Reactivity: Precipitation, Acid-Base, and Redox Reactions

Adapted in part with permission from an experiment developed by Professor CJ Scaife, Union College, Schenectady, NY.

• To explore reactions of cations and anions in aqueous solution
• To devise solubility rules
• To explore acid-base reactions in aqueous solution
• To explore redox reactions in aqueous solution

Precipitation Reactions A reaction of two water soluble salts in which cation and anion partners are "traded" is called a metathesis or double displacement reaction. Such a reaction is represented schematically below, where A and B are two different cations, X and Y two different anions.

For such a reaction to proceed to the right, there must be a driving force. The driving force is provided, according to Le Chatelier's principle, by removing one of the products from solution. This will occur if one of the products precipitates (i.e., leaves solution as a solid), is a weak electrolyte (i.e., is essentially unionized in solution), or decomposes to form a gas, which escapes from solution. Reaction of two salts to form solid, weak electrolyte, or gas is spontaneous, meaning that it occurs without outside assistance.

The formation of a precipitate is illustrated in equation form below. Three equations, of increasing simplicity, are shown.

These are called, respectively, the total equation (TE), the ionic equation (IE), and the net ionic equation (NIE). The TE shows what reagents to obtain from the stockroom to carry out a given reaction. An IE is obtained by writing all soluble strong electrolytes from the TE in dissociated form; it bridges the total equation and the net ionic equation. The net ionic equation is obtained by omitting ions that appear on both sides of the ionic equation. Canceled ions are called spectator ions because they do not take part in the chemical reaction. Therefore, the NIE shows in simplest form what species take part in the reaction.

Solubility is the amount of solute (in g or moles) that dissolves in a given amount of solvent (in mL, L, or g). If quite a bit of a solid dissolves in a solvent, the solid is soluble in that solvent. If only a small amount of a solid dissolves in the solvent, the solid is insoluble, or sparingly soluble, in the solvent. A solubility of 1.0 gram solute per 100 mL solvent is considered the border line between soluble and insoluble. For example, 74.5 g of calcium chloride, CaCl₂, dissolves in 100 mL H₂O; thus CaCl₂ is soluble in water. In contrast, only 2.9 x 10^-4 g silver chloride, AgCl, dissolves in 100 mL water, so it is insoluble in water.

Acid-Base Reactions A compound that contains a hydrogen ion, H⁺, in combination with an anion is called an acid. Examples of common acids are HCl (hydrochloric acid), HBr (hydrobromic acid), HNO₃ (nitric acid), H₂SO₄ (sulfuric acid), and H₃PO₄ (phosphoric acid). A compound that contains a cation in combination with the hydroxide ion, OH^-, is called a base. Examples of common bases are NaOH (sodium hydroxide), KOH (potassium hydroxide), and Ca(OH)₂ (calcium hydroxide). Other compounds that do not contain hydroxide nonetheless behave as bases in the presence of acids. Examples of such compounds are sodium carbonate, Na₂CO₃, sodium hydrogen carbonate, NaHCO₃, sodium phosphate, Na₃PO₄, ammonia, NH₃, and many others. These substances behave as bases because their anions react with the hydrogen ion of acids to form weak electrolytes.

Reaction of an acid and a base always involves the combination of the hydrogen ion from the acid with a basic anion. The anion is OH^- for the common bases. For the other bases above, the anion varies from substance to substance. In the case of ammonia, the molecule itself combines with the hydrogen ion to produce the ammonium ion, NH₄⁺. Several acid-base reactions are shown below:

All of these reactions can be written in simpler form as net ionic reactions. In (5), Cl^- and Na⁺ are spectator ions and can be omitted, leaving the net ionic equation in (9):

You should be sure you understand that equation (9) is also the net ionic equation for reactions (6) and (7). The net ionic equation for reaction (8) is given in (10):

Redox Reactions Reduction-oxidation, or redox, reactions involve a transfer of electrons from one substance to another. For example, reaction (11) involves the transfer of electrons from sodium metal to chlorine gas. The result is table salt, NaCl, in which sodium is present as the Na⁺ cation, and chlorine is present as the Cl- anion.

One electron is transferred per atom of sodium that reacts, so two electrons are transferred in process (11). Similarly, reaction (12) involves the transfer of electrons from iron metal to oxygen gas.

Based on the formula for the product, we conclude that iron is present as the Fe³⁺ cation. Thus 3 electrons are transferred per iron atom reacted, for a total of 12 electrons transferred in all. These 12 electrons are accepted by six oxygen atoms, each of which becomes an oxide ion, O^2-.

When an atom loses electrons, we say that it has been oxidized or that it has undergone oxidation. An atom that gains electrons is said to have been reduced or to have undergone reduction. In equation (11), sodium is oxidized (to Na⁺) and chlorine is reduced (to Cl^-). In equation (12), iron is oxidized, and oxygen is reduced.

In many cases it is not so easy to tell from the chemical equation whether or not redox has occurred. For example, equation (13) shows a redox reaction in which some chemical experience is required to identify the atoms involved in losing and gaining electrons:

In your later studies of chemistry, you will be able to determine that the nitrogen atoms in N₂H₄ each transfer 2 electrons to the nitrogen atoms of N₂O₄, each of which accepts 4 electrons. Thus the nitrogen atoms of N₂H₄ are oxidized, and those of N₂O₄ are reduced. However, this is not at all obvious from inspection of the equation!

In this experiment, you will examine reactions of all 3 types discussed above. First you investigate many cation-anion combinations to discover which produce insoluble products by carrying out double displacement reactions that form solids. The cations to be investigated are Na⁺, K⁺, NH₄⁺, Mg²⁺, Ca²⁺, Ba²⁺, Al³⁺, Pb²⁺, Cr³⁺, Mn²⁺, Fe²⁺, Co²⁺, Ni²⁺, Cu²⁺, Zn²⁺, Ag⁺. Solutions of each of these cations with the NO₃^- anion are available in the lab. The anions to be investigated are NO₃^-, Cl^-, Br^-, I^-, CO₃^2-, SO₄^2-, PO₄^3-, PO₄^3-, CrO₄^2-, OH^-, S^2-. Solutions of each of these anions with the sodium cation, Na⁺, are available in the lab.

Second, you will explore a number of acid-base reactions to become familiar with some of the "signals" that such a reaction has taken place.

Finally, you will examine several simple redox reactions in aqueous solution.

To minimize time, cost, and safety hazard, and to maximize the efficiency of carrying out and cleaning up the experiments, you will work on a very small scale (microscale). Based on observation, you will propose answers to the following questions:

1. What cation(s) form only soluble salts with the anions investigated? What cation(s) form only insoluble salts with the anions? What anions form only soluble salts with the cations investigated? What anions form only insoluble salts with the cations?
2. Formulate a set of solubility rules to summarize your experimental observations for the anions. For anions that form primarily soluble salts, a rule might take the form, "All chlorides are soluble except..." For anions that form primarily insoluble salts, a rule might take the form: "All hydroxides are insoluble except..." Group several anions in a single rule when possible.
3. Did you observe an effect in Experiment 1? Describe it. How did the observed effects in Experiments 2 and 3 differ from the effect in Experiment 1?
4. Describe and explain your results in Experiment 4.
5. Describe and explain your results in Experiment 5.
6. Describe and explain your results in Experiment 6.
7. Describe and explain your results in Experiment 7.
8. Describe and explain your results in Experiment 8.

• 8.5 x 11" plastic reaction sheet
• cation/anion grid
• paper clips
• disposable micropipets
• magnifying lens
• Kimwipes
• microstirrers (melting point capillaries)
• wash bottle
• well plates
• red litmus
• 6 M NaOH
• 1 M Na₂S
• solutions containing cations ( + nitrate anion) and anions (+ sodium cation)
• 2 M HCl
• 2 M H₂SO₄
• 2 M H₃PO₄
• 2 M NaOH
• 0.1 M HCl
• 0.1 M NaOH
• sodium hydrogen carbonate
• sodium carbonate
• zinc strips or wire
• copper wire
• silver wire
• iron wire
• calcium turnings

All reagents used in this experiment are to be considered harmful. Wear your goggles. Avoid ingestion of solutions or solids.

Precipitation Reactions Record all data and observations in your lab notebook. Obtain a plastic reaction sheet, a paper grid, two paper clips, and three microstirrers (toothpicks). Wipe the sheet with a Kimwipe to remove any residue (contaminants may cause confusing reactions). Attach the grid to the underside of the plastic sheet with paper clips. Practice delivering individual drops of water from a disposable micropipet until you can do it competently.

Several test tube racks, each containing test tubes with cation and anion solutions, and micropipets for dispensing solution, have been placed in the lab. You will share the solutions and pipets with others, so be quick, careful, and courteous. Do not keep a test tube/micropipet set at your bench for longer than you need it. Start with the cation nitrate solutions shown in the left column of the grid. The cation nitrate solutions were prepared by dissolving the nitrate salt of the cation in water (e.g., Ca(NO₃)₂(s) was used for Ca²⁺). Obtain a test tube containing a solution of any cation except Na⁺ (which will be omitted, since it is already present in all the anion solutions). Use the pipet to place one drop of the cation solution in the appropriate cation circle in the leftmost column on the grid. Place nine more drops in the smaller circles horizontally across the grid. Center the drops in the circles. Return the pipet to the test tube, and the test tube to the rack. Continue with other cation solutions until you have a complete matrix of cation drops (160 total).

Proceed similarly with the sodium anion solutions in the top row of the grid. Sodium anion solutions were prepared by dissolving the sodium salt of the anion in water (e.g., NaBr(s) was used for Br^-). Obtain a test tube containing solution of any anion except NO₃^- (omitted, since it is already present in all cation solutions). Use the pipet to place one drop of the anion solution in the appropriate anion circle in the top row on the grid. Place 16 more drops vertically down the sheet. Place each drop directly on the drop of cation solution. Complete the matrix of anion solutions (153 drops in all). Use only one drop of each solution, and be careful not to touch the tip of the dropper to other solutions on the sheet. Finally, place one drop NaNO₃(aq) centered in the oval in the upper left corner of the grid.

Record observations on the boxed grid (see following pages). Indicate whether solid is formed, and colors of solids and/or remaining solutions. Cloudiness often suggests solid formation. If in doubt about solid formation, view the drop with a magnifying lens and/or stir the drop carefully with a microstirrer. Use a colored highlighter to indicate on the boxed grid where precipitate formed.

UNDER A FUME HOOD, add 6 M NaOH solution dropwise to the 250-mL beaker until the mixture is basic to litmus (3-5 drops required). Add 1 drop excess. Add 10 drops 1 M Na₂S solution to insure complete precipitation of aqueous cations as insoluble sulfide salts. Finally, heat the mixture using a Bunsen burner just to boiling for 3 minutes to coagulate the precipitate. Pour the mixture into the waste bottle in the hood labelled HEAVY METAL MIXTURE. This will later be filtered to remove heavy metal waste prior to disposal. Wash your hands well.

Acid-Base Reactions Obtain about 15 mL of 0.1 M NaOH, 10 mL of 0.1 M HCl, and 10 mL of 2 M NaOH in small beakers. In 3 small test tubes, obtain 1 mL each of 2 M HCl, 2M H₂SO₄, and 2 M H₃PO₄. Carry out the following experiments using supplied reagents.

Experiment 1: Reaction of HCl with NaOH. Transfer about 3 mL of 2 M NaOH into the small test tube that contains 1 mL of 2 M HCl, and STIR with a glass stirring rod. Record observations and write a net ionic equation for the reaction. Place the test tube in a small beaker so that you may compare it with subsequent experiments.

Experiment 2: Reaction of H₂SO₄ with NaOH. Transfer about 3 mL of 2 M NaOH into the small test tube that contains 1 mL of 2 M H₂SO₄ and STIR. Record observations. How do your results compare with results of Experiment 1? Write a net ionic equation for the reaction. Place the test tube in the beaker with the test tube from experiment 1.

Experiment 3: Reaction of H₃PO₄ with NaOH. Transfer about 3 mL of 2 M NaOH into the small test tube that contains 1 mL of 2 M H₃PO₄ and STIR. Record observations. How do your results compare with results of Experiments 1 and 2? Write a net ionic equation for the reaction. Hold the test tubes from all 3 experiments in your hand. What do you observe?

Experiment 4: Reaction of HCl with Carbonates. Weigh 0.1 mmole of sodium hydrogen carbonate (baking soda, NaHCO₃) and place it in a small test tube. Carefully add 0.1 M HCl dropwise, counting drops, until reaction stops (how will you know?). Then weigh 0.1 mmole of calcium carbonate (limestone/chalk, CaCO₃) and place it in a small test tube. Carefully add 0.1 M HCl dropwise with stirring, counting drops, until reaction stops. Discard reaction solutions in the sink. Write net ionic equations for the reactions.

Experiment 5: Reaction of HCl with insoluble metal hydroxides. Into a small test tube, transfer about 0.5 mL of copper nitrate (Cu(NO₃)₂) solution and add 1 mL of 0.1 M NaOH. STIR. Record observations. Then add 0.1 M HCl dropwise from a Pasteur pipet with stirring until reaction is complete (how will you know?). Record observations then discard reaction solution in the heavy metal waste. Write a net ionic equation for the reaction.

Into a small test tube, transfer about 0.33 mL of aluminum nitrate (Al(NO₃)₃) solution and add 1 mL of 0.1 M NaOH. STIR. Record observations. Then add 0.1 M HCl dropwise from a Pasteur pipet, with stirring, until reaction is complete. Record observations then discard reaction solution in the heavy metal waste. Write a net ionic equation for the reaction.

Into a small test tube, transfer about 0.33 mL of aluminum nitrate solution and add 1 mL of 0.1 M NaOH. Record observations. Then add 0.1 M NaOH dropwise from a Pasteur pipet, with stirring, until reaction is complete. Record observations then discard reaction solution in the heavy metal waste. Write a net ionic equation for the reaction.

Redox Reactions Obtain a 12-well plate and a magnifier from the instructor. Carry out the following experiments in the wells of the plate.

Experiment 6: Reaction of zinc with copper ion. Place about 10 drops of 0.2 M copper nitrate solution in a well. Submerge the end of a piece of zinc strip or wire and observe. When reaction is complete, remove the zinc strip and examine it. Record all observations.

Now place 10 drops of 0.2 M zinc nitrate (Zn(NO₃)₂) solution in a well and submerge the end of a piece of copper wire. Observe for 1 minute. Record all observations.

Experiment 7: Reaction of metals with hydrochloric acid. Place about 10 drops of 2 M HCl in a well. Submerge the end of a piece of zinc strip or wire and observe. When reaction is complete (how will you know?), remove the zinc strip and examine it. Record all observations.

Repeat the experiment using a small piece of calcium instead of zinc. Record observations.

Repeat the experiment using a small piece of iron. Record observations.

Repeat the experiment using a small piece of copper. Record observations.

Experiment 8: Reaction of silver and copper. Place about 10 drops of 0.2 M silver nitrate (AgNO₃) solution in a well. Submerge the end of a piece of copper wire and observe with the magnifier. When reaction is complete, remove the copper wire and examine it. Record all observations.

Place about 10 drops of 0.2 M copper nitrate solution in a well. Submerge the end of a piece of silver wire and observe with the magnifier. After a few minutes, remove the silver wire and examine it. Record all observations.

Disposal Methods

Empty the wells into a beaker, and dump the beaker contents in the heavy metal waste.

1. C. Scaife, Union College, communication.

Preparation
Reactivity: Precipitation, Acid-Base, and Redox Reactions

Preparation Questions