 2+ S 2.5 x 10-22 3.0x10 16 1.0x10-10 S Group IV/V Ba 2+ S S S 5.0x10 9 1.1x10 10 1.2 x 10 - Ca 2+ S S S 4.5x10 9 2.4x10 4 7.1 x 10-4 2.1 x 10-33 Mg 2+ S S 1.6x10 12 4.0x10 5 S 1.0 x 10-24 Na + S S S S S S S S K + S S S S S S S S NH + 4 S S S S S S S S 10
Chem 1B Saddleback College Dr. White 7 Ion to be confirmed Table 5: Soluble Complexes of Metal Cations (colors observed in the absence of interfering ions) NH 3 OH - Other Ag + Ag(NH 3) + 2 AgCl - 2 Pb 2+ Pb(OH) 2-4 PbCl 2-4 Cu 2+ Cu(NH 3) 2+ 4 blue Cu(OH) 2-4 blue CuCl 2-4 blue Sb 3+ Sb(OH) 3-6 SbS 3 3-, Sb(Cl) 3-6 yellow Sn 4+ Al 3+ Ni 2+ Ni(NH 3) 2+ 6 blue Zn 2+ Zn(NH 3) 2+ 4 Co 2+ Co(NH 3) 2+ 6 pink Al 3+ Reagent to add Group I Cations (Insoluble Chlorides) Lead (II): Pb 2+ Sn(OH) 6 2- Al(OH) 4 - Zn(OH) 4 2- Al(OH) 4 - SnS 3 2-, Sb(Cl) 3-6 yellow NiCl 6 4- yellow Co(SCN) 4 2- blue Fe 3+ FeCl 4 -, FeSCN 2+ yellow red Table 6: Confirmation Tests for the Cations Result Net Ionic Reaction & Reaction Type K 2CrO 4 Yellow ppt Pb 2+ (aq) + CrO 2-4 (aq) PbCrO 4 (s) Silver: Ag + HCl White ppt Ag + (aq) + Cl - (aq) AgCl (s) Mercury (I): Hg 2 2+ HCl then NH 3 White ppt forms after the addition of HCl, then it changes to grey/black after the addition of NH 3 Hg 2Cl 2 (s) + 2 NH 3(aq) HgClNH 2(s) + Hg(s) + NH 4 + (aq) + Cl - (aq) REDOX RXN
Chem 1B Saddleback College Dr. White 8 Ion to be confirmed Reagent to add Result Group II Cations (Insoluble sulfides in an acidic ph) Copper (II): Cu 2+ Na 2S 2O 4 (s), heat Blue solution turns and a reddish- brown or black ppt forms Net Ionic Reaction & Reaction Type Cu(NH 3) 4 2+(aq) + S 2O 4 2- (aq) + 2H 2O (l) Cu(s) + 2SO 3 2- (aq) + 4H + (aq) REDOX RXN Bismuth: Bi 3+ a. Sn 2+ in a basic solution a. Black ppt a. 2 Bi(OH) 3 (s) + 3Sn(OH) 4 2- (aq) 2Bi(s) +3Sn(OH) 4 2- (aq) REDOX RXN b. White ppt b. HCl (usually done in a beaker) b. Bi 3+ (aq) + Cl - (aq) + H 2O (l) BiOCl(s) + 2H + (aq) REDOX RXN Tin (II): Sn 2+ HgCl 2 Gray ppt 2Hg 2+ (aq) + Sn 2+ (aq) + 2Cl - (aq) Hg 2Cl 2 (s) + Sn 4+ (aq) REDOX RXN Antimony (III): Sb 3+ Na 2S 2O 3 (s) Peach (Red- Orange) ppt (note that this is the same confirmation rxn for mercury) S 2O 2-3 (aq) + H 2O (l) SO 2-4 (aq) + H 2S (aq) 2SbOOH (s) + 2H 2S (aq) Sb 2OS 2 (s) + 3H 2O
Chem 1B Saddleback College Dr. White 9 Ion to be confirmed Reagent to add Result Group III Cations: Insoluble in a Basic ph Nickel (II) Ni 2+ Dimethyl- glyoxime reagent (DMG) Rose- Red ppt Net Ionic Reaction & Reaction Type Ni 2+ (aq) + 2DMG (aq) Ni(DMG) 2 (s) Cobalt (II) KNO 2 Yellow ppt Co 2+ (aq) + 7NO - 2 (aq) + 3K + (aq) + 2H + Co 2+ K 3Co(NO 2) 6 (s) + NO (g) + H 2O REDOX RXN Iron (II) and Iron (III): Fe 2+ and Fe 3+ a. KSCN a. Deep Red Solution a. Fe 3+ (aq) + SCN - (aq) FeSCN 2+ (aq) COMPLEX ION FORMATION RXN b. NH 3 b. Brown ppt b. Fe 3+ (aq) + 3OH - (aq) Fe(OH) 3 (s) Aluminum: Al 3+ a. NH 3 a. White Gelatinous ppt a. Al 3+ (aq) + 3OH - (aq) Al(OH) 3 (s) b. catechol violet b. blue solution b. COMPLEX ION RXN Zinc: Zn 2+ K 4Fe(CN) 6 Grey- white ppt (may look green) 3Zn 2+ (aq) + 2K + (aq) + 2Fe(CN) 6 4- (aq) K 2Zn 3[Fe(CN) 6] 2 (s)
Chem 1B Saddleback College Dr. White 10 Ion to be confirmed Reagent to add Result Net Ionic Reaction & Reaction Type Group IV Cations: Barium: Ba 2+ a. K 2CrO 4 a. Yellow ppt a. Ba 2+ (aq) + CrO 4 2- (aq) BaCrO 4 (s) b. H 2SO 4 b. White ppt c. Green Flame b. Ba 2+ (aq) + SO 4 2- (aq) BaSO 4 (s) c. Flame test Calcium: Ca 2+ a. K 2C 2O 4 in a basic solution (NH 3) a. White ppt a. Ca 2+ (aq) + C 2O 4 2- (aq) CaC 2O 4 (s) b. red- orange flame b. flame test Magnesium: Mg 2+ magnesium reagent, 4- (p- nitrophenylazo )resorcinol, and NaOH Blue Lake (a precipitate of Mg(OH) 2 with adsorbed magnesium reagent) Mg 2+ (aq) + 2OH - (aq) Mg(OH) 2 (s) Sodium: Na + Flame test Yellow Flame N/A Potassium: K + Flame Test Violet Flame N/A Ammonium: NH 4 + NaOH, heat Litmus turns from red to blue from NH 3 gas NH 4 + (aq) + OH - (aq) H 2O (l) + NH 3 (g)
Chem 1B Saddleback College Dr. White 11 Name: Lab Day/Time: Qualitative Analysis Introduction Pre-Lab Questions Due at the beginning of Lab for Experiment 5: Qualitative Analysis of Group I Cations 1. What are the groups in qualitative analysis based on? 2. Instead of measuring out volulmes in a graduated cylinder, how should you measure out volumes? In this method what constitutes 1 ml? 3. How do you successfully separate a precipitate and a supernatant? 4. How do you test if a solution is basic? How do you test if a solution is acidic? 5. Flow charts are used to summarize the procedure and your results in Qualitative Analysis labs. How do you show that a mixture was centrifuged and separated into separate test tubes? 6. A common reagent in qualitative analysis is 6 M NH 3. What are the uses of 6 M NH 3? 7. You have a solution that contains either Sn 2+ or Ca 2+. List a way that you could determine which of the cations is present (Hint: use the solubility rules).
Chem 1B Saddleback College Dr. White 12 8. You have a solution that contains only one of the group III cations. The solution is green. What ion is most likely present? (see table 4) 9. You have an unknown sample that may contain Cu 2+ or Al 3+. 6 M NH 3 is added to the solution and a blue solution results. Which cation is present? Explain. (see table 5) 10. You have a solution that Co 2+ or Fe 3+. You add KSCN and the solution turns red. Which ion is present? (see table 6)