Handout of Experiment Identification of Unknown Ionic compound.docx

Identification of an Ionic Compound

Scenario: Our hapless chemistry professor Dan is once again in trouble. This time, he is being questioned by the police for sending his bitter rival Ben an envelope containing another unknown white powder. Dan claims that he was just joking, and the powder is simply table salt, NaCl. Initial tests on the powder do reveal that it is an ionic compound of some sort. Your job will be to determine the actual identity of Dan’s mystery powder.

Reading: This lab is based on Project 11 in “Cooperative Chemistry Laboratory Manual, 5^th edition” by Melanie Cooper (2012).

Goals: In this experiment, your group will be given an unknown ionic compound. Your job is to a) determine the identity of your compound, b) investigate the physical and chemical properties of your compound, and c) synthesize a 2.0 g sample of this compound. In week 1 of this experiment, your group should aim to identify your unknown. In week 2, you will synthesize the unknown and test the accuracy of your synthesis. In the week 3 of the lab, your group will prepare a report that summarizes your findings.

Background: You will utilize two main classes of chemical reactions in this project: precipitation reactions and acid/base neutralization reactions. All compounds can be identified and synthesized by a combination of these classes of reactions.

Precipitation reactions involve mixing two aqueous solutions of soluble salts. The cations and anions of the dissolved salts mix, and any insoluble products will precipitate from solution as a solid. The solid can then be filtered out from the solution (the filtrate). For example, when aqueous silver(I) nitrate and calcium chloride are mixed, a white solid is produced. This solid is silver(I) chloride. Calcium and nitrate ions remain dissolved in the solution.

2 AgNO_{3 (aq)} + CaCl_{2 (aq)} 🡪 2 AgCl _(s) + Ca(NO₃)_{2 (aq)}

In general, any salt with an alkali metal or ammonium cation will be soluble. Similarly, all nitrate and acetate salts are soluble. More complete lists of soluble and insoluble salts can be found in your Chem 121 textbook or on the internet.

Acid-base reactions in aqueous solution involve the transfer of a hydrogen cation (a proton, H⁺) from a proton-donor molecule (the acid) to a proton-acceptor (the base). Some acids (e.g. HCl, HBr, HNO₃, H₂SO₄, HClO₄) are called “strong acids”, while others (e.g. acetic acid, H₃PO₄, citric acid) are called “weak acids”. You can think of strong acids as molecules that really want to give away a proton. If you add HCl to water, the HCl will very quickly donate its protons to water molecules, resulting in a solution of chloride anions and hydronium cations, H₃O⁺. Weak acids on the other hand are willing to donate a proton, though they do so to a lesser degree. Similarly, strong bases (e.g. NaOH, KOH) very much want to accept a proton. Weak bases (e.g. NH₃) are willing to accept a proton if necessary (from a strong acid). Some molecules (such as water) can act as both acids and bases in chemical reactions. These molecules are termed amphoteric.

When an acid and base are combined, a neutralization reaction takes place. If the base is a hydroxide (and thus a strong base), then the products of the neutralization are water and a salt. For example, acetic acid (vinegar) and lithium hydroxide combine to form lithium acetate and water.

CH₃CO₂H _(aq) + LiOH _(aq) 🡪 LiCH₃CO_{2 (aq)} + H₂O _(l)

One common example of a base you may encounter in this experiment is the carbonate anion, CO₃²⁻ (or bicarbonate, HCO₃⁻). Carbonates can combine with acids to form CO₂ gas and water.

CO₃²⁻ _(aq) + 2 HCl _(aq) 🡪 CO_{2 (g)} + H₂O _(l) + 2 Cl⁻

One final class of reactions that you might see in this experiment are decomposition reactions. For example, potassium bicarbonate will decompose if heated to form potassium carbonate, carbon dioxide gas, and water vapor:

2 KHCO_{3 (s)} + heat 🡪 K₂CO_{3 (s)} + CO_{2 (g)} + H₂O_(g)

Reaction stoichiometry: Given a balanced chemical reaction and a known mass of starting reactants, you can calculate the mass you should expect to produce of a product. For example, you mix 0.456 g of solid AgNO₃ into water and then add 10.0 mL of 0.500 M CaCl₂. What mass of silver chloride should you collect by filtration?

moles AgNO₃ = 0.456 g / 169.87 g/mol = 0.00268 mol AgNO₃

moles CaCl₂ = 0.500 M × 0.0100 L = 0.00500 mol CaCl₂ (note that 1 M = 1 mol/L)

Since 0.00500 mol / 1 > 0.00268 mol / 2, the silver nitrate is the limiting reagent. (divide the moles of each reactant by the stoichiometric coefficient)

moles AgCl = 0.00268 mol AgNO₃ × 1 AgCl / 1 AgNO₃ = 0.00268 mol AgCl

mass AgCl = 0.00268 mol × 143.32 g/mol = 0.384 g AgCl

You will need to use stoichiometry in order to plan your synthesis of your unknown compound. Your goal is to synthesize 2.0 g of the salt.

List of provided compounds:

Samples of all possible unknowns + sodium bicarbonate, NaHCO₃

Aqueous and solid bases: KOH, NaOH, Mg(OH)₂, Ca(OH)_2­, NH₃ / NH₄OH

Aqueous acids: HCl, H₂SO₄, acetic acid, HNO₃

Other aqueous solutions: AgNO₃, BaCl₂, FeSO₄

Organic solvents: ethanol, acetone

Possible unknowns: There are 25 possible unknowns in this experiment, chosen from the list of cations and anions shown below. Each compound contains exactly one type of cation and one type of anion (e.g. potassium carbonate, ammonium acetate). Your instructor will let you know if we are missing any of these 25 possible salts.

Available cations: sodium, potassium, calcium, magnesium, ammonium

Available anions: chloride, carbonate, sulfate, nitrate, acetate

Safety: Barium compounds must be disposed of in the provided waste container in the fume hood! All other solutions and solids can be disposed of in the sink / trashcan. Note that some nitrate compounds can react with certain strong acids to produce toxic NO₂ gas. If you see a brown or orange gas being produced, immediately take your reaction beaker to the fume hood. Do not breathe in the brown gas! You must wear goggles, a lab coat, and gloves during this lab, and please wash your hands before leaving.

Tests for Anions and Cations: The following are a variety of reactions that you might use to test for the presence of particular anions and cations in an ionic sample. In general, you should use small quantities of materials in these tests. Note that if a test is ambiguous (or you don’t know how to identify a positive result), then you should perform the test on a known compound first to provide a point of reference.

Solubility tests: Not all ionic compounds are soluble in water (or ethanol or acetone). A list of soluble and insoluble salts is shown below. You can thus use this solubility data to narrow down the options for your unknown. For example, if your compound is soluble and you know it contains a phosphate anion, then the cation must be an alkali metal or the ammonium cation. Note that no compound is infinitely soluble. In order to test solubility, add a very small amount of your salt to a beaker of DI water. Stir and let sit for several minutes before deciding a salt is insoluble.

Always soluble: Alkali metal cations Ammonium cations, NH₄⁺

(in water) Nitrates, NO₃⁻ Acetates, C₂H₃O₂⁻

Chlorates, ClO₃⁻ Perchlorates, ClO₄⁻

Usually soluble: Chlorides, bromides, and iodides (but not with Ag⁺, Pb²⁺, Hg₂²⁺)

Fluorides (not with Pb^2+, alkaline earth metals)

Sulfates, SO₄²⁻ (not with Ca²⁺, Sr²⁺, Ba²⁺, Pb²⁺)

Usually insoluble: Carbonates, CO₃²⁻ (except with alkali metals and NH₄⁺)

Phosphates, oxalates, and chromates (except as with CO₃²⁻)

Hydroxides, OH⁻ (except alkali metals, NH₄⁺, and Mg²⁺)

The solubility (in g/L) of various salts can be found online or in the CRC Handbook. You can determine the solubility of your compound by carefully and quantitatively adding your salt to a known volume of hot water until no more dissolves. Allow the water to cool, filter out and weigh any undissolved solids, and then calculate the solubility.

pH test: You can test the pH (acidity or basicity) of your unknown by first dissolving it in water and then touching the solution to pH paper.

Flame tests for cations: Many metal cations produce bright colors when strongly heated in a flame. To test your sample, first dissolve it in DI water. (If it won’t dissolve in water, then try dissolving it in dilute hydrochloric acid.) Then light a Bunsen burner and obtain a clean nichrome wire. (Heat the wire in the flame for a while to clean it.) Dip the wire in your sample solution and place it in the hottest part of the flame. You may observe a brief flash of color due to your sample. (The orange color that stays around for many seconds is due to the wire itself and not your sample.)

Note that sodium produces a very, very bright yellow-orange light. This yellow light can make it difficult to detect certain cations, especially potassium. To counteract this sodium color, you can observe your flame test through a piece of blue glass. This blue glass will filter out the yellow color, making it easier to see other colors.

Table 1: Flame test colors and intensities

Test for ammonium cation: Dissolve your sample in water to create an aqueous solution. Mix roughly 2 mL of your sample with 2 mL of 6M NaOH. Mix well and smell the resulting mixture. If ammonium cations are present, then you should smell ammonia. A piece of pH paper held above the mixture should also turn blue, indicating the presence of a base (ammonia gas). (Note that if you aren’t getting a scent but you suspect that you have ammonium cations, then you might try adding some NaOH to your solid sample. If you don’t know what ammonia smells like, then try this test on a known ammonium salt.)

Tests for anions: All of the following tests (except for the carbonate test) require you to first dissolve your sample in DI water. If your sample won’t fully dissolve, then use a dilute slurry of your sample.

Test for chloride anions: Mix 1 mL of your sample with 1 mL of 6M nitric acid, HNO₃. Then add 1 mL of silver nitrate, AgNO₃. A white precipitate indicates the presence of either chloride anions or sulfate anions. (A yellow precipitate will form if bromide or iodide anions are present.)

AgNO_{3 (aq)} + Cl⁻_(aq) 🡪 AgCl _(s) + NO₃⁻_(aq)

Test for sulfate anions: Mix 1 mL of your sample with 1 mL of 6M HCl and 1 mL of BaCl₂ solution. A white precipitate (BaSO₄) indicates the presence of sulfate anions. Note that barium cations should not go down the drain! Please dispose of all mixtures containing barium in the appropriate waste container.

BaCl_{2 (aq)} + SO₄²⁻_(aq) 🡪 BaSO_{4 (s)} + 2 Cl⁻_(aq)

Test for carbonate anions: Add a small scoop of your solid sample to a watch glass, and then add several drops of 6M HCl. If your sample effervesces (bubbles or foams) then you have carbonate anions.

2 HCl _(aq) + CO₃²⁻_(aq) 🡪 CO_{2 (g)} + H₂O _(l) + 2 Cl⁻_(aq)

Test for nitrate anions: Place 1 mL of your sample in a test tube in the fume hood, and then carefully add 3 mL of concentrated sulfuric acid, H₂SO₄. (Be careful! Sulfuric acid can and will burn you and make holes in your clothing!) Mix well and allow it to cool. Tip the test tube at a 45° angle and then slowly add 2 mL of FeSO₄ solution so that this latter solution floats on the sulfuric acid / sample mixture. A brown ring will form at the interface of these solutions if nitrate anions are present. (Note that it sometimes takes a while for the ring to form. I suggest you set up the test tube and then allow it to sit undisturbed for up to 30 minutes. If a ring hasn’t formed after this time, then you don’t have nitrate anions.)

Test for acetate anions: Add 2 mL of your sample to a test tube, and add 1 drop of concentrated sulfuric acid (be careful!). Add 1 mL of ethanol, and heat the solution for several minutes in a hot water bath. If acetate is present, then you should smell a fruity odor due to the presence of ethyl acetate. (I strongly suggest that you first try this test on a known acetate compound so that you know what you should be smelling!)

C₂H₃O₂⁻ + C₂H₅OH 🡪 CH₃CO₂C₂H₅ (ethyl acetate)

Procedure: As noted in the goals, you have three objectives in this lab.

Objective 1: Identify your unknown compound

• Using tests for anions and cations, carefully determine the identity of your unknown salt. You will likely need to repeat some tests if the result is unclear.
• Perform all tests on your unknown. If you skip tests, it is very easy to make a mistake and arrive at an incorrect conclusion!
• Once you think you have identified your salt, perform the tests on a sample of the known compound. Do you get the same results?

Objective 2: Describe the properties of your unknown

• You can test properties such as solubility, pH, and reactivity to acid in the lab.
• You should also research your compound and learn as much about it as possible. What are the physical and chemical properties of your compound? What is it used for in industry or biology? Be sure to cite all outside sources you use.

Objective 3: Synthesize a 2.0 g sample of your compound

• Each compound can be synthesized using a combination of precipitation, decomposition, and/or neutralization reactions.
• Use a stoichiometric calculation to determine how much of each reagent you need. You should include these calculations in your notebook, and you should also calculate the percent yield of your compound.
• Once you have your synthesized compound, how do you know that you made the correct material? Is your compound a mixture of salts? To test this, you should perform the cation and anion tests on your compound!