Reactivity: Precipitation

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

• To make thermodynamic predictions of cation/anion reactions in aqueous solution
• To carry out reactions of cations and anions in aqueous solution
• To devise solubility rules

A spontaneous process is one that takes place without assistance from outside. Examples are melting of ice at room temperature, rusting of iron, and combustion of natural gas in a furnace. A process is spontaneous either because it is exothermic (has DH < 0) or because it increases disorder (has DS > 0) or both. The effects of these two quantities on spontaneity are neatly summarized in terms of a single quantity, called the free energy, which is related to enthalpy (H) and entropy (S) by equation 1.

D G = D H - TDS (1)

Processes for which D G < 0 are spontaneous; those for which D G > 0 are not.

It is useful for chemists to tabulate the standard free energies of formation, D G_f^o, for many chemical substances. D G_f^o is the change in free energy which accompanies the formation (subscript f) of one mole of a substance in its standard state (superscript ^o) from its elements in their standard states. For example, the formation of one mole of CaCO₃(s) from its elements is shown below:

Ca(s) + C(s) + 3/2 O₂(g) ® CaCO₃(s)

The free energy change which accompanies this process is, by definition, the standard free energy of formation of CaCO₃. D G^o for any chemical reaction may be calculated from the D G_f^o values of reactants and products using equation 2. This is what makes D G_f^o so useful.

D G^o for reaction = S D G_f^o(products) - S D G_f^o(reactants) (2)

To calculate the free energy change for a chemical reaction under conditions other than standard, we need only correct DG^o for the non-standard conditions using equation 3.

DG = DG^o + RT ln Q(3)

Here R is the gas constant, in kJ/mole-K, T is the Kelvin temperature, and Q is the mass action quotient.

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.

AX + BY ® AY + BX(4)

For such a reaction to proceed to the right, there must be a thermodynamic 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. The free energy change, DG, for such a process is negative. This experiment deals with reactions that (potentially) form solids.

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

Ag⁺(aq) + NO₃^-(aq) + Na⁺(aq) + Cl^-(aq) ® AgCl(s) + Na⁺(aq) + NO₃^-(aq) (6)

Ag⁺(aq) + Cl^-(aq) ---> AgCl(s) (7)

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.

We can readily calculate DG for any net ionic equation from ion concentrations and DG_f^o values for reactants and products. Table 1 gives standard free energies of formation, DG_f^o, for 1 M solutions of several cations (first and second columns) and anions (third and fourth columns). The fifth and sixth columns show DG_f^o for several solid salts formed from combinations of the listed cations and anions. Note that two of the solids crystallize with waters of hydration (e.g., CaSO₄·2H₂O). Equations for formation of these solids must include the appropriate number of water molecules on the left side. DG_f^o for water is -237.18 kJ/mole.

DG_f^o Values in kJ/mol for Cations, Anions, and Solids

(Data from the CRC Handbook of Chemistry and Physics, 63rd Edition, CRC Press, Inc., Boca Raton, FL, 1982)

Using data from the table, we calculate DG^o for equation 7 as follows:

DG^o = DG_f^o (AgCl(s)) - DG_f^o(Ag⁺ (aq)) - DG_f^o(Cl^-(aq))

= -109.789 - 77.107 - (-131.228) kJ/mole

= -55.668 kJ/mole

This number is substantially negative. Formation of solid AgCl(s) from its ions under standard conditions (i.e., [Ag⁺] = [Cl^-] = 1 M) is spontaneous--it will occur of its own accord. If ion concentrations were 0.1 M rather than 1 M, would precipitation still occur? Use equation 3:

DG = DG^o + RT ln Q

DG = DG^o + RT ln 1/[Ag⁺][Cl^-]

= -55.668 + (8.314 x 10^-3)(298) ln [1/(0.1)(0.1)]

= -44.250 kJ/mole

Precipitation is still spontaneous, but somewhat less so than at 1 M ion concentrations.

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. Formation of a soluble salt from its ions has DG^o > 0; formation of an insoluble salt from its ions has DG^o < 0.

In this experiment, you investigate many cation-anion combinations to discover which produce insoluble products. First, before coming to lab, predict whether or not solid will form when specified cation/anion pairs are brought together. Then, the cation/anion combinations that you predicted, and many others as well, will be investigated in the laboratory to learn about double displacement reactions which form solids. 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. How accurate were your thermodynamic predictions? Prepare a table showing predicted and observed results for each cation/anion combination.
2. 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?
3. 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.

• 8.5 x 11" plastic reaction sheet
• cation/anion grid
• paper clips
• disposable micropipets
• magnifying lens
• Kimwipes
• microstirrers (toothpicks)
• wash bottle
• red litmus
• 6 M NaOH
• 1 M Na₂S
• solutions containing cations ( + nitrate anion) and anions (+ sodium cation)

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

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.

Clean the plastic reaction sheet by curling it and emptying the solutions and solids into a 250-mL beaker. Rinse remaining material into the beaker using a distilled water wash bottle. USE NO MORE THAN A TOTAL OF 100 mL OF WATER TO WASH THE SHEET. Finally, rinse the working side of the sheet with tap water and dry carefully with a towel, then with a Kimwipe. When the sheet is completely dry, return it and the paper grid sheet to the instructor.

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.

1. C. Scaife, Union College, communication.

Preparation
Reactivity: Precipitation

Preparation Questions