Testing Anions

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Goals
This lab introduces qualitative analysis, the area of chemistry concerned with the identification of substances by their chemical reactions. You will observe the reactions of some simple salts, analyze common household chemicals, and identify an unknown sample by testing its reactivity. This experiment also provides a review of balancing chemical equations and reaction stoichiometry.


Introduction

Identifying samples of unknown substances is an important part of chemistry, with applications in such fields as medicine, environmental science, and geology. Materials can be characterized by their chemical reactions and by methods involving instruments. Often identification of substances by their typical reactions is easy, quick, accurate, and inexpensive, in comparison to instrumental methods. In this lab you will observe some characteristic reactions of five negatively charged ions (anions). Your observations of carbonate (CO₃^-2), hydrogen carbonate (HCO₃^-), sulfate (SO₄^-2), chloride (Cl^-), and iodide (I^-) reactions will then help you to identify an unknown sample.

Chemistry & Background
It is important to understand the structure of the five anions, as you observe their reactions and balance the corresponding chemical equations. Chloride and iodide ions are formed when a chlorine or iodine atom gains an extra electron, to form a negatively charged ion. These anions form ionic compounds, or salts, with positively charged cations, often metal ions, such as sodium, Na⁺, as in sodium chloride, NaCl. The formula of an ionic compound is the ratio of cation and anion that gives a balanced, neutral charge. Carbonate, hydrogen carbonate, and sulfate ions are polyatomic ions. They consist of groups of atoms covalently bonded together, with an overall charge. They also form ionic compounds by the attraction of an oppositely charged ion, like sodium or calcium, Ca²⁺, as in calcium carbonate, CaCO₃. Polyatomic atoms are discussed in your textbook in section 2.9.

The first step in writing chemical equations for aqueous ionic reactions is to recognize the anions and cations and write them as dissolved species. In aqueous solution, soluble ionic compounds separate into cations and anions, which are surrounded by water molecules. For example, sodium chloride dissolves in water to form sodium cations, Na⁺, and chloride anions, Cl^-. Ionic compounds containing polyatomic ions also dissolve in water to form separate anions and cations, but the polyatomic ions remain intact. For example, sodium hydrogen carbonate dissolves in water to form sodium cations, Na⁺, and hydrogen carbonate anions, HCO₃^-. When writing chemical equations for the reactions observed, we generally write the net ionic equation, which shows only those ions that react, not those which are spectator ions. For more information on balancing chemical equations, see the following section.

A Note on Balanced Chemical Equations
All chemical equations in your lab notebook and write-up should be properly balanced. This means that each side of the equation should have the same number of moles of each element, so that matter is not created or destroyed. In addition, the electrical charges must be the same on both sides of the equation. When writing chemical equations for ionic species, one generally writes the net ionic equation which shows only those species which react. Ions which do not react, but serve to balance the electrical charge of reacting ions, are called spectator ions.

For example, when a solution of sodium chloride reacts with silver nitrate, the balanced equation can be written in complete form asNaCl(aq) + AgNO₃(aq)  AgCl(s) + NaNO₃(aq)
We can break each dissolved ionic compound into cations and anions to give

Na⁺ (aq) + Cl^- (aq) + Ag⁺ (aq) + NO₃^- (aq) 
AgCl(s) + Na⁺ (aq) + NO₃^- (aq)

Note that the insoluble precipitate, AgCl, is not broken down into separate ions, since it is not disolved. The sodium and nitrate ions are spectators to the precipitation reaction between silver and chloride ions. If we remove the spectator ions, the net ionic equation can be written

Cl^- (aq) + Ag⁺ (aq)  AgCl(s)

For more on balancing chemical equations, refer to the sections of your textbook listed on the Introduction and Goals page.

Anion Reactions
The following paragraphs present the chemical reactions that you will observe in the lab this week. Compounds containing the carbonate ion (or the related ion bicarbonate or hydrogen carbonate, HCO₃^-) react with acids to give carbon dioxide (CO₂), a colorless, odorless, and slightly acidic gas. For example, the reaction of baking soda (NaHCO₃) with acids such as those in lemon juice or buttermilk releases carbon dioxide which causes baked goods to rise:

HCO₃^- + H⁺  CO₂ (g) + H₂O

In this equation the sodium ion, Na⁺, does not participate in the reaction and is a spectator ion. It is not included in the net ionic equation, given above.

Many other gases are colorless and odorless, like CO₂. To confirm that the gas formed is carbon dioxide, it can be allowed to react with Ba(OH)₂, barium hydroxide, to form a white insoluble precipitate of barium carbonate, BaCO₃:

CO₂ + Ba²⁺ (aq) + 2 OH^- (aq) BaCO₃ (s) + H₂O

Eggshells, seashells, chalk, and limestone are some of the common substances containing calcium carbonate.
Chloride salts react with sulfuric acid to make hydrogen chloride gas, which is also an acid:

Cl^- + H⁺ (aq) HCl (g).

In this reaction, the sulfate ion, SO₄^-2, from the sulfuric acid, is a spectator ion and is not included in the net ionic equation.

Another useful reaction for identification of the chloride ion is its reaction with silver nitrate to form a white precipitate of silver chloride:

Cl^- + Ag⁺ (aq)  AgCl (s)

As discussed above, nitrate and sodium are spectator ions and do not appear in the net ionic equation.

To test the properties of the sulfate ion, you will examine the reactions of Epsom salts. This salt, MgSO₄^.7H₂O, is used as a purgative and, as an aqueous solution, to soak tired feet. Since it dissolves in water to form Mg²⁺ and SO₄^2- ions, it will not react with a solution containing SO₄^2- and HSO₄^- ions, such as H₂SO₄. It will, however, form insoluble precipitate of BaSO₄when in the presence of Ba²⁺ ion:

Ba²⁺ + SO₄^2-  BaSO₄ (s)

Iodide salts can be identified by their reaction with chlorine (Cl₂). Chlorine is a toxic, pale, yellow-green gas with an irritating odor and low solubility in water. Its structure consists of two chlorine atoms covalently bonded together, with a single bond of two shared electrons. A convenient laboratory source of chlorine is commercial bleach, which is usually a 5 percent aqueous solution of sodium hypochlorite, NaOCl. The solution behaves as if it contains dissolved chlorine. This is because hypochlorite and chloride ions are in equilibrium with chlorine and hydroxide ions:

OCl^- (aq) + Cl^- (aq) + H₂O  Cl₂ (aq) + 2 OH^- (aq)

Note that sodium is again a spectator ion and does not appear in the net ionic equation. In this experiment, you will use a solution of bleach to supply Cl₂ (aq) for reaction with iodide ion. The products of this reaction are molecular iodine and chloride ion.

Cl₂ + 2 I^-  I₂ + 2 Cl^-

Iodine (I₂) exists as a purple solid and dissolves sparingly in water. When iodine (I₂) is in solution with iodide ion (I^-), the brown triiodide ion forms:

I₂ (aq) + I^- (aq)   I₃^- (aq)
Purple  colorless  brown

Thus, reaction of iodide ion with a solution of bleach produces iodine, which then reacts with remaining iodide ion to form the brown triiodide ion. This brown color will help you to recognize the presence of iodide and iodine in your reaction solutions.

A second test for iodide ion is formation of a silver iodide precipitate. Like the chloride ion, iodide reacts with silver nitrate solution to give a precipitate, although AgI is yellow, while AgCl is white:

Ag⁺ + I^-  AgI (s)

In this reaction, sodium and nitrate are spectator ions and are not included in the net ionic equation.

Yet another test for iodide is the reaction of an iodide salt with sulfuric acid to form a brown solid and cause gas evolution. In this series of reactions, iodide ion is first oxidized to I₂. The gas formed is hydrogen sulfide (H₂S), from the reduction of sulfuric acid. Since both iodide ion (I^-) and molecular iodine (I₂) are present, the triiodide ion forms, giving the characteristic brown color.


8 I^- (aq) + 10 H⁺ (aq) + SO₄^2- 
H₂S (g) + 4 I₂(aq) + 4 H₂O

I₂ (aq) + I^- (aq)  I₃^- (aq)

A final test for the presence of iodide ion can be performed using starch to detect I₂, produced from I^- in one of the above reactions. Starch contains amylose, a polymer of a-D glucose that reacts to form a characteristic blue-black complex with I₂ in the presence of I^-. The iodine and iodide form an I₅^- chain which fits inside the helix formed by the sugar chains. This complex is a dark blue-black that is easily observed. Starch can be used to confirm the results of the reaction between iodide and chlorine or iodide and H₂SO₄, since both reactions produce with I₂ in the presence of I^-. 

Key Questions Here are some questions to think about before, during, and after this week's experiment: 1. What happens when an ionic compound, like NaCl, dissolves in water? 2. What is the logical scheme for identifying each anion? Can you draw a flowchart for the identification of each? 3. Would it be possible to identify each of the four anions if they were mixed together? How? Or why not? 4. What is the general class of each test reaction? Precipitation? Acid-Base? Oxidation-Reduction?

Prelab Problems
1. Complete and balance the following equation:

CH₃CO₂H(aq) + NaHCO₃(aq) -->

2. Complete and balance the following equation:

BaCl₂(aq)+ H₂SO₄(aq) -->


3. Complete and balance the following equation:

AgNO₃(aq) + NaI(aq) -->

4. Draw a flowchart for the analysis of an unknown sample. Start with the addition of H₂SO₄ to the sample and draw arrows for the possible results and conclusions that can be drawn. Include confirmation tests for each possible ion and balanced chemical equations. Below is a partial chart to help you get started.




5. If you had gotten a single solid sample in your test tube, but couldn't remember if it was NaCl or Na₂CO₃, how could you test the powder and decide which substance was present?

6. Use the web applet Qualitative Analysis of Anions to test a web-based unknown sample. Practice with known samples and unknowns that you can reveal before trying a "true unknown". Follow the procedure and use your flowchart from prelab problem 4. In your lab notebook, record the code number of your "true unknown" and its identity for your TA to grade. Like all prelab problems, this one should be done independently, not in collaboration with a lab partner or friend. You can access the applet by clicking on the link below.
Qualitative Analysis of Anions Applet

Procedure

TheCarbonate Ion, CO₃^-2 and Hydrogen Carbonate Ion, HCO₃^-

Place a pea-sized amount of solid baking soda in a small clean test tube that has been rinsed with distilled water. Add to it 1 or 2 drops of 18 M sulfuric acid (H₂SO₄) (CAUTION! STRONG ACID) and record your observations. Confirm the presence of CO₂ by dripping a drop of Ba(OH)₂ solution down the side of the test tube, as the gas bubbles are forming. Watch carefully for precipitate in the drop of solution running down the side of the test tube. This delicate procedure is illustrated on the ChemLab website, in the Procedure Overview section. Write a balanced chemical equation and describe your observations.

Repeat this experiment using vinegar in place of the sulfuric acid. Vinegar contains acetic acid, CH₃CO₂H. Write balanced chemical equations to describe your observations.

Test common blackboard chalk for this ion as follows: place a small piece of chalk in a dry test tube and add a few drops of 2 M HCl. Test the escaping gas by carefully running a drop of barium hydroxide solution from an eye dropper down the inner wall of the tube. Record your observations and write balanced chemical equations.

The Chloride Ion Cl^-
Put an amount of sodium chloride the size of a small pea in a small, clean test tube that has been rinsed with distilled water. Take the tube to the fume hood and carefully add 1 or 2 drops of 18 M sulfuric acid. (CAUTION, avoid breathing the gas formed). Carefully lower a piece of moist pH indicator paper into the tube, as the gas is evolved, to get more information about the gas. Indicator paper is impregnated with a colored compound which is sensitive to acids and bases. It is used to categorize solutions and gases as strongly acidic (red), neutral (no significant color change), or basic (blue). Record your observations and write balanced chemical equations.

Using a similar amount of NaCl in a new small test tube, add 15 drops distilled water and one drop of 3 M nitric acid (CAUTION! STRONG ACID). Then add 3-4 drops of 0.1 M silver nitrate and mix the contents. What happens? Answer with a balanced chemical equation. The purpose of the nitric acid, HNO₃, in this reaction is to prevent the precipitation of undesired silver salts, like AgOH, which can occur with non-acidic conditions. It is not a reactant in the balanced net ionic equation, for this reaction of chloride ion and silver nitrate.

Now test Hanover tap water for the presence of chloride ion. Place about 2 mL of tap water (without NaCl) into a clean test tube and add 1 drop of 3 M HNO₃. Then add the silver nitrate. Look carefully for precipitate. You may wish to look through the test tube the long way, through the opening, and compare to a test tube containing only distilled water. What can you conclude about the presence of chloride ions in tap water? Write a balanced chemical equation to describe your observations.

The Sulfate Ion SO₄^-2
Place a pea-sized amount of Epsom salts in a small, clean test tube that has been rinsed with distilled water and add 1-2 drops of 18 M sulfuric acid. Compare your observations with those made above for the carbonate ion and record them in your notebook. Write a chemical equation to describe your observations.

Dissolve a similar amount of Epsom salts in about 1 mL of distilled water in a new test tube. Add 1 drop of 3 M nitric acid and then 1-2 drops of 0.2 M BaCl₂. Compare the result of this experiment to those seen above for the carbonate ion. Write balanced chemical equations for your observations.

The Iodide Ion I^-
To test for the iodide ion, dissolve a pea-sized amount of sodium iodide (NaI) in a test tube in 1 mL of distilled water; then add 5 drops of bleach (CAUTION, avoid getting bleach on your skin). Write a balanced chemical equation to match your observations. Confirm the presence of I₂ and I^- by observing the reaction between a corn starch packing peanut and the reaction mixture in your test tube. Tear off a piece of a packing peanut and push it into the solution in your test tube with a stirring rod. Record your observations.

As a second test, dissolve a small amount of NaI in 1 mL of distilled water and add a drop of 3 M HNO₃, then 3-4 drops of 0.1 M silver nitrate solution. Record your observations and write a balanced chemical equation.

For a third test, place a very small amount of NaI in a dry test tube. In the fume hood, carefully add 1-2 drops of 18 M sulfuric acid. (CAUTION, avoid breathing the gases formed). Again, record your observations and write a balanced chemical equation. Confirm the presence of I₂ and I^- by observing the reaction between a corn starch packing peanut and the reaction mixture in your test tube. First add about 1 mL of distilled water. Then tear off a piece of a packing peanut and push it into the solution in your test tube with a stirring rod. Record your observations.

Testing for Carbonate Ion, CO32-
		Baking soda is sodium hydrogen carbonate or sodium bicarbonate. It reacts with acid the same way that carbonate ion does.
H+ (aq) + HCO3- (aq)  H2O (l) + CO2 (g)
		Chalk is calcium carbonate, CaCO3. 2 H+ (aq) + CO32- (aq)  H2O (l) + CO2 (g)
		Reaction with Ba2+ ion is a test for CO2 gas. CO2 (g) + Ba2+ (aq) + H2O (l)   BaCO3 (s) + 2 H+(aq) The barium carbonate precipitate is visible in the drop of Ba(OH)2 solution, running down the side of the test tube.
Testing for Chloride Ion, CI-
		The source of chloride ion in this experiment is NaCl, sodium chloride or table salt.
Cl- (aq) + H+ (aq)  HCl (g)
		To test for HCl gas, put wet pH paper in the test tube, as the gas is evolved. HCl is an acid and should change the color of the pH paper accordingly.
Another test for chloride ion is reaction with silver nitrate, AgNO3. Nitric acid is added to aid the precipitation.
		Cl- (aq) + Ag+ (aq)  AgCl (s)
		The precipitation of AgCl may be less obvious for tap water, but is chloride ion present?
Testing for Sulfate, SO42-
		The sulfate in this experiment is supplied by magnesium sulfate, MgSO47H2O, or Epsom Salts.
		Sulfate anion plus sulfuric acid SO42-(aq) + H+ (aq)  ??
Add nitric acid (HNO3) and barium chloride (BaCl2)
		SO42- (aq) + Ba2+ (aq)  BaSO4(s)
Testing for Iodide, I-
		Sodium iodide, NaI, is the source of iodide anion for this experiment.
		Reaction with bleach involves three steps. The brown color shows the presence of I3- ions. Hypochlorite ion yields chlorine: OCl- (aq) + Cl- (aq) + H2O   Cl2 (aq) + 2 OH- Chlorine reacts with iodide anion: Cl2 (aq) + 2 I- (aq)   I2 (aq) + 2 Cl- (aq) Triiodide ion is formed: I2 (aq) + I- (aq)  I3- (aq)
		Starch reacts with iodine and iodide to form a characteristic blue/black complex. A corn starch packing peanut is shown here.
Silver ion reacts with iodide to form silver iodide, AgI.
		Ag+ (aq) + I- (aq)  AgI (s)
Reaction of iodide with sulfuric acid produces hydrogen sulfide gas and brown triiodide solution in a series of reactions: I- (aq) + H2SO4 (aq)  HI (aq) + HSO4- (aq) 8 HI + H2SO4 (aq)  H2S (g) + 4 I2 (aq) + 4 H2O I2 (aq) + I- (aq)  I3- (aq)
		Again, starch is used to confirm the presence of iodine and iodide.
Read More http://www.dartmouth.edu/~chemlab/chem3-5/qual_an/overview/procedure.html